Oxygen

Singlet oxygen

2007-08-14T11:07:00.001-07:00

Singlet oxygen
From Wikipedia, the free encyclopedia
Jump to: navigation, search
This article needs additional references or sources for verification.
Please help to improve this article by adding reliable references. Unverifiable material may be challenged and removed.
This article has been tagged since July 2007.

Molecular orbital diagram for singlet oxygen. Quantum mechanics predicts that this configuration with the paired electrons is higher in energy than the triplet ground state.Singlet oxygen is the common name used for the two metastable states of molecular oxygen (O2) with higher energy than the ground state triplet oxygen [1]. The energy difference between the lowest energy of O2 in the singlet state and the lowest energy in the triplet state is about 3625 kelvin (Te (a¹Δg <- X³Σg-) = 7918.1 cm-1.)

Molecular oxygen differs from most molecules in having an open-shell triplet ground state, O2(X³Σg-). Molecular orbital theory predicts two low-lying excited singlet states O2(a¹Δg) and O2(b¹Σg+) (for nomenclature see article on Molecular term symbol). These electronic states differ only in the spin and the occupancy of oxygen's two degenerate antibonding πg-orbitals (see degenerate energy level). The O2(b¹Σg+)-state is very short lived and relaxes quickly to the lowest lying excited state, O2(a¹Δg). Thus, the O2(a¹Δg)-state is commonly referred to as singlet oxygen.

Contents [hide]
1 Physics
2 Chemistry
3 Biochemistry
4 External links
5 References

[edit] Physics
The energy difference between ground state and singlet oxygen is 94.2 kJ/mol and corresponds to a transition in the near-infrared at ~1270 nm. In the isolated molecule, the transition is strictly forbidden by spin, symmetry and parity selection rules, making it one of nature's most forbidden transitions. In other words, direct excitation of ground state oxygen by light to form singlet oxygen is very improbable. As a consequence, singlet oxygen in the gas phase is extremely long lived (72 minutes). Interaction with solvents, however, reduces the lifetime to microsecond or even nanoseconds.

Direct detection of singlet oxygen is possible through its extremely weak phosphorescence at 1270 nm, which is not visible to the eye. However, at high singlet oxygen concentrations, the fluorescence of the so-called singlet oxygen dimol (simultaneous emission from two singlet oxygen molecules upon collision) can be observed as a red glow at 634 nm [2].

[edit] Chemistry
The chemistry of singlet oxygen is different from that of ground state oxygen. Singlet oxygen can participate in Diels-Alder reactions and ene reactions. It can be generated in a photosensitized process by energy transfer from dye molecules such as rose bengal, methylene blue or porphyrins, or by chemical processes such as spontaneous decomposition of hydrogen trioxide in water or the reaction of hydrogen peroxide with hypochlorite [3]. Singlet oxygen reacts with an alkene -C=C-CH- by abstraction of the allylic proton in an ene reaction type reaction to the allyl hydroperoxide HO-O-C-C=C. It can then be reduced to the allyl alcohol. With some substrates dioxetanes are formed and cyclic dienes such as 1,3-Cyclohexadiene form [4+2]cycloaddition adducts. [4].

[edit] Biochemistry
In photosynthesis, singlet oxygen can be produced from the light-harvesting chlorophyll molecules. One of the roles of carotenoids in photosynthetic systems is to prevent damage caused by produced singlet oxygen by either removing excess light energy from chlorophyll molecules or quenching the singlet oxygen molecules directly.

In mammalian biology, singlet oxygen is a form of reactive oxygen species, which is linked to oxidation of LDL cholesterol and resultant cardiovascular effects. Polyphenol antioxidants can scavenge and reduce concentrations of reactive oxygen species and may prevent such deleterious oxidative effects [5].

Singlet oxygen is the active species in photodynamic therapy.

[edit] External links
The NIST webbook on oxygen
Photochemistry & Photobiology tutorial on Singlet Oxygen
Demonstration of the Red Singlet Oxygen Dimol Emission (Purdue University)

[edit] References
^ David R. Kearns (1971). "Physical and chemical properties of singlet molecular oxygen". Chemical Reviews 71 (4): 395 - 427. DOI:10.1021/cr60272a004.
^ Interpretation of the atmospheric oxygen bands; electronic levels of the oxygen molecule R.S. Mulliken Nature (journal) Volume 122, Page 505 1928
^ Physical Mechanisms of Generation and Deactivation of Singlet Oxygen C. Schweitzer, R. Schmidt Chemical Reviews Volume 103, Pages 1685-1757 2003
^ Carey, Francis A.; Sundberg, Richard J.; (1984). Advanced Organic Chemistry Part A Structure and Mechanisms (2nd ed.). New York N.Y.: Plenum Press. ISBN 0-306-41198-9.
^ Cell and Molecular Cell Biology concepts and experiments Fourth Edition. Gerald Karp. Page 223 2005
Retrieved from "http://en.wikipedia.org/wiki/Singlet_oxygen"
Categories: Articles lacking reliable references from July 2007 | Reagents for organic chemistry | Spectroscopy | Physical chemistry | Oxygen

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
日本語
Polski
Русский

This page was last modified 10:45, 30 July 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

neutron

2007-08-14T11:05:00.000-07:00

Neutron
From Wikipedia, the free encyclopedia
Jump to: navigation, search
This article is about the subatomic particle. For other uses, see Neutron (disambiguation).
This article is a discussion of neutrons in general. For the specific case of a neutron found outside the nucleus, see free neutron.
Neutron

The quark structure of the neutron.
Composition: one up, two down
Family: Fermion
Group: Quark
Interaction: Gravity, Electromagnetic, Weak, Strong
Antiparticle: Antineutron
Discovered: James Chadwick[1]
Symbol: n
Mass: 1.674 927 29(28) × 10−27kg
939.565 560(81) MeV/c²
1.008665 u
Electric charge: 0 C
Spin: ½
In physics, the neutron is a subatomic particle with no net electric charge and a mass of 939.573 MeV/c² or 1.008 664 915 (78) u (1.6749 × 10−27 kg, slightly more than a proton). Its spin is ½. Its antiparticle is called the antineutron. The neutron, along with the proton, is a nucleon.

The nucleus of all atoms (except the lightest isotope of hydrogen, which has only a single proton) consists of protons and neutrons. The number of neutrons determines the isotope of an element. For example, the carbon-12 isotope has 6 protons and 6 neutrons, while the carbon-14 isotope has 6 protons and 8 neutrons. Isotopes are atoms of the same element that have the same atomic number but different masses due to a different number of neutrons.

A neutron consists of two down quarks and one up quark. Since it has three quarks, it is classified it as a baryon.

Contents [hide]
1 Neutron Stability and Beta Decay
2 Interactions
3 Detection
4 Uses
5 Sources
6 Discovery
7 Anti-Neutron
8 Current developments
8.1 Electric dipole moment
8.2 Tetraneutrons
9 Protection
10 See also
10.1 Fields concerning neutrons
10.2 Types of neutrons
10.3 Objects containing neutrons
10.4 Neutron sources
10.5 Processes involving neutrons
11 References

[edit] Neutron Stability and Beta Decay

The Feynman diagram of the neutron beta decay processOutside the nucleus, free neutrons are unstable and have a mean lifetime of 885.7±0.8 seconds (about 15 minutes), decaying by emission of a negative electron and antineutrino to become a proton:[2] . This decay mode, known as beta decay, can also transform the character of neutrons within unstable nuclei.

Inside of a bound nucleus, protons can also transform via beta decay into neutrons. In this case, the transformation may occur by emission of a positive electron (also called a positron or an antielectron) and neutrino (instead of an antineutrino): . The transformation of a proton to a neutron inside of a nucleus is also possible through electron capture: . Positron capture by neutrons in nuclei that contain an excess of neutrons would also be possible, but is hindered due to the fact positrons are repelled by the nucleus, and furthermore, quickly annihilate when they encounter negative electrons.

When bound inside of a nucleus, the instability of a single neutron to beta decay is balanced against the instability that would be acquired by the nucleus as a whole if an additional proton were to participate in repulsive interactions with the other protons that are already present in the nucleus. As such, although free neutrons are unstable, bound neutrons are not necessarily so. The same reasoning explains why protons, which are stable in empty space, may transform into neutrons when bound inside of a nucleus.

Beta decay and electron capture are types of radioactive decay and are both governed by the weak interaction.

[edit] Interactions
The neutron interacts through all four fundamental interactions: the electromagnetic, weak nuclear, strong nuclear and gravitational interactions.

Although the neutron has zero net charge, it may interact electromagnetically in two ways: first, the neutron has a magnetic moment of the same order as the proton;[3] second, it is composed of electrically charged quarks. Thus, the electromagnetic interaction is primarily important to the neutron in deep inelastic scattering and in magnetic interactions.

The neutron experiences the weak interaction through beta decay into a proton, electron and electron antineutrino. It experiences the gravitational force as does any energetic body; however, gravity is so weak that it may be neglected in particle physics experiments.

The most important force to neutrons is the strong interaction. This interaction is responsible for the binding of the neutron's three quarks into a single particle. The residual strong force is responsible for the binding of neutrons and protons together into nuclei. This nuclear force plays the leading role when neutrons pass through matter. Unlike charged particles or photons, the neutron cannot lose energy by ionizing atoms. Rather, the neutron goes on its way unchecked until it makes a head-on collision with an atomic nucleus. For this reason, neutron radiation is extremely penetrating.

[edit] Detection
Main article: neutron detection
The common means of detecting a charged particle by looking for a track of ionization (such as in a cloud chamber) does not work for neutrons directly. Neutrons that elastically scatter off atoms can create an ionization track that is detectable, but the experiments are not as simple to carry out; other means for detecting neutrons, consisting of allowing them to interact with atomic nuclei, are more commonly used.

A common method for detecting neutrons involves converting the energy released from such reactions into electrical signals. The nuclides 3He, 6Li, 10B, 233U, 235U, 237Np and 239Pu are useful for this purpose. A good discussion on neutron detection is found in chapter 14 of the book Radiation Detection and Measurement by Glenn F. Knoll (John Wiley & Sons, 1979).

[edit] Uses
The neutron plays an important role in many nuclear reactions. For example, neutron capture often results in neutron activation, inducing radioactivity. In particular, knowledge of neutrons and their behavior has been important in the development of nuclear reactors and nuclear weapons.

Cold, thermal and hot neutron radiation is commonly employed in neutron scattering facilities, where the radiation is used in a similar way one uses X-rays for the analysis of condensed matter. Neutrons are complementary to the latter in terms of atomic contrasts by different scattering cross sections; sensitivity to magnetism; energy range for inelastic neutron spectroscopy; and deep penetration into matter.

The development of "neutron lenses" based on total internal reflection within hollow glass capillary tubes or by reflection from dimpled aluminum plates has driven ongoing research into neutron microscopy and neutron/gamma ray tomography.[4][5][6]

One use of neutron emitters is the detection of light nuclei, particularly the hydrogen found in water molecules. When a fast neutron collides with a light nucleus, it loses a large fraction of its energy. By measuring the rate at which slow neutrons return to the probe after reflecting off of hydrogen nuclei, a neutron probe may determine the water content in soil.

[edit] Sources
Due to the fact that free neutrons are unstable, they can be obtained only from nuclear disintegrations, nuclear reactions, and high-energy reactions (such as in cosmic radiation showers or accelerator collisions). Free neutron beams are obtained from neutron sources by neutron transport. For access to intense neutron sources, researchers must go to specialist facilities, such as the ISIS facility in the UK, which is currently the world's most intense pulsed neutron and muon source.

Neutrons' lack of total electric charge prevents engineers or experimentalists from being able to steer or accelerate them. Charged particles can be accelerated, decelerated, or deflected by electric or magnetic fields. However, these methods have no effect on neutrons except for a small effect of a magnetic field because of the neutron's magnetic moment.

[edit] Discovery
In 1930 Walther Bothe and H. Becker in Germany found that if the very energetic alpha particles emitted from polonium fell on certain light elements, specifically beryllium, boron, or lithium, an unusually penetrating radiation was produced. At first this radiation was thought to be gamma radiation although it was more penetrating than any gamma rays known, and the details of experimental results were very difficult to interpret on this basis. The next important contribution was reported in 1932 by Irène Joliot-Curie and Frédéric Joliot in Paris. They showed that if this unknown radiation fell on paraffin or any other hydrogen-containing compound it ejected protons of very high energy. This was not in itself inconsistent with the assumed gamma ray nature of the new radiation, but detailed quantitative analysis of the data became increasingly difficult to reconcile with such a hypothesis. Finally (later in 1932) the physicist James Chadwick in England performed a series of experiments showing that the gamma ray hypothesis was untenable. He suggested that in fact the new radiation consisted of uncharged particles of approximately the mass of the proton, and he performed a series of experiments verifying his suggestion. Such uncharged particles were eventually called neutrons, apparently from the Latin root for neutral and the Greek ending -on (by imitation of electron and proton).

[edit] Anti-Neutron
Main article: antineutron
The antineutron is the antiparticle of the neutron. It was discovered by Bruce Cork in the year 1956, a year after the antiproton was discovered.

CPT-symmetry puts strong constraints on the relative properties of particles and antiparticles and, therefore, is open to stringent tests. The fractional difference in the masses of the neutron and antineutron is (9±5)×10−5. Since the difference is only about 2 standard deviations away from zero, this does not give any convincing evidence of CPT-violation.[3]

[edit] Current developments

[edit] Electric dipole moment
An experiment at the Institut Laue-Langevin (ILL) has attempted to measure an electric dipole, or separation of charges, within the neutron, and is consistent with an electric dipole moment of zero. These results are important in developing theories that go beyond the Standard Model. See FRONTIERS article, and the experiment's web page.

[edit] Tetraneutrons
The existence of stable clusters of four neutrons, or tetraneutrons, has been hypothesised by a team led by Francisco-Miguel Marqués at the CNRS Laboratory for Nuclear Physics based on observations of the disintegration of beryllium-14 nuclei. This is particularly interesting, because current theory suggests that these clusters should not be stable.

[edit] Protection
Exposure to neutrons can be hazardous, since the interaction of neutrons with molecules in the body can cause disruption to molecules and atoms, and can also cause reactions which give rise to other forms of radiation. The normal expectations of radiation protection apply: avoid exposure, stay as far from the source as possible, and keep exposure time to the minimum. Some thought must however be given to how to protect oneselves from such exposure. For other types of radiation, e.g. alpha particles, beta particles, or gamma rays, material of a high atomic number and with high density makes for good shielding; frequently lead is used. However, this approach will not work with neutrons, since the absorption of neutrons does not increase straightforwardly with atomic number as it does with alpha, beta, and gamma radiation. Instead one needs to look at the particular interactions neutrons have with matter (see the section on detection above). For example, hydrogen rich materials are often used since ordinary hydrogen scatters neutrons, so this often means simple concrete blocks, or paraffin loaded plastic blocks may be the best protection.

[edit] See also

[edit] Fields concerning neutrons
particle physics
quark model
chemistry
Neutron Detection
Neutron Scattering

[edit] Types of neutrons
nucleon
fast neutron
free neutron
thermal neutron
neutron radiation and the Sievert radiation scale
neutron temperature, used to classify neutron types

[edit] Objects containing neutrons
nucleus
dineutron
tetraneutron
neutronium
neutron star

[edit] Neutron sources
Neutron sources
Neutron generator

[edit] Processes involving neutrons
neutron transport
neutron diffraction
neutron bomb

[hide]v • d • eParticles in physics
elementary particles Elementary fermions: Quarks: u · d · s · c · b · t • Leptons: e · μ · τ · νe · νμ · ντ
Elementary bosons: Gauge bosons: γ · g · W± · Z0 • Ghosts
Composite particles Hadrons: Baryons(list)/Hyperons/Nucleons: p · n · Δ · Λ · Σ · Ξ · Ω · Ξb • Mesons(list)/Quarkonia: π · K · ρ · J/ψ · Υ
Other: Atomic nucleus • Atoms • Molecules • Positronium
Hypothetical elementary particles Superpartners: Axino · Dilatino · Chargino · Gluino · Gravitino · Higgsino · Neutralino · Sfermion · Slepton · Squark
Other: Axion · Dilaton · Goldstone boson · Graviton · Higgs boson · Tachyon · X · Y · W' · Z'
Hypothetical composite particles Exotic hadrons: Exotic baryons: Pentaquark • Exotic mesons: Glueball · Tetraquark
Other: Mesonic molecule
Quasiparticles Davydov soliton · Exciton · Magnon · Phonon · Plasmon · Polariton · Polaron

[edit] References
^ 1935 Nobel Prize in Physics
^ Particle Data Group Summary Data Table on Baryons
^ a b Particle Data Group's Review of Particle Physics 2006
^ Nature 357, 390-391 (04 June 1992); doi:10.1038/357390a0
^ Physorg.com, "New Way of 'Seeing': A 'Neutron Microscope'"
^ NASA.gov: "NASA Develops a Nugget to Search for Life in Space"
Retrieved from "http://en.wikipedia.org/wiki/Neutron"
Categories: Neutron | Fundamental physics concepts

Decay Product

2007-08-14T11:03:00.000-07:00

Decay product
From Wikipedia, the free encyclopedia
Jump to: navigation, search
In nuclear physics, a decay product, also known as a daughter product, daughter isotope or daughter nuclide, is a nuclide resulting from the radioactive decay of a parent isotope or precursor nuclide. The daughter product may be stable or it may decay to form a daughter product of its own. The daughter of a daughter product is sometimes called a granddaughter product.

Decay products are extremely important in understanding radioactive decay and the management of radioactive waste.

In practice nearly all decay products are themselves radioactive. The result of this is that most radionuclides do not have simply a decay product, but rather a decay chain, leading eventually to a stable nuclide. For elements above lead in atomic number, this is nearly always an isotope of lead. Lead is generally the stable point at which decay chains stop.

In many cases members of the decay chain are far more radioactive than the original nuclide. Thus, although uranium is not dangerously radioactive when pure, some pieces of naturally-occurring pitchblende are quite dangerous owing to their radium content. Similarly, thorium gas mantles are very slightly radioactive when new, but become far more radioactive after only a few months of storage.

Although it cannot be predicted whether any given atom of a radioactive substance will decay at any given time, the decay products of a radioactive substance are extremely predictable. Because of this, decay products are important to scientists in many fields who need to know the quantity or type of the parent product. Such studies are done to measure pollution levels (in and around nuclear facilities) and for other matters.

Retrieved from "http://en.wikipedia.org/wiki/Decay_product"
Categories: Nuclear physics | Nuclear chemistry

Decay Energy

2007-08-14T11:02:00.000-07:00

Decay energy
From Wikipedia, the free encyclopedia
Jump to: navigation, search
The decay energy is the energy released by a nuclear decay.

The difference between the mass of the reactants and the mass of products is often written as Q:

Q = (mass of reactants) - (mass of products)
This can be expressed as energy by Albert Einstein's famous formula E=mc².

Types of radioactive decay include

gamma radiation
beta decay
alpha decay

[edit] External links
University of Waterloo science
This chemistry article is a stub. You can help Wikipedia by expanding it.

Retrieved from "http://en.wikipedia.org/wiki/Decay_energy"
Categories: Chemistry stubs | Nuclear chemistry

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
File upload wizard
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
العربية
Asturianu
Català
Deutsch
Español
Français
Magyar
Nederlands
日本語
Polski
Português
Svenska

This page was last modified 20:19, 21 May 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Radioactive Decay

2007-08-14T10:58:00.000-07:00

Radioactive decay
From Wikipedia, the free encyclopedia
Jump to: navigation, search
"Radioactive" and "Radioactivity" redirect here. For other uses see Radioactive (disambiguation).
For decay rate in a more general context see Particle decay.
Radioactive decay is the process in which an unstable atomic nucleus loses energy by emitting radiation in the form of particles or electromagnetic waves. This decay, or loss of energy, results in an atom of one type, called the parent nuclide transforming to an atom of a different type, called the daughter nuclide. For example: a carbon-14 atom (the "parent") emits radiation and transforms to a nitrogen-14 atom (the "daughter.") This is a random process on the atomic level, in that it is impossible to predict when a particular atom will decay, but given a large number of similar atoms, the decay rate, on average, is predictable.

The trefoil symbol is used to indicate radioactive material.The SI unit of radioactive decay is the becquerel (Bq). One Bq is defined as one transformation (or decay) per second. Since any reasonably-sized sample of radioactive material contains many atoms, a Bq is a tiny measure of activity; amounts on the order of TBq (terabecquerels) or GBq (gigabecquerels) are commonly used. Another unit of decay is the curie, which was originally defined as the radioactivity of one gram of pure radium, and is equal to 3.7 × 1010 Bq.

Contents [hide]
1 Explanation
2 Discovery
3 Modes of decay
4 Decay chains and multiple modes
5 Occurrence and applications
6 Radioactive decay rates
6.1 Activity measurements
7 Decay timing
8 References
9 See also
10 External links

[edit] Explanation
The neutrons and protons that constitute nuclei, as well as other particles that may approach them, are governed by several interactions. The strong nuclear force, not observed at the familiar macroscopic scale, is the most powerful force over subatomic distances. The electrostatic force is also significant. Of lesser importance is the weak nuclear force.

The interplay of these forces is simple. Some configurations of the particles in a nucleus have the property that, should they shift ever so slightly, the particles could fall into a lower-energy arrangement (with the extra energy moving elsewhere). One might draw an analogy with a snowfield on a mountain: while friction between the snow crystals can support the snow's weight, the system is inherently unstable with regards to a lower-potential-energy state, and a disturbance may facilitate the path to a greater entropy state (i.e., towards the ground state where heat will be produced, and thus total energy is distributed over a larger number of quantum states). Thus, an avalanche results. The total energy does not change in this process, but because of entropy effects, avalanches only happen in one direction, and the end of this direction, which is dictated by the largest number of chance-mediated ways to distribute available energy, is what we commonly refer to as the "ground state."

Such a collapse (a decay event) requires a specific activation energy. In the case of a snow avalanche, this energy classically comes as a disturbance from outside the system, although such disturbances can be arbitrarily small. In the case of an excited atomic nucleus, the arbitrarily small disturbance comes from quantum vacuum fluctuations. A nucleus (or any excited system in quantum mechanics) is unstable, and can thus spontaneously stabilize to a less-excited system. This process is driven by entropy considerations: the energy does not change, but at the end of the process, the total energy is more diffused in spacial volume. The resulting transformation alters the structure of the nucleus. Such a reaction is thus a nuclear reaction, in contrast to chemical reactions, which also are driven by entropy, but which involve changes in the arrangement of the outer electrons of atoms, rather than their nuclei.

Some nuclear reactions do involve external sources of energy, in the form of collisions with outside particles. However, these are not considered decay. Rather, they are examples of induced nuclear reactions. Nuclear fission and fusion are common types of induced nuclear reactions.

[edit] Discovery
Radioactivity was first discovered in 1896 by the French scientist Henri Becquerel while working on phosphorescent materials. These materials glow in the dark after exposure to light, and he thought that the glow produced in cathode ray tubes by X-rays might somehow be connected with phosphorescence. So he tried wrapping a photographic plate in black paper and placing various phosphorescent minerals on it. All results were negative until he tried using uranium salts. The result with these compounds was a deep blackening of the plate.

However, it soon became clear that the blackening of the plate had nothing to do with phosphorescence because the plate blackened when the mineral was kept in the dark. Also non-phosphorescent salts of uranium and even metallic uranium blackened the plate. Clearly there was some new form of radiation that could pass through paper that was causing the plate to blacken.

Alpha particles may be completely stopped by a sheet of paper, beta particles by aluminum shielding. Gamma rays, however, can only be reduced by much more substantial obstacles, such as a very thick piece of lead.At first it seemed that the new radiation was similar to the then recently discovered X-rays. However further research by Becquerel, Marie Curie, Pierre Curie, Ernest Rutherford and others discovered that radioactivity was significantly more complicated. Different types of decay can occur, but Rutherford was the first to realize that they all occur with the same mathematical approximately exponential formula (see below).

As for types of radioactive radiation, it was found that an electric or magnetic field could split such emissions into three types of beams. For lack of better terms, the rays were given the alphabetic names alpha, beta, and gamma, names they still hold today. It was immediately obvious from the direction of electromagnetic forces that alpha rays carried a positive charge, beta rays carried a negative charge, and gamma rays were neutral. From the magnitude of deflection, it was also clear that alpha particles were much more massive than beta particles. Passing alpha rays through a thin glass membrane and trapping them in a discharge tube allowed researchers to study the emission spectrum of the resulting gas, and ultimately prove that alpha particles are in fact helium nuclei. Other experiments showed the similarity between beta radiation and cathode rays; they are both streams of electrons, and between gamma radiation and X-rays, which are both high energy electromagnetic radiation.

Although alpha, beta, and gamma are most common, other types of decay were eventually discovered. Shortly after discovery of the neutron in 1932, it was discovered by Enrico Fermi that certain rare decay reactions give rise to neutrons as a decay particle. Isolated proton emission was also eventually observed in some elements. Shortly after the discovery of the positron in cosmic ray products, it was realized that the same process that operates in classical beta decay can also produce positrons (positron emission), analogously to negative electrons. Each of the two types of beta decay acts to move a nucleus toward a ratio of neutrons and protons which has the least energy for the combination. Finally, in a phenomenon called cluster decay, specific combinations of neutrons and protons other than alpha particles were found to occasionally spontaneously be emitted from atoms.

Still other types of radioactive decay were found which emit previously seen particles, but by different mechanisms. An example is internal conversion, which results in electron and sometimes high energy photon emission, even though it involves neither beta nor gamma decay.

The early researchers also discovered that many other chemical elements besides uranium have radioactive isotopes. A systematic search for the total radioactivity in uranium ores also guided Marie Curie to isolate a new element polonium and to separate a new element radium from barium; the two elements' chemical similarity would otherwise have made them difficult to distinguish.

The dangers of radioactivity and of radiation were not immediately recognized. Acute effects of radiation were first observed in the use of X-rays when the Serbo-Croatian-American electric engineer Nikola Tesla intentionally subjected his fingers to X-rays in 1896. He published his observations concerning the burns that developed, though he attributed them to ozone rather than to the X-rays. Fortunately his injuries healed later.

The genetic effects of radiation, including the effects on cancer risk, were recognized much later. It was only in 1927 that Hermann Joseph Muller published his research that showed the genetic effects. In 1946 he was awarded the Nobel prize for his findings.

Before the biological effects of radiation were known, many physicians and corporations had begun marketing radioactive substances as patent medicine and Radioactive quackery; particularly alarming examples were radium enema treatments, and radium-containing waters to be drunk as tonics. Marie Curie spoke out against this sort of treatment, warning that the effects of radiation on the human body were not well understood (Curie later died from aplastic anemia assumed due to her own work with radium, but later examination of her bones showed that she had been a careful laboratory worker and had a low burden of radium; a better candidate for her disease was her long exposure to unshielded X-ray tubes while a volunteer medical worker in WW I). By the 1930s, after a number of cases of bone-necrosis and death in enthusiasts, radium-containing medical products had nearly vanished from the market.

[edit] Modes of decay
Radionuclides can undergo a number of different reactions. These are summarized in the following table. A nucleus with positive charge (atomic number) Z and atomic weight A is represented as (A, Z).

Mode of decay Participating particles Daughter nucleus
Decays with emission of nucleons:
Alpha decay An alpha particle (A=4, Z=2) emitted from nucleus (A-4, Z-2)
Proton emission A proton ejected from nucleus (A-1, Z-1)
Neutron emission A neutron ejected from nucleus (A-1, Z)
Double proton emission Two protons ejected from nucleus simultaneously (A-2, Z-2)
Spontaneous fission Nucleus disintegrates into two or more smaller nuclei and other particles -
Cluster decay Nucleus emits a specific type of smaller nucleus (A1, Z1) larger than an alpha particle (A-A1, Z-Z1) + (A1,Z1)
Different modes of beta decay:
Beta-Negative decay A nucleus emits an electron and an antineutrino (A, Z+1)
Positron emission, also Beta-Positive decay A nucleus emits a positron and a neutrino (A, Z-1)
Electron capture A nucleus captures an orbiting electron and emits a neutrino - The daughter nucleus is left in an excited and unstable state (A, Z-1)
Double beta decay A nucleus emits two electrons and two antineutrinos (A, Z+2)
Double electron capture A nucleus absorbs two orbital electrons and emits two neutrinos - The daughter nucleus is left in an excited and unstable state (A, Z-2)
Electron capture with positron emission A nucleus absorbs one orbital electron, emits one positron and two neutrinos (A, Z-2)
Double positron emission A nucleus emits two positrons and two neutrinos (A, Z-2)
Transitions between states of the same nucleus:
Gamma decay Excited nucleus releases a high-energy photon (gamma ray) (A, Z)
Internal conversion Excited nucleus transfers energy to an orbital electron and it is ejected from the atom (A, Z)

Radioactive decay results in a reduction of summed rest mass, which is converted to energy (the disintegration energy) according to the formula E = mc2. This energy is released as kinetic energy of the emitted particles. The energy remains associated with a measure of mass of the decay system invariant mass, inasmuch the kinetic energy of emitted particles contributes also to the total invariant mass of systems. Thus, the sum of rest masses of particles is not conserved in decay, but the system mass or system invariant mass (as also system total energy) is conserved.

[edit] Decay chains and multiple modes
The daughter nuclide of a decay event is usually also unstable, sometimes even more unstable than the parent. If this is the case, it will proceed to decay again. A sequence of several decay events, producing in the end a stable nuclide, is a decay chain.

Many radionuclides have several different observed modes of decay. Bismuth-212, for example, has three. Thus a given nuclide may lead to several different decay chains.

Of the commonly occurring forms of radioactive decay, the only one that changes the number of aggregate protons and neutrons (nucleons) contained in the nucleus is alpha emission, which reduces it by four. Thus, the number of nucleons modulo 4 is preserved across any decay chain.

[edit] Occurrence and applications
According to the Big Bang theory, radioactive isotopes of the lightest elements (H, He, and traces of Li) were produced very shortly after the emergence of the universe. However, these nuclides are so highly unstable that virtually none of them have survived to today. Most radioactive nuclei are therefore relatively young, having formed in stars (particularly supernovae) and during ongoing interactions between stable isotopes and energetic particles. For example, carbon-14, a radioactive nuclide with a half-life of only 5730 years, is constantly produced in Earth's upper atmosphere due to interactions between cosmic rays and nitrogen.

Radioactive decay has been put to use in the technique of radioisotopic labeling, used to track the passage of a chemical substance through a complex system (such as a living organism). A sample of the substance is synthesized with a high concentration of unstable atoms. The presence of the substance in one or another part of the system is determined by detecting the locations of decay events.

On the premise that radioactive decay is truly random (rather than merely chaotic), it has been used in hardware random-number generators. Because the process is not thought to vary significantly in mechanism over time, it is also a valuable tool in estimating the absolute ages of certain materials. For geological materials, the radioisotopes and certain of their decay products become trapped when a rock solidifies, and can then later be used (subject to many well-known qualifications) to estimate the date of the solidification. These include checking the results of several simultaneous processes and their products against each other, within the same sample.

[edit] Radioactive decay rates
The decay rate, or activity, of a radioactive substance are characterized by:

Constant quantities:

half life — symbol t1 / 2 — the time for half of a substance to decay.
mean lifetime — symbol τ — the average lifetime of any given particle.
decay constant — symbol λ — the inverse of the mean lifetime.
(Note that although these are constants, they are associated with statistically random behavior of substances, and predictions using these constants are less accurate for small number of atoms.)
Time-variable quantities:

Total activity — symbol A — number of decays an object undergoes per second.
Number of particles — symbol N — the total number of particles in the sample.
Specific activity — symbol SA — number of decays per second per amount of substance. The "amount of substance" can be the unit of either mass or volume.)
These are related as follows:

where
is the initial amount of active substance — substance that has the same percentage of unstable particles as when the substance was formed.

[edit] Activity measurements
The units in which activities are measured are: becquerel (symbol Bq) = number of disintegrations per second; curie (Ci) = 3.7 × 1010 disintegrations per second. Low activities are also measured in disintegrations per minute (dpm).

[edit] Decay timing
See also: exponential decay
As discussed above, the decay of an unstable nucleus is entirely random and it is impossible to predict when a particular atom will decay. However, it is equally likely to decay at any time. Therefore, given a sample of a particular radioisotope, the number of decay events –dN expected to occur in a small interval of time dt is proportional to the number of atoms present. If N is the number of atoms, then the probability of decay (– dN/N) is proportional to dt:

Particular radionuclides decay at different rates, each having its own decay constant (λ). The negative sign indicates that N decreases with each decay event. The solution to this first-order differential equation is the following function:

This function represents exponential decay. It is only an approximate solution, for two reasons. Firstly, the exponential function is continuous, but the physical quantity N can only take non-negative integer values. Secondly, because it describes a random process, it is only statistically true. However, in most common cases, N is a very large number and the function is a good approximation.

In addition to the decay constant, radioactive decay is sometimes characterized by the mean lifetime. Each atom "lives" for a finite amount of time before it decays, and the mean lifetime is the arithmetic mean of all the atoms' lifetimes. It is represented by the symbol τ, and is related to the decay constant as follows:

A more commonly used parameter is the half-life. Given a sample of a particular radionuclide, the half-life is the time taken for half the radionuclide's atoms to decay. The half life is related to the decay constant as follows:

This relationship between the half-life and the decay constant shows that highly radioactive substances are quickly spent, while those that radiate weakly endure longer. Half-lives of known radionuclides vary widely, from more than 1019 years (such as for very nearly stable nuclides, e.g. 209Bi), to 10-23 seconds for highly unstable ones.

[edit] References
"Radioactivity", Encyclopædia Britannica. 2006. Encyclopædia Britannica Online. 18 Dec. 2006

[edit] See also
Nuclear pharmacy
Nuclear physics
Radioactivity in biology
Poisson process
Radiation
Radiation therapy
Radioactive contamination
Radiometric dating
Actinides in the environment
Half-life
Fallout shelter
Particle decay

[edit] External links
Look up radioactivity in
Wiktionary, the free dictionary.General information
General information, with emphasis on different modes
Some numerical calculations based on the Uranium-232 decay chain
Nomenclature of nuclear chemistry
Some theoretical questions of nuclear stability
Decay heat rate|quantity calculation
Specific activity and related topics.
The Lund/LBNL Nuclear Data Search - Contains tabulated information on radioactive decay types and energies.
Retrieved from "http://en.wikipedia.org/wiki/Radioactive_decay"
Categories: Exponentials | Radioactivity

Half Life Period

2007-08-14T10:57:00.000-07:00

Half-life
From Wikipedia, the free encyclopedia
Jump to: navigation, search
This article is about the scientific and mathematical term. For other uses, see Half-life (disambiguation).
The half-life of a quantity, subject to exponential decay, is the time required for the quantity to decay to half of its initial value. The concept originated in the study of radioactive decay, but applies to many other fields as well, including phenomena which are described by non-exponential decays.

The term half-life was coined in 1907, but it was always referred to as half-life period. It was not until the early 1950s that the word period was dropped from the name. [1]

Number of
half-lives
elapsed Fraction
remaining As
power
of 2
0 1/1 1 / 20
1 1/2 1 / 21
2 1/4 1 / 22
3 1/8 1 / 23
4 1/16 1 / 24
5 1/32 1 / 25
6 1/64 1 / 26
7 1/128 1 / 27
... ... ...
N 1 / 2N 1 / 2N
The table at right shows the reduction of the quantity in terms of the number of half-lives elapsed.

It can be shown that, for exponential decay, the half-life t1 / 2 obeys this relation:

where

ln(2) is the natural logarithm of 2 (approximately 0.693), and
λ is the decay constant, a positive constant used to describe the rate of exponential decay.
The half-life is related to the mean lifetime τ by the following relation:

Contents [hide]
1 Examples
2 Decay by two or more processes
3 Derivation
4 Experimental determination
5 See also
6 References
7 External links

[edit] Examples
Main article: Exponential decay--Applications and examples
The constant λ can represent many different specific physical quantities, depending on what process is being described.

In an RC circuit or RL circuit, λ is the reciprocal of the circuit's time constant. For simple RC and RL circuits, λ equals 1 / RC or R / L, respectively.
In first-order chemical reactions, λ is the reaction rate constant.
In radioactive decay, it describes the probability of decay per unit time: dN = λNdt, where dN is the number of nuclei decayed during the time dt, and N is the quantity of radioactive nuclei.
In biology (specifically pharmacokinetics), from MeSH: Half-Life: The time it takes for a substance (drug, radioactive nuclide, or other) to lose half of its pharmacologic, physiologic, or radiologic activity. Year introduced: 1974 (1971).

[edit] Decay by two or more processes
Some quantities decay by two processes simultaneously (see Decay by two or more processes). In a fashion similar to the previous section, we can calculate the new total half-life T1 / 2 and we'll find it to be:

or, in terms of the two half-lives t1 and t2

i.e., half their harmonic mean.

[edit] Derivation
Quantities that are subject to exponential decay are commonly denoted by the symbol N. (This convention suggests a decaying number of discrete items. This interpretation is valid in many, but not all, cases of exponential decay.) If the quantity is denoted by the symbol N, the value of N at a time t is given by the formula:

where N0 is the initial value of N (at t = 0)

When t = 0, the exponential is equal to 1, and N(t) is equal to N0. As t approaches infinity, the exponential approaches zero. In particular, there is a time such that

Substituting into the formula above, we have

[edit] Experimental determination
The half-life of a process can be determined easily by experiment. In fact, some methods do not require advance knowledge of the law governing the decay rate, be it exponential decay or another pattern.

Most appropriate to validate the concept of half-life for radioactive decay, in particular when dealing with a small number of atoms, is to perform experiments and correct computer simulations. See in [1] how to test the behavior of the last atoms. Validation of physics-math models consists in comparing the model's behavior with experimental observations of real physical systems or valid simulations (physical and/or computer). The references given here describe how to test the validity of the exponential formula for small number of atoms with simple simulations, experiments, and computer code.

In radioactive decay, the exponential model does not apply for a small number of atoms (or a small number of atoms is not within the domain of validity of the formula or equation or table). The DIY experiments use pennies or M&M's candies. [2], [3]. A similar experiment is performed with isotopes of a very short half-life, for example, see Fig 5 in [4]. See how to write a computer program that simulates radioactive decay including the required randomness in [5] and experience the behavior of the last atoms. Of particular note, atoms undergo radioactive decay in whole units, and so after enough half-lives the remaining original quantity becomes an actual zero rather than asymptotically approaching zero as with continuous systems.

[edit] See also
Look up half-life in
Wiktionary, the free dictionary.Exponential decay
Mean lifetime
Elimination half-life
For non-exponential decays, see half-life in the article Rate equation

[edit] References
^ John Ayto "20th Century Words" (1999) Cambridge University Press.

[edit] External links
Time constant [6]
Retrieved from "http://en.wikipedia.org/wiki/Half-life"
Categories: Radioactivity | Exponentials | Chemical kinetics

Natural Abundance

2007-08-14T10:56:00.000-07:00

Natural abundance
From Wikipedia, the free encyclopedia
Jump to: navigation, search
This article or section is in need of attention from an expert on the subject.
WikiProject Chemistry or the Chemistry Portal may be able to help recruit one.
If a more appropriate WikiProject or portal exists, please adjust this template accordingly.

In chemistry, natural abundance (NA) refers to the prevalence of isotopes of a chemical element as naturally found on a planet. The relative atomic mass (a weighted average) of these isotopes is the atomic weight listed for the element in the periodic table. The abundance of an isotope varies from planet to planet but remains relatively constant in time.

As an example, uranium has three naturally occurring isotopes: U-238, U-235 and U-234. Their respective NA is 99.2745%, 0.72% and 0.0055%. For example, if 100,000 uranium atoms were analyzed, one would expect to find approximately 99,275 U-238 atoms, 720 U-235 atoms, and no more than 5 or 6 U-234 atoms. This is because U-238 is much more stable than U-235 or U-234, as the half-life of each isotope reveals: 4.468×109 years for U-238 compared to 7.038×108 years for U-235 and 245,500 years for U-234.

[edit] See also
Abundance of the chemical elements

This chemistry article is a stub. You can help Wikipedia by expanding it.

This physics-related article is a stub. You can help Wikipedia by expanding it.

Retrieved from "http://en.wikipedia.org/wiki/Natural_abundance"
Categories: Chemistry articles needing expert attention | Articles needing expert attention | Chemical properties | Chemistry stubs | Physics stubs

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
Afrikaans
العربية
Asturianu
Català
Deutsch
Español
한국어
Italiano
Magyar
Nederlands
日本語
Plattdüütsch
Português
Slovenščina
中文

This page was last modified 16:16, 5 July 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Isotope

2007-08-14T10:55:00.001-07:00

From Wikipedia, the free encyclopedia
Jump to: navigation, search
For other uses, see Isotope (disambiguation).
Isotopes are any of the several different forms of an element each having different atomic mass (mass number). Isotopes of an element have nuclei with the same number of protons (the same atomic number) but different numbers of neutrons. Therefore, isotopes have different mass numbers, which give the total number of nucleons—the number of protons plus neutrons.

A nuclide is any particular atomic nucleus with a specific atomic number Z and mass number A; it is equivalently an atomic nucleus with a specific number of protons and neutrons. Collectively, all the isotopes of all the elements form the set of nuclides. The distinction between the terms isotope and nuclide has somewhat blurred, and they are often used interchangeably. Isotope is best used when referring to several different nuclides of the same element; nuclide is more generic and is used when referencing only one nucleus or several nuclei of different elements. For example, it is more correct to say that an element such as fluorine consists of one stable nuclide rather than that it has one stable isotope.

In IUPAC nomenclature, isotopes and nuclides are specified by the name of the particular element, implicitly giving the atomic number, followed by a hyphen and the mass number (e.g. helium-3, carbon-12, carbon-13, iodine-131 and uranium-238). In symbolic form, the number of nucleons is denoted as a superscripted prefix to the chemical symbol (e.g. 3He, 12C, 13C, 131I and 238U).

The term isotope was coined in 1913 by Margaret Todd, a Scottish doctor, during a conversation with Frederick Soddy (to whom she was distantly related by marriage). Soddy, a chemist at Glasgow University, explained that it appeared from his investigations as if several elements occupied each position in the periodic table. Hence Todd suggested the Greek for "at the same place" as a suitable name. Soddy adopted the term and went on to win the Nobel Prize for Chemistry in 1921 for his work on radioactive substances.

In the bottom right corner of JJ Thomson's photographic plate are markings for the two isotopes of neon: neon-20 and neon-22.In 1913, as part of his exploration into the composition of canal rays, JJ Thomson channeled a stream of ionized neon through a magnetic and an electric field and measured its deflection by placing a photographic plate in its path. Thomson observed two patches of light on the photographic plate (see image on right), which suggested two different parabolas of deflection. Thomson concluded that some of the atoms in the gas were of higher mass than the rest.

Contents [hide]
1 Variation in properties between isotopes
2 Occurrence in nature
3 Molecular mass of isotopes
4 Applications of isotopes
4.1 Use of chemical properties
4.2 Use of nuclear properties
5 See also
6 External links

[edit] Variation in properties between isotopes
A neutral atom has the same number of electrons as protons. Thus, different isotopes of a given element all have the same number of protons and electrons and the same electronic structure; because the chemical behavior of an atom is largely determined by its electronic structure, isotopes exhibit nearly identical chemical behavior. The main exception to this is the kinetic isotope effect: due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element.

This "mass effect" is most pronounced for protium (1H) vis-à-vis deuterium (2H), because deuterium has twice the mass of protium. For heavier elements the relative mass difference between isotopes is much less, and the mass effect is usually negligible.

Similarly, two molecules which differ only in the isotopic nature of their atoms (isotopologues) will have identical electronic structure and therefore almost indistinguishable physical and chemical properties (again with deuterium providing the primary exception to this rule). The vibrational modes of a molecule are determined by its shape and by the masses of its constituent atoms. Consequently, isotopologues will have different sets of vibrational modes. Since vibrational modes allow a molecule to absorb photons of corresponding energies, isotopologues have different optical properties in the infrared range.

Although isotopes exhibit nearly identical electronic and chemical behavior, their nuclear behavior varies dramatically. Atomic nuclei consist of protons and neutrons bound together by the strong nuclear force. Because protons are positively charged, they repel each other. Neutrons, which are electrically neutral, allow some separation between the positively charged protons, reducing the electrostatic repulsion. Neutrons also stabilize the nucleus because at short ranges they attract each other and protons equally by the strong nuclear force, and this also offsets the electrical repulsion between protons. For this reason, one or more neutrons are necessary for two or more protons to be bound into a nucleus. As the number of protons increases, additional neutrons are needed to form a stable nucleus; for example, although the neutron to proton ratio of 3He is 1:2, the neutron/proton ratio of 238U is greater than 3:2. If too many or too few neutrons are present, the nucleus is unstable and subject to nuclear decay.

[edit] Occurrence in nature
Most elements have several different isotopes that can be found in nature. The relative abundance of an isotope is strongly correlated with its tendency toward nuclear decay; short-lived nuclides quickly decay away, while their long-lived counterparts endure. However, this does not mean that short-lived species disappear entirely; many are continually produced through the decay of longer-lived nuclides. Also, short-lived isotopes such as those of promethium have been detected in the spectra of stars, where they presumably are being continuously made by stellar nucleosynthesis. The tabulated atomic masses of elements are averages that account for the presence of multiple isotopes with different masses.

According to generally accepted cosmology, virtually all nuclides other than isotopes of hydrogen and helium (and traces of some isotopes of lithium, beryllium and boron-- see big bang nucleosynthesis) were built in stars and supernovae. Their respective abundances here result from the quantities formed by these processes, their spread through the galaxy, and their rates of decay. After the initial coalescence of the solar system, isotopes were redistributed according to mass. The isotopic composition of elements is different on different planets, making it possible to determine the origin of meteorites.

Isotopes of oxygen

2007-08-14T10:53:00.000-07:00

Isotopes of oxygen
From Wikipedia, the free encyclopedia
Jump to: navigation, search
Oxygen (O)
Standard atomic mass: 15.9994(3) u
The isotopes of oxygen include 3 stable nuclei and 14 unstable nuclei.

Contents [hide]
1 Table
1.1 Notes
2 References
3 See also

[edit] Table
nuclide
symbol Z(p) N(n)
isotopic mass (u)
half-life nuclear
spin representative
isotopic
composition
(mole fraction) range of natural
variation
(mole fraction)
excitation energy
12O 8 4 12.034405(20) 580(30)E-24 s [0.40(25) MeV] 0+
13O 8 5 13.024812(10) 8.58(5) ms (3/2-)
14O 8 6 14.00859625(12) 70.598(18) s 0+
15O 8 7 15.0030656(5) 122.24(16) s 1/2-
16O 8 8 15.99491461956(16) STABLE 0+ 0.99757(16) 0.99738-0.99776
17O 8 9 16.99913170(12) STABLE 5/2+ 0.00038(1) 0.00037-0.00040
18O 8 10 17.9991610(7) STABLE 0+ 0.00205(14) 0.00188-0.00222
19O 8 11 19.003580(3) 26.464(9) s 5/2+
20O 8 12 20.0040767(12) 13.51(5) s 0+
21O 8 13 21.008656(13) 3.42(10) s (1/2,3/2,5/2)+
22O 8 14 22.00997(6) 2.25(15) s 0+
23O 8 15 23.01569(13) 82(37) ms 1/2+#
24O 8 16 24.02047(25) 65(5) ms 0+
25O 8 17 25.02946(28)# <50 ns (3/2+)#
26O 8 18 26.03834(28)# <40 ns 0+
27O 8 19 27.04826(54)# <260 ns 3/2+#
28O 8 20 28.05781(64)# <100 ns 0+

[edit] Notes
The precision of the isotope abundances and atomic mass is limited through variations. The given ranges should be applicable to any normal terrestrial material.
Values marked # are not purely derived from experimental data, but at least partly from systematic trends. Spins with weak assignment arguments are enclosed in parentheses.
Uncertainties are given in concise form in parentheses after the corresponding last digits. Uncertainty values denote one standard deviation, except isotopic composition and standard atomic mass from IUPAC which use expanded uncertainties.

[edit] References
Isotope masses from Ame2003 Atomic Mass Evaluation by G. Audi, A.H. Wapstra, C. Thibault, J. Blachot and O. Bersillon in Nuclear Physics A729 (2003).
Isotopic compositions and standard atomic masses from Atomic weights of the elements. Review 2000 (IUPAC Technical Report). Pure Appl. Chem. Vol. 75, No. 6, pp. 683-800, (2003) and Atomic Weights Revised (2005).
Half-life, spin, and isomer data selected from these sources. Editing notes on this article's talk page.
Audi, Bersillon, Blachot, Wapstra. The Nubase2003 evaluation of nuclear and decay properties, Nuc. Phys. A 729, pp. 3-128 (2003).
National Nuclear Data Center, Brookhaven National Laboratory. Information extracted from the NuDat 2.1 database (retrieved Sept. 2005).
David R. Lide (ed.), Norman E. Holden in CRC Handbook of Chemistry and Physics, 85th Edition, online version. CRC Press. Boca Raton, Florida (2005). Section 11, Table of the Isotopes.

[edit] See also
Oxygen isotope ratio cycle
Oxygen

Isotopes of nitrogen Isotopes of oxygen Isotopes of fluorine
Index to isotope pages

Retrieved from "http://en.wikipedia.org/wiki/Isotopes_of_oxygen"
Categories: Oxygen | Isotopes

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article

This page was last modified 07:27, 27 June 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

CAS registry number

2007-08-14T10:52:00.000-07:00

CAS registry number
From Wikipedia, the free encyclopedia
Jump to: navigation, search
CAS registry numbers are unique numerical identifiers for chemical compounds, polymers, biological sequences, mixtures and alloys. They are also referred to as CAS numbers, CAS RNs or CAS #s.

Chemical Abstracts Service (CAS), a division of the American Chemical Society, assigns these identifiers to every chemical that has been described in the literature. The intention is to make database searches more convenient, as chemicals often have many names. Almost all molecule databases today allow searching by CAS number.

As of June 2007, there were 31,745,275 organic and inorganic substances and 59,039,087 sequences in the CAS registry.[1] Around 50,000 new numbers are added each week.

CAS also maintains and sells a database of these chemicals, known as the CAS registry.

Contents [hide]
1 Format
2 Isomers, enzymes, and mixtures
3 Searches
4 Notes
5 See also
6 External links

[edit] Format
A CAS registry number is separated by hyphens into three parts, the first consisting of up to 6 digits, the second consisting of two digits, and the third consisting of a single digit serving as a check digit. The numbers are assigned in increasing order and do not have any inherent meaning. The checksum is calculated by taking the last digit times 1, the next digit times 2, the next digit times 3 etc., adding all these up and computing the sum modulo 10. For example, the CAS number of water is 7732-18-5: the checksum is calculated as (8×1 + 1×2 + 2×3 + 3×4 + 7×5 + 7×6) = 105; 105 mod 10 = 5.

[edit] Isomers, enzymes, and mixtures
Different stereoisomers of a molecule receive different CAS numbers: D-glucose has 50-99-7, L-glucose has 921-60-8, α-D-glucose has 26655-34-5, etc. Occasionally, whole classes of molecules receive a single CAS number: the group of alcohol dehydrogenases has 9031-72-5. An example of a mixture with a CAS number is mustard oil (8007-40-7).

[edit] Searches
When using CAS numbers for database searches, it is useful to include the numbers of closely related compounds. For instance, to search for information about cocaine (CAS 50-36-2), one should consider including cocaine hydrochloride (CAS 53-21-4), since that is the most common form of cocaine when used as a drug.

[edit] Notes
^ CAS Registry Number and Substance Counts

[edit] See also
EC number (Enzyme Commission)
EC# (EINECS and ELINCS)
International Chemical Identifier (InChI)
PubChem
SMILES
UN number
Chemical database

[edit] External links
CAS registry description, by the Chemical Abstracts Service
To find the CAS number of a compound given its name, formula or structure, the following free resources can be used:

PubChem
R&D Chemicals
NIH ChemIDplus
NIST Chemistry WebBook
NCI Database Browser
Chemfinder
European chemical Substances Information System (ESIS) - useful for finding EC#
Retrieved from "http://en.wikipedia.org/wiki/CAS_registry_number"
Categories: Chemical numbering schemes | American Chemical Society

Speed of Sound

2007-08-14T10:50:00.000-07:00

Speed of sound
From Wikipedia, the free encyclopedia
Jump to: navigation, search
This page is about the physical speed of sound waves in a medium. For other uses of the term and related terms, see Speed of sound (disambiguation).
Sound measurements
Sound pressure p
Sound pressure level (SPL)
Particle velocity v
Particle velocity level (SVL)
(Sound velocity level)
Particle displacement ξ
Sound intensity I
Sound intensity level (SIL)
Sound power Pac
Sound power level (SWL)
Sound energy density E
Sound energy flux q
Acoustic impedance Z
Speed of sound c
Sound is a vibration that travels through an elastic medium as a wave. The speed of sound describes how much distance such a wave travels in a given amount of time. The speed varies with the medium employed (for example, sound waves move faster through water than through air), as well as with the properties of the medium, especially temperature. The term is commonly used to refer specifically to the speed of sound in air. At sea level, at a temperature of 21 °C (70 °F) and under normal atmospheric conditions, the speed of sound is 344 m/s (1238 km/h, or 769 mph, or 1128 ft/s or 661.5 kt).

The speed of sound is sometimes used in describing the nature of substances (see the article on sodium).

In conventional use and in scientific literature sound velocity, v, and sound speed, c, are used synonymously and should not be confused with sound particle velocity (also symbolized as v), which is the velocity of the individual particles.

In the Earth's atmosphere, the speed varies with atmospheric conditions; the most important factor is the temperature. Air pressure has almost no effect on sound speed. It has no effect at all in an ideal gas approximation, because pressure and density both contribute to sound velocity equally, and in an ideal gas the two effects cancel out, leaving only the effect of temperature. Sound usually travels more slowly with greater altitude, due to reduced temperature, creating a negative sound speed gradient. In the stratosphere, the speed of sound increases with height due to heating within the ozone layer, producing a positive sound speed gradient.

Humidity has a small, but measurable effect on sound speed. Sound travels slightly (0.1%-0.6%) faster in humid air. The approximate speed of sound in 0% humidity (dry) air, in metres per second (m·s-1), at temperatures near 0 °C, can be calculated from:

where is the temperature in degrees Celsius (°C).

This equation is derived from the first two terms of the Taylor expansion of the following equation:

The value of 331.3 m/s, which represents the 0 °C speed, is probably the most defensible based on theoretical (and some measured) values of the specific heat ratio, γ. Commonly found values for the speed of sound at 0 °C may vary from 331.2 to 331.6 due to the assumptions made when it is calculated. If ideal gas γ is assumed to be 7/5 = 1.4 exactly, the 0 °C speed is calculated (see section below) to be 331.3 m/s, the coefficient used above.

This equation is correct to a wider temperature range, but still depends on the approximation of heat capacity being independent of temperature, and will fail particularly at higher temperatures. It gives good predictions in relatively dry, cold, low pressure conditions, such as the Earth's stratosphere. A derivation of these equations will be given in a later section.

Contents [hide]
1 Basic concept
2 Details
2.1 Speed in solids
2.2 Speed in fluids
2.3 Speed in ideal gases and in air
2.3.1 Temperatures in Celsius
2.4 Tables
3 Effect of frequency and gas composition
4 Mach number
5 Experimental methods
5.1 Single-shot timing methods
5.2 Other methods
6 Gradients
7 References
8 See also
9 External links

[edit] Basic concept
The transmission of sound can be explained using a toy model consisting of an array of balls interconnected by springs. For a real material the balls represent molecules and the springs represent the bonds between them. Sound passes through the model by compressing and expanding the springs, transmitting energy to neighboring balls, which transmit energy to their springs, and so on. The speed of sound through the model depends on the stiffness of the springs (stiffer springs transmit energy more quickly). Effects like dispersion and reflection can also be understood using this model.

In a real material, the stiffness of the springs is called the elastic modulus, and the mass corresponds to the density. All other things being equal, sound will travel more slowly in denser materials, and faster in stiffer ones. For instance, sound will travel faster in iron than uranium, and faster in hydrogen than nitrogen, due to the lower density of the first material of each set. At the same time, sound will travel faster in iron than hydrogen, because the internal bonds in a solid like iron are much stronger than the gaseous bonds between hydrogen molecules. In general, solids will have a higher speed of sound than liquids, and liquids will have a higher speed of sound than gases.

Some textbooks mistakenly state that the speed of sound increases with increasing density. This is usually illustrated by presenting data for three materials, such as air, water and steel. With only these three examples it indeed appears that speed is correlated to density, yet including only a few more examples would show this assumption to be incorrect.

[edit] Details
In general, the speed of sound c is given by

where

C is a coefficient of stiffness
ρ is the density
Thus the speed of sound increases with the stiffness of the material, and decreases with the density. For general equations of state, if classical mechanics is used, the speed of sound c is given by

where differentiation is taken with respect to adiabatic change.

If relativistic effects are important, the speed of sound may be calculated from the relativistic Euler equations.

In a non-dispersive medium sound speed is independent of sound frequency, so the speeds of energy transport and sound propagation are the same. For audible sounds air is a non-dispersive medium. But air does contain a small amount of CO2 which is a dispersive medium, and it introduces dispersion to air at ultrasonic frequencies (> 28 kHz).[citation needed]

In a dispersive medium sound speed is a function of sound frequency. The spatial and temporal distribution of a propagating disturbance will continually change. Each frequency component propagates at its own phase velocity, while the energy of the disturbance propagates at the group velocity. The same phenomenon occurs with light waves -- see optical dispersion for a description.

[edit] Speed in solids
In a solid, there is a non-zero stiffness both for volumetric and shear deformations. Hence, in a solid it is possible to generate sound waves with different velocities dependent on the deformation mode.

In a solid rod (with thickness much smaller than the wavelength) the speed of sound is given by:

where

E is Young's modulus
ρ (rho) is density
Thus in steel the speed of sound is approximately 5100 m·s-1.

In a solid with lateral dimensions much larger than the wavelength, the sound velocity is higher. It is found by replacing Young's modulus with the plane wave modulus, which can be expressed in terms of the Young's modulus and Poisson's ratio as:

[edit] Speed in fluids
In a fluid the only non-zero stiffness is to volumetric deformation (a fluid does not sustain shear forces).

Hence the speed of sound in a fluid is given by

where

K is the bulk modulus of the fluid
The speed of sound in water is of interest to anyone using underwater sound as a tool, whether in a laboratory, a lake or the ocean. Examples are sonar, acoustic communication and acoustical oceanography. See Discovery of Sound in the Sea for other examples of the uses of sound in the ocean (by both man and other animals). In fresh water, sound travels at about 1497 m/s at 25 °C. See Technical Guides - Speed of Sound in Pure Water for an online calculator.

Sound speed as a function of depth at a position north of Hawaii in the Pacific Ocean derived from the 2005 World Ocean Atlas. The SOFAR channel is centered on the minimum in sound speed at ca. 750-m depth.In salt water that is free of air bubbles or suspended sediment, sound travels at about 1500 m/s. The speed of sound in seawater depends on pressure (hence depth), temperature (a change of 1 °C ~ 4 m/s), and salinity (a change of 1‰ ~ 1 m/s), and empirical equations have been derived to accurately calculate sound speed from these variables. Other factors affecting sound speed are minor. For more information see Dushaw et al. (1993).

A simple empirical equation for the speed of sound in sea water with reasonable accuracy for the world's oceans is due to Mackenzie (1981)

c(T, S, z) = a1 + a2T + a3T2 + a4T3 + a5(S - 35) + a6z + a7z2 + a8T(S - 35) + a9Tz3
where T, S, and z are temperature in degrees Celsius, salinity in parts per thousand and depth in metres, respectively. The constants a1, a2, ..., a9 are:

a1 = 1448.96, a2 = 4.591, a3 = -5.304×10-2, a4 = 2.374×10-4, a5 = 1.340, a6 = 1.630×10-2, a7 = 1.675×10-7, a8 = -1.025×10-2, a9 = -7.139×10-13
with check value 1550.744 m/s for T=25 °C, S=35‰, z=1000 m. This equation is accurate to O(0.2 m/s). See Technical Guides - Speed of Sound in Sea-Water for an online calculator.

Other equations for sound speed in sea water have slightly greater accuracy, but are far more complicated, e.g., that by V. A. Del Grosso (1974) and the Chen-Millero-Li Equation (1994).[1] [2]

[edit] Speed in ideal gases and in air
For a gas, K (the bulk modulus in equations above, equivalent to C, the coefficient of stiffness in solids) is approximately given by

thus
Where:

γ is the adiabatic index also known as the isentropic expansion factor. It is the ratio of specific heats of a gas at a constant-pressure to a gas at a constant-volume(Cp / Cv), and arises because a classical sound wave induces an adiabatic compression, in which the heat of the compression does not have enough time to escape the pressure pulse, and thus contributes to the pressure induced by the compression.
p is the pressure.
ρ is the density
Using the ideal gas law to replace p with NRT/V, and replacing ρ with NM/V, the equation for an ideal gas becomes:

where

cideal is the speed of sound in an ideal gas.
R (approximately 8.3145 J·mol-1·K-1) is the molar gas constant.[1]
k is the Boltzmann constant
γ (gamma) is the adiabatic index (sometimes assumed 7/5 = 1.400 for diatomic molecules from kinetic theory, assuming from quantum theory a temperature range at which thermal energy is fully partitioned into rotation (rotations are fully excited), but none into vibrational modes. Gamma is actually experimentally measured over a range from 1.3991 to 1.403 at 0 degrees Celsius, for air. Gamma is assumed from kinetic theory to be exactly 5/3 = 1.6667 for monoatomic molecules such as noble gases).
T is the absolute temperature in kelvins.
M is the molar mass in kilograms per mole. The mean molar mass for dry air is about .0289645 kg/mole.
m is the mass of a single molecule in kilograms.
This equation applies only when the sound wave is a small perturbation on the ambient condition, and the certain other noted conditions are fulfilled, as noted below. Calculated values for cair have been found to vary slightly from experimentally determined values.:[3]

Newton famously considered the speed of sound before most of the development of thermodynamics and so incorrectly used isothermal calculations instead of adiabatic. His result was missing the factor of γ but was otherwise correct.

Numerical substitution of the above values gives the ideal gas approximation of sound velocity for gases, which is accurate at relatively low gas pressures and densities (for air, this includes standard Earth sea-level conditions). Also, for diatomic gases the use of requires that the gas exist in a temperature range high enough that rotational heat capacity is fully excited (i.e., molecular rotation is fully used as a heat energy "partition" or reservoir); but at the same time the temperature must be low enough that molecular vibrational modes contribute no heat capacity (i.e., insigificant heat goes into vibration, as all vibrational quantum modes above the minimum-energy-mode, have energies too high to be populated by a significant number of molecules at this temperature). For air, these conditions are fulfilled at room temperature, and also temperatures considerably below room temperature (see tables below). See the section on gases in heat capacity for a more complete discussion of this phenomenon.

[edit] Temperatures in Celsius
If temperatures in degrees Celsius(°C) are to be used to calculate air speed in the region near 273 kelvins, then Celsius temperature may be used.

For dry air, where (theta) is the temperature in degrees Celsius(°C).

Making the following numerical substitutions: , where is the molar gas constant, , and using the ideal diatomic gas value of

Then:

Using the first two terms of the Taylor expansion:

The derivation includes the two approximate equations which were given in the introduction. For Celsius temperatures which are negative, the second term of the equation right hand side, is negative.

[edit] Tables
In the standard atmosphere:

T0 is 273.15 K (= 0 °C = 32 °F), giving a theoretical value of 331.3 m·s-1 (= 1086.9 ft/s = 1193 km·h-1 = 741.1 mph = 644.0 knots). Values ranging from 331.3-331.6 may be found in reference literature, however.
T20 is 293.15 K (= 20 °C = 68 °F), giving a value of 343.2 m·s-1 (= 1126.0 ft/s = 1236 km·h-1 = 767.8 mph = 667.2 knots).
T25 is 298.15 K (= 25 °C = 77 °F), giving a value of 346.1 m·s-1 (= 1135.6 ft/s = 1246 km·h-1 = 774.3 mph = 672.8 knots).

In fact, assuming an ideal gas, the speed of sound c depends on temperature only, not on the pressure or density (since these change in lockstep for a given temperature and cancel out). Air is almost an ideal gas. The temperature of the air varies with altitude, giving the following variations in the speed of sound using the standard atmosphere - actual conditions may vary.

Effect of temperature
in °C c in m·s-1 ρ in kg·m-3 Z in N·s·m-3
−10 325.2 1.342 436.1
−5 328.3 1.317 432.0
0 331.3 1.292 428.4
+5 334.3 1.269 424.3
+10 337.3 1.247 420.6
+15 340.3 1.225 416.8
+20 343.2 1.204 413.2
+25 346.1 1.184 409.8
+30 349.0 1.165 406.3

is the temperature in °C
c is the speed of sound in m·s-1
ρ is the density in kg·m-3
Z is the characteristic acoustic impedance in N·s·m-3 (Z=ρ·c)
Given normal atmospheric conditions, the temperature, and thus speed of sound, varies with altitude:

Altitude Temperature m·s-1 km·h-1 mph knots
Sea level 15 °C (59 °F) 340 1225 761 661
11 000 m−20 000 m
(Cruising altitude of commercial jets,
and first supersonic flight) -57 °C (-70 °F) 295 1062 660 573
29 000 m (Flight of X-43A) -48 °C (-53 °F) 301 1083 673 585

[edit] Effect of frequency and gas composition
The medium in which a sound wave is travelling does not always respond adiabatically, and as a result the speed of sound can vary with frequency.[4]

The limitations of the concept of speed of sound due to extreme attenuation are also of concern. The attenuation which exists at sea level for high frequencies applies to successively lower frequencies as atmospheric pressure decreases, or as the mean free path increases. For this reason, the concept of speed of sound (except for frequencies approaching zero) progressively loses its range of applicability at high altitudes.:[3]

The molecular composition of the gas contributes both as the mass (M) of the molecules, and their heat capacities, and so both have an influence on speed of sound. In general, at the same molecular mass, monatomic gases have slightly higher sound speeds (over 9% higher) due to the fact that they have a higher γ (5/3 = 1.67) than diatomics do (7/5 = 1.4). Thus, at the same molecular mass, the sound speed of a monatomic gas goes up by a factor of

= 1.09

This gives the 9 % difference, and would be a typical ratio for sound speeds at room temperature in helium vs. deuterium, each with a molecular weight of 4. Sound travels faster in helium than deuterium because adiabatic compression heats helium more, since the helium molecules can store heat energy from compression only in translation, but not rotation. Thus helium molecules (monatomic molecules) travel faster in a sound wave and transmit sound faster. (Sound generally travels at about 70% of the mean molecular velocity in gases).

Note that in this example we have assumed that temperature is low enough that heat capacities are not influenced by molecular vibration (see heat capacity). However, vibrational modes simply cause gammas which decrease toward 1, since vibration modes in a polyatomic gas gives the gas additional ways to store heat which do not affect temperature, and thus do not affect molecular velocity and sound velocity. Thus, the effect of higher temperatures and vibrational heat capacity acts to increase the difference between sound speed in monatomic vs. polyatomic molecules, with the speed remaining greater in monatomics.

[edit] Mach number
Main article: Mach number.
Mach number, a useful quantity in aerodynamics, is the ratio of an object's speed to the speed of sound in the medium through which it is passing (again, usually air). At altitude, for reasons explained, Mach number is a function of temperature.

Aircraft flight instruments, however, operate using pressure differential to compute Mach number; not temperature. The assumption is that a particular pressure represents a particular altitude and, therefore, a standard temperature. Aircraft flight instruments need to operate this way because the impact pressure sensed by a Pitot tube is dependent on altitude as well as speed.

Assuming air to be an ideal gas, the formula to compute Mach number in a subsonic compressible flow is derived from the Bernoulli equation for M<1:[5]

where

M is Mach number
qc is impact pressure and
P is static pressure.
The formula to compute Mach number in a supersonic compressible flow is derived from the Rayleigh Supersonic Pitot equation:

where

M is Mach number
qc is impact pressure measured behind a normal shock
P is static pressure.
As can be seen, M appears on both sides of the equation. The easiest method to solve the supersonic M calculation is to enter both the subsonic and supersonic equations into a computer spread sheet such as Microsoft Excel, OpenOffice.org Calc, or some equivalent program. First determine if M is indeed greater than 1.0 by calculating M from the subsonic equation. If M is greater than 1.0 at that point, then use the value of M from the subsonic equation as the initial condition in the supersonic equation. Then perform a simple iteration of the supersonic equation, each time using the last computed value of M, until M converges to a value--usually in just a few iterations.[5]

[edit] Experimental methods
A range of different methods exist for the measurement of sound in air.

[edit] Single-shot timing methods
The simplest concept is the measurement made using two microphones and a fast recording device such as a digital storage scope. This method uses the following idea.

If a sound source and two microphones are arranged in a straight line, with the sound source at one end, then the following can be measured:

1. The distance between the microphones (x), called microphone basis. 2. The time of arrival between the signals (delay) reaching the different microphones (t)

Then v = x / t

An older method is to create a sound at one end of a field with an object that can be seen to move when it creates the sound. When the observer sees the sound-creating device act they start a stopwatch and when the observer hears the sound they stop their stopwatch. Again using v = x / t you can calculate the speed of sound. A separation of at least 200 m between the two experimental parties is required for good results with this method.

[edit] Other methods
In these methods the time measurement has been replaced by a measurement of the inverse of time (frequency).

Kundt's tube is an example of an experiment which can be used to measure the speed of sound in a small volume. It has the advantage of being able to measure the speed of sound in any gas. This method uses a powder to make the nodes and antinodes visible to the human eye. This is an example of a compact experimental setup.

A tuning fork can be held near the mouth of a long pipe which is dipping into a barrel of water. In this system it is the case that the pipe can be brought to resonance if the length of the air column in the pipe is equal to ({1+2n}λ/4) where n is an integer. As the antinodal point for the pipe at the open end is slightly outside the mouth of the pipe it is best to find two or more points of resonance and then measure half a wavelength between these.

Here it is the case that v = fλ

[edit] Gradients
Normally sound waves spread out in an inverse square law, and rapidly dissipate.

However, in the ocean (the 'deep sound channel' or SOFAR channel), variations in the speed of sound create a layer where the speed of sound is at a minimum and any sound waves generated spread out in a substantially flat layer- refraction keeps the sound constrained to spread out in an inverse law, which carries the sound very much further.

A similar effect occurs in the atmosphere and Project_Mogul was an attempt to detect the sound of nuclear weapons using this principle, and was successfully able to detect a nuclear explosion at a considerable distance.

[edit] References
^ Article Abstract
^ dushaw-jasa-93
^ a b U.S. Standard Atmosphere, 1976, U.S. Government Printing Office, Washington, D.C., 1976.
^ A B Wood, A Textbook of Sound (Bell, London, 1946)
^ a b Olson, Wayne M. (2002). "AFFTC-TIH-99-02, Aircraft Performance Flight Testing." (PDF). Air Force Flight Test Center, Edwards AFB, CA, United States Air Force.
Del Grosso, V. A., 1974. New equation for the speed of sound in natural waters (with comparisons to other equations). J. Acoust. Soc. Am., 56, pp. 1084-1091.
Dushaw, B. D., P. F. Worcester, B. D. Cornuelle, and B. M. Howe, 1993. On equations for the speed of sound in seawater. J. Acoust. Soc. Am., 93, pp. 255-275.
Mackenzie, K. V., 1981. Discussion of sea water sound-speed determinations. J. Acoust. Soc. Am., 70, pp. 801-806.
Applied Physics Laboratory - University of Washington, 1994. APL-UW High-Frequency Ocean Environmental Acoustic Models Handbook. APL-UW TR 9407, pp. I1-I2.

[edit] See also
Sound barrier
SOFAR channel
Underwater acoustics

[edit] External links
Calculation: Speed of sound in air and the temperature
Speed of sound - temperature matters, not air pressure
Properties Of The U.S. Standard Atmosphere 1976
How to measure the speed of sound in a laboratory
Speed of sound in water and water steam as function of pressure & temperature
Teaching resource for 14-16yrs on sound including speed of sound
Technical Guides - Speed of Sound in Pure Water
Technical Guides - Speed of Sound in Sea-Water
Retrieved from "http://en.wikipedia.org/wiki/Speed_of_sound"
Categories: Sound measurements | All articles with unsourced statements | Articles with unsourced statements since February 2007 | Chemical properties | Fluid dynamics | Units of velocity

Paramagnetism

2007-08-14T10:43:00.000-07:00

Paramagnetism
From Wikipedia, the free encyclopedia
Jump to: navigation, search

Simple Illustration of a paramagnetic probe made up from miniature magnets.Paramagnetism is a form of magnetism which occurs only in the presence of an externally applied magnetic field. Paramagnetic materials are attracted to magnetic fields, hence have a relative magnetic permeability greater than one (or, equivalently, a positive magnetic susceptibility). However, unlike ferromagnets which are also attracted to magnetic fields, paramagnets do not retain any magnetization in the absence of an externally applied magnetic field.

Contents [hide]
1 Introduction
2 Curie's law
3 Paramagnetic materials
3.1 Elements
3.2 Compounds
4 See also
5 References
6 External links

[edit] Introduction
Constituent atoms or molecules of paramagnetic materials have permanent magnetic moments (dipoles), even in the absence of an applied field. This generally occurs due to the presence of unpaired electrons in the atomic/molecular electron orbitals. In pure paramagnetism, the dipoles do not interact with one another and are randomly oriented in the absence of an external field due to thermal agitation, resulting in zero net magnetic moment. When a magnetic field is applied, the dipoles will tend to align with the applied field, resulting in a net magnetic moment in the direction of the applied field. In the classical description, this alignment can be understood to occur due to a torque being provided on the magnetic moments by an applied field, which tries to align the dipoles parallel to the applied field. However, the truer origins of the alignment can only be understood via the quantum-mechanical properties of spin and angular momentum.

If there is sufficient energy exchange between neighbouring dipoles they will interact, and may spontaneously align or anti-align and form magnetic domains, resulting in ferromagnetism (permanent magnets) or antiferromagnetism, respectively. Paramagnetic behaviour can also be observed in ferromagnetic materials that are above their Curie temperature, and in antiferromagnets above their Néel temperature.

In general paramagnetic effects are quite small: the magnetic susceptibility is of the order of 10−3 to 10−5 for most paramagnets, but may be as high as 10-1 for synthetic paramagnets such as ferrofluids.

[edit] Curie's law
For low levels of magnetisation, the magnetisation of paramagnets is approximated by Curie's law:

where

M is the resulting magnetization
B is the magnetic flux density of the applied field, measured in teslas
T is absolute temperature, measured in kelvins
C is a material-specific Curie constant
This law indicates that the susceptibiliy of paramagnetic materials is inversely proportional to their temperature. However, Curie's law is only valid under conditions of low magnetisation, since it does not consider the saturation of magnetisation that occurs when the atomic dipoles are all aligned in parallel (after everything is aligned, increasing the external field will not increase the total magnetisation since there can be no further alignment).

[edit] Paramagnetic materials

[edit] Elements
Elements can be paramagnetic if they have unpaired electrons.

The following are some examples of paramagnetic elements:

Aluminium Al [13] (metal) — Al is the preferred paramagnetic material in theoretical designs of lunar mass driver applications using regolith as an ore.
Barium Ba [56] (metal)
Calcium Ca [20] (metal) [Ar]4s2 — diamagnetic
Oxygen. O [8] (non-metal)
Platinum Pt [78] (metal)
Sodium Na [11] (metal)
Strontium Sr [38] (metal)
Uranium U [92] (metal)
Magnesium Mg [12] (metal) 1s2 2s2 2p6 3s2 — diamagnetic
Technetium Tc [43] (artificial)
Dysprosium Dy [66] (metal) — ferromagnetic

[edit] Compounds
Many salts of the d and f transitional metal group show paramagnetic behaviour.

Examples are:

Copper sulphate
Dysprosium oxide
Ferric chloride
Ferric oxide
Holmium oxide
Manganese chloride
Some simple molecules contain unpaired electrons and are thus paramagnetic. The most common is the diatomic oxygen molecule.

[edit] See also
Pierre Curie
Ferromagnetism

Magnetic states
diamagnetism – superdiamagnetism – paramagnetism – superparamagnetism – ferromagnetism – antiferromagnetism – ferrimagnetism – metamagnetism – spin glass

[edit] References
Charles Kittel, Introduction to Solid State Physics (Wiley: New York, 1996).
Neil W. Ashcroft and N. David Mermin, Solid State Physics (Harcourt: Orlando, 1976).
John David Jackson, Classical Electrodynamics (Wiley: New York, 1999).

[edit] External links
Classification of Magnetic Materials by Applied Alloy Chemistry Group at University of Birmingham.
Retrieved from "http://en.wikipedia.org/wiki/Paramagnetism"
Categories: Electric and magnetic fields in matter | Magnetism | Fundamental physics concepts

Magnetism

2007-08-14T10:41:00.000-07:00

Magnetism
From Wikipedia, the free encyclopedia
Jump to: navigation, search
For other senses of this word, see magnetism (disambiguation) or Electromagnetic induction.

Magnetic lines of force of a bar magnet shown by iron filings on paperIn physics, magnetism is one of the phenomena by which materials exert attractive or repulsive forces on other materials. Some well known materials that exhibit easily detectable magnetic properties are nickel, iron and their alloys; however, all materials are influenced to greater or lesser degree by the presence of a magnetic field.

Contents [hide]
1 History
2 Physics of magnetism
2.1 Magnetism, electricity, and special relativity
2.2 Magnetic fields and forces
2.3 Magnetic dipoles
2.4 Magnetic monopoles
2.5 Atomic magnetic dipoles
3 Types of magnets
3.1 Electromagnets
3.2 Permanent and temporary magnets
3.2.1 Magnetic metallic elements
3.2.2 Composites
3.2.2.1 Ceramic or ferrite
3.2.2.2 Alnico
3.2.2.3 Injection molded
3.2.2.4 Flexible
3.2.3 Rare earth magnets
3.2.3.1 Samarium-cobalt
3.2.3.2 Neodymium-iron-boron (NIB)
3.2.4 Single-molecule magnets (SMMs) and single-chain magnets (SCMs)
3.2.5 Nano-structured magnets
4 Units of electromagnetism
4.1 SI units related to magnetism
4.2 Other units
5 See also
6 References
7 External links

[edit] History
Aristotle attributes the first scientific theory on magnetism to Thales, who lived from about 625 BC to about 545 BC. [1] In China, the earliest literary reference to magnetism lies in a 4th century BC book called Book of the Devil Valley Master (鬼谷子): "The lodestone makes iron come or it attracts it."[1] The earliest mention of the attraction of a needle appears in a work composed between 20 and 100 AD (Louen-heng): "A lodestone attracts a needle."[2] The ancient Chinese scientist Shen Kuo (1031-1095) was the first person to write of the magnetic needle compass and improved the accuracy of navigation by employing the astronomical concept of true north (Dream Pool Essays, 1088 AD), and by the 12th century the Chinese were known to use the lodestone compass for navigation. Alexander Neckham, by 1187, was the first in Europe to describe the compass and its use for navigation.

An understanding of the relationship between electricity and magnetism began in 1819 with work by Hans Christian Oersted, a professor at the University of Copenhagen, discovered more or less by accident that an electric current could influence a compass needle. This landmark experiment is known as Oersted's Experiment. Several other experiments followed, with André-Marie Ampère, Carl Friedrich Gauss, Michael Faraday, and others finding further links between magnetism and electricity. James Clerk Maxwell synthesized and expanded these insights into Maxwell's equations, unifying electricity, magnetism, and optics into the field of electromagnetism. In 1905, Einstein used these laws in motivating his theory of special relativity[3], in the process showing that electricity and magnetism are fundamentally interlinked and inseparable.

Electromagnetism has continued to develop into the twentieth century, being incorporated into the more fundamental theories of gauge theory, quantum electrodynamics, electroweak theory, and finally the standard model.

[edit] Physics of magnetism

[edit] Magnetism, electricity, and special relativity
As a consequence of Einstein's theory of special relativity, electricity and magnetism are understood to be fundamentally interlinked. Both magnetism without electricity, and electricity without magnetism, are inconsistent with special relativity, due to such effects as length contraction, time dilation, and the fact that the magnetic force is velocity-dependent. However, when both electricity and magnetism are taken into account the resulting theory (electromagnetism) is fully consistent with special relativity[4][5]. In particular, a phenomenon that appears purely electric to one observer may be purely magnetic to another, or more generally the relative contributions of electricity and magnetism are dependent on the frame of reference. Thus, special relativity "mixes" electricity and magnetism into a single, inseparable phenomenon called electromagnetism (analogously to how special relativity "mixes" space and time into spacetime).

[edit] Magnetic fields and forces
Main article: magnetic field
The phenomenon of magnetism is "mediated" by the magnetic field -- i.e., an electric current or magnetic dipole creates a magnetic field, and that field, in turn, imparts magnetic forces on other particles that are in the fields.

To an excellent approximation (but ignoring some quantum effects---see quantum electrodynamics), Maxwell's equations (which simplify to the Biot-Savart law in the case of steady currents) describe the origin and behavior of the fields that govern these forces. Therefore magnetism is seen whenever electrically charged particles are in motion---for example, from movement of electrons in an electric current, or in certain cases from the orbital motion of electrons around an atom's nucleus. They also arise from "intrinsic" magnetic dipoles arising from quantum effects, i.e. from quantum-mechanical spin.

The same situations which create magnetic fields (charge moving in a current or in an atom, and intrinsic magnetic dipoles) are also the situations in which a magnetic field has an effect, creating a force. Following is the formula for moving charge; for the forces on an intrinsic dipole, see magnetic dipole.

When a charged particle moves through a magnetic field B, it feels a force F given by the cross product:

where is the electric charge of the particle, is the velocity vector of the particle, and is the magnetic field. Because this is a cross product, the force is perpendicular to both the motion of the particle and the magnetic field. It follows that the magnetic force does no work on the particle; it may change the direction of the particle's movement, but it cannot cause it to speed up or slow down. The magnitude of the force is

where is the angle between the and vectors.

One tool for determining the direction of the velocity vector of a moving charge, the magnetic field, and the force exerted is labeling the index finger "V", the middle finger "B", and the thumb "F" with your right hand. When making a gun-like configuration (with the middle finger crossing under the index finger), the fingers represent the velocity vector, magnetic field vector, and force vector, respectively. See also right hand rule.

[edit] Magnetic dipoles
Main article: magnetic dipole
A very common type of magnetic field seen in nature is a dipoles, having a "South pole" and a "North pole"; terms dating back to the use of magnets as compasses, interacting with the Earth's magnetic field to indicate North and South on the globe. Since opposite ends of magnets are attracted, the 'north' magnetic pole of the earth must be magnetically 'south'.

A magnetic field contains energy, and physical systems stabilize into the configuration with the lowest energy. Therefore, when placed in a magnetic field, a magnetic dipole tends to align itself in opposed polarity to that field, thereby canceling the net field strength as much as possible and lowering the energy stored in that field to a minimum. For instance, two identical bar magnets placed side-to-side normally line up North to South, resulting in a much smaller net magnetic field, and resist any attempts to reorient them to point in the same direction. The energy required to reorient them in that configuration is then stored in the resulting magnetic field, which is double the strength of the field of each individual magnet. (This is, of course, why a magnet used as a compass interacts with the Earth's magnetic field to indicate North and South).

An alternative, equivalent formulation, which is often easier to apply but perhaps offers less insight, is that a magnetic dipole in a magnetic field experiences a torque and a force which can be expressed in terms of the field and the strength of the dipole (i.e., its magnetic dipole moment). For these equations, see magnetic dipole.

[edit] Magnetic monopoles
Main article: Magnetic monopole
Regular bar magnets have two "poles": north and south. Like poles will repel each other, and unlike poles will attract. These poles cannot be separated: when a bar magnet is cut in half across the axis joining those poles, each of the resulting pieces are normal (albeit smaller) bar magnets, each with its own north pole and south pole.

A monopole (if they exist) would be an isolated north pole, not attached to a south pole, or vice versa. It would have "magnetic charge" analogous to electric charge. A bar magnet cannot be broken into monopoles because its magnetism (see ferromagnetism is ultimately due to circulating current (electrons orbiting nuclei) and intrinsic spin dipoles, both of which are at root dipolar in nature. Indeed, despite systematic searches since 1931, as of 2006, magnetic monopoles have never been observed, and could very well not exist.[6]

Nevertheless, some theoretical physics models predict the existence of magnetic monopoles. Paul Dirac observed in 1931 that, because electricity and magnetism show a certain symmetry, just as quantum theory predicts that individual positive or negative electric charges can be observed without the opposing charge, isolated South or North magnetic poles should be observable. Using quantum theory Dirac showed that if magnetic monopoles exist, then one could explain the quantization of electric charge---that is, why the observed elementary particles carry charges that are multiples of the charge of the electron. (Since then, other explanations of the quantization of electric charge, not involving monopoles, have been formulated; see spontaneous symmetry breaking and electroweak theory.)

Certain grand unified theories predict the existence of monopoles which, unlike elementary particles, are solitons (localized energy packets). Using these models to estimate the number of monopoles created in the big bang, the initial results that contradicted cosmological observations---the monopoles would have been so plentiful and massive that they would have long since halted the expansion of the universe. However, the idea of inflation (for which this problem served as a partial motivation) was successful in solving this problem, creating models in which monopoles existed but were rare enough to be consistent with current observations.[7]

[edit] Atomic magnetic dipoles
The physical cause of the magnetism of objects, as distinct from electrical currents, is the atomic magnetic dipole. Magnetic dipoles, or magnetic moments, result on the atomic scale from the two kinds of movement of electrons. The first is the orbital motion of the electron around the nucleus; this motion can be considered as a current loop, resulting in an orbital dipole magnetic moment. The second, much stronger, source of electronic magnetic moment is due to a quantum mechanical property called the spin dipole magnetic moment (although current quantum mechanical theory states that electrons neither physically spin, nor orbit the nucleus).

Dipole moment of a bar magnet.The overall magnetic moment of the atom is the net sum of all of the magnetic moments of the individual electrons. Because of the tendency of magnetic dipoles to oppose each other to reduce the net energy, in an atom the opposing magnetic moments of some pairs of electrons cancel each other, both in orbital motion and in spin magnetic moments. Thus, in the case of an atom with a completely filled electron shell or subshell, the magnetic moments normally completely cancel each other out and only atoms with partially-filled electron shells have a magnetic moment, whose strength depends on the number of unpaired electrons.

The differences in configuration of the electrons in various elements thus determine the nature and magnitude of the atomic magnetic moments, which in turn determine the differing magnetic properties of various materials. Several forms of magnetic behavior have been observed in different materials, including:

Diamagnetism
Paramagnetism
Molecular magnet
Ferromagnetism
Antiferromagnetism
Ferrimagnetism
Metamagnetism
Spin glass
Superparamagnetism
Magnetars, stars with extremely powerful magnetic fields, are also known to exist.

[edit] Types of magnets

[edit] Electromagnets
Since all magnetism is caused by moving charges, all magnets are in fact electromagnets. However, we usually refer to magnets made from electrical wire wound around a magnetic material, such as iron as electromagnets. This form of magnet is useful in cases where a magnet must be switched on or off; for instance, large cranes to lift junked automobiles.

For the case of electric current moving through a wire, the resulting field is directed according to the "right hand rule." If the right hand is used as a model, and the thumb of the right hand points along the wire from positive towards the negative side ("conventional current", the reverse of the direction of actual movement of electrons), then the magnetic field will wrap around the wire in the direction indicated by the fingers of the right hand. As can be seen geometrically, if a loop or helix of wire is formed such that the current is traveling in a circle, then all of the field lines in the center of the loop are directed in the same direction, resulting in a magnetic dipole whose strength depends on the current around the loop, or the current in the helix multiplied by the number of turns of wire. In the case of such a loop, if the fingers of the right hand are directed in the direction of conventional current flow (i.e., positive to negative, the opposite direction to the actual flow of electrons), the thumb will point in the direction corresponding to the North pole of the dipole.

[edit] Permanent and temporary magnets
A permanent magnet retains its magnetism without an external magnetic field whereas a temporary magnet is only magnetic while within another magnetic field. Inducing magnetism in steel results in a permanent magnet but iron loses its magnetism when the inducing field is withdrawn. A temporary magnet such as iron is thus a good material for electromagnets. Magnets are made by stroking with another magnet, tapping while fixed in a magnetic field or placing inside a solenoid coil supplied with a direct current. A permanent magnet may be de-magnetised by subjecting it to heating or sharp blows or placing it inside a solenoid supplied with a reducing alternating current.

[edit] Magnetic metallic elements
Many materials have unpaired electron spins, and the majority of these materials are paramagnetic. When the spins interact with each other in such a way that the spins align spontaneously, the materials are called ferromagnetic (what is often loosely termed as "magnetic"). Due to the way their regular crystalline atomic structure causes their spins to interact, some metals are (ferro)magnetic when found in their natural states, as ores. These include iron ore (magnetite or lodestone), cobalt and nickel, as well the rare earth metals gadolinium and dysprosium (when at a very low temperature). Such naturally occurring (ferro)magnets were used in the first experiments with magnetism. Technology has since expanded the availability of magnetic materials to include various manmade products, all based, however, on naturally magnetic elements.

[edit] Composites

[edit] Ceramic or ferrite
Ceramic, or ferrite, magnets are made of a sintered composite of powdered iron oxide and barium/strontium carbonate ceramic. Due to the low cost of the materials and manufacturing methods, inexpensive magnets (or nonmagnetized ferromagnetic cores, for use in electronic component such as radio antennas, for example) of various shapes can be easily mass produced. The resulting magnets are noncorroding, but brittle and must be treated like other ceramics.

[edit] Alnico
Alnico magnets are made by casting or sintering a combination of aluminium, nickel and cobalt with iron and small amounts of other elements added to enhance the properties of the magnet. Sintering offers superior mechanical characteristics, whereas casting delivers higher magnetic fields and allows for the design of intricate shapes. Alnico magnets resist corrosion and have physical properties more forgiving than ferrite, but not quite as desirable as a metal.

[edit] Injection molded
Injection molded magnets are a composite of various types of resin and magnetic powders, allowing parts of complex shapes to be manufactured by injection molding. The physical and magnetic properties of the product depend on the raw materials, but are generally lower in magnetic strength and resemble plastics in their physical properties.

[edit] Flexible
Flexible magnets are similar to injection molded magnets, using a flexible resin or binder such as vinyl, and produced in flat strips or sheets. These magnets are lower in magnetic strength but can be very flexible, depending on the binder used.

[edit] Rare earth magnets
Main article: Rare-earth magnet
'Rare earth' (lanthanoid) elements have a partially occupied f electron shell (which can accommodate up to 14 electrons.) The spin of these electrons can be aligned, resulting in very strong magnetic fields, and therefore these elements are used in compact high-strength magnets where their higher price is not a concern.

[edit] Samarium-cobalt
Samarium-cobalt magnets are highly resistant to oxidation, with higher magnetic strength and temperature resistance than alnico or ceramic materials. Sintered samarium-cobalt magnets are brittle and prone to chipping and cracking and may fracture when subjected to thermal shock.

[edit] Neodymium-iron-boron (NIB)
Neodymium magnets, more formally referred to as neodymium-iron-boron (NdFeB) magnets, have the highest magnetic field strength, but are inferior to samarium cobalt in resistance to oxidation and temperature. This type of magnet has traditionally been expensive, due to both the cost of raw materials and licensing of the patents involved. This high cost limited their use to applications where such high strengths from a compact magnet are critical. Use of protective surface treatments such as gold, nickel, zinc and tin plating and epoxy resin coating can provide corrosion protection where required. Beginning in the 1980s, NIB magnets have increasingly become less expensive and more popular in other applications such as controversial children's magnetic building toys. Even tiny neodymium magnets are very powerful and have important safety considerations.[8]

[edit] Single-molecule magnets (SMMs) and single-chain magnets (SCMs)
In the 1990s it was discovered that certain molecules containing paramagnetic metal ions are capable of storing a magnetic moment at very low temperatures. These are very different from conventional magnets that store information at a "domain" level and theoretically could provide a far denser storage medium than conventional magnets. In this direction research on monolayers of SMMs is currently under way. Very briefly, the two main attributes of an SMM are:

a large ground state spin value (S), which is provided by ferromagnetic or ferrimagnetic coupling between the paramagnetic metal centres.
a negative value of the anisotropy of the zero field splitting (D)
Most SMM's contain manganese, but can also be found with vanadium, iron, nickel and cobalt clusters. More recently it has been found that some chain systems can also display a magnetization which persists for long times at relatively higher temperatures. These systems have been called single-chain magnets.

[edit] Nano-structured magnets
Some nano-structured materials exhibit energy waves called magnons that coalesce into a common ground state in the manner of a Bose-Einstein condensate.

See results from NIST published April 2005,[9] or[10]

[edit] Units of electromagnetism

[edit] SI units related to magnetism
editSI electromagnetism units
Symbol [citation needed] Name of Quantity Derived Units Unit Base Units
I Magnitude of current ampere (SI base unit) A A = W/V = C/s
q Electric charge, Quantity of electricity coulomb C A·s
V Potential difference or Electromotive force volt V J/C = kg·m2·s−3·A−1
R, Z, X Resistance, Impedance, Reactance ohm Ω V/A = kg·m2·s−3·A−2
ρ Resistivity ohm metre Ω·m kg·m3·s−3·A−2
P Power, Electrical watt W V·A = kg·m2·s−3
C Capacitance farad F C/V = kg−1·m−2·A2·s4
Elastance reciprocal farad F−1 V/C = kg·m2·A−2·s−4
ε Permittivity farad per metre F/m kg−1·m−3·A2·s4
χe Electric susceptibility (dimensionless) - -
G, Y, B Conductance, Admittance, Susceptance siemens S Ω−1 = kg−1·m−2·s3·A2
σ Conductivity siemens per metre S/m kg−1·m−3·s3·A2
B Magnetic flux density, Magnetic induction tesla T Wb/m2 = kg·s−2·A−1 = N·A−1·m−1
Φm Magnetic flux weber Wb V·s = kg·m2·s−2·A−1
H Magnetic field strength,Magnetic field intensity ampere per metre A/m A·m−1
Reluctance ampere-turn per weber A/Wb kg−1·m−2·s2·A2
L Inductance henry H Wb/A = V·s/A = kg·m2·s−2·A−2
μ Permeability henry per metre H/m kg·m·s−2·A−2
χm Magnetic susceptibility (dimensionless)
Π and Π * Electric and Magnetic hertzian vector potentials n/a n/a

[edit] Other units
gauss-The gauss, abbreviated as G, is the cgs unit of magnetic flux density or magnetic induction (B).
oersted-The oersted is the CGS unit of magnetic field strength.
maxwell-is the CGS unit for the magnetic flux.
μo -common symbol for the permeability of free space (4πx10-7 N/(ampere-turn)2).

[edit] See also
Wikibooks has more about this subject:
School science how-toElectrostatics
Magnetostatics
Electromagnetism
Plastic magnet
Magnet
Magnetic field
Magnetic bearing
Magnet therapy
Magnetic circuit
Michael Faraday
Micromagnetism
James Clerk Maxwell
Coercivity
Spin wave
Spontaneous magnetization
Sensor

Magnetic states
diamagnetism – superdiamagnetism – paramagnetism – superparamagnetism – ferromagnetism – antiferromagnetism – ferrimagnetism – metamagnetism – spin glass

[edit] References
Griffiths, David J. (1998). Introduction to Electrodynamics (3rd ed.). Prentice Hall. ISBN 0-13-805326-X.
Tipler, Paul (2004). Physics for Scientists and Engineers: Electricity, Magnetism, Light, and Elementary Modern Physics (5th ed.). W. H. Freeman. ISBN 0-7167-0810-8.
Furlani, Edward P. (2001). Permanent Magnet and Electromechanical Devices: Materials, Analysis and Applications. Academic Press. ISBN 0-12-269951-3.
^ Li Shu-hua, “Origine de la Boussole 11. Aimant et Boussole,” Isis, Vol. 45, No. 2. (Jul., 1954), p.175
^ Li Shu-hua, “Origine de la Boussole 11. Aimant et Boussole,” Isis, Vol. 45, No. 2. (Jul., 1954), p.176
^ A. Einstein: "On the Electrodynamics of Moving Bodies", June 30, 1905. http://www.fourmilab.ch/etexts/einstein/specrel/www/.
^ A. Einstein: "On the Electrodynamics of Moving Bodies", June 30, 1905. http://www.fourmilab.ch/etexts/einstein/specrel/www/.
^ Griffiths, David J. (1998). Introduction to Electrodynamics, 3rd ed., Prentice Hall. ISBN 0-13-805326-X. , chapter 12
^ Milton mentions some inconclusive events (p.60) and still concludes that "no evidence at all of magnetic monopoles has survived" (p.3). Milton, Kimball A. (June 2006). "Theoretical and experimental status of magnetic monopoles". Reports on Progress in Physics 69 (6): 1637-1711. DOI:10.1088/0034-4885/69/6/R02. .
^ Guth, Alan (1997). The Inflationary Universe: The Quest for a New Theory of Cosmic Origins. Perseus. ISBN 0-201-32840-2. .
^ Magnet Man, Magnet Basics - Safety Considerations accessed 6 October 2006.
^ Nanomagnets Bend The Rules. Retrieved on November 14, 2005.
^ Nanomagnets bend the rules. Retrieved on November 14, 2005.

[edit] External links
Look up Magnetism in
Wiktionary, the free dictionary.Electromagnetism - a chapter from an online textbook
Magnetic Force and Field on Project PHYSNET
On the Magnet, 1600 First scientific book on magnetism by the father of electrical engineering. Full English text, full text search.

Electromagnetism
Electricity · Magnetism
Electrostatics
Electric charge
Coulomb's law
Electric field
Gauss's law
Electric potential
Electric dipole moment
Magnetostatics
Ampère's Circuital law
Magnetic field
Magnetic flux
Biot-Savart law
Magnetic dipole moment
Electrodynamics
Electrical current
Lorentz force law
Electromotive force
(EM) Electromagnetic induction
Faraday-Lenz law
Displacement current
Maxwell's equations
(EMF) Electromagnetic field
(EM) Electromagnetic radiation
Electrical Network
Electrical conduction
Electrical resistance
Capacitance
Inductance
Impedance
Resonant cavities
Waveguides
Tensors in Relativity
Electromagnetic tensor
Electromagnetic stress-energy tensor
This box: view • talk • edit
Retrieved from "http://en.wikipedia.org/wiki/Magnetism"
Categories: All articles with unsourced statements | Articles with unsourced statements since February 2007 | Electric and magnetic fields in matter | Magnetism

Van der Waals radius

2007-08-14T10:39:00.000-07:00

Van der Waals radius
From Wikipedia, the free encyclopedia
Jump to: navigation, search
Element radius (Å)
Hydrogen 1.20
Carbon 1.7
Nitrogen 1.55
Oxygen 1.52
Fluorine 1.35
Phosphorus 1.9
Sulfur 1.85
Chlorine 1.8
The van der Waals radius of an atom is the radius of an imaginary hard sphere which can be used to model the atom for many purposes. Van der Waals radii are determined from measurements of atomic spacing between pairs of unbonded atoms in crystals.

The van der Waals radius is named after Johannes Diderik van der Waals, winner of the 1910 Nobel Prize in Physics.

Real gases do not behave exactly as predicted. In some cases the deviation can be extremely large. For example, ideal gases could never become liquids or solids, no matter how much they were cooled or compressed. Modifications of the ideal gas law, , were therefore proposed. Particularly useful and well known is the van der Waals equation of state: , where a and b are adjustable parameters determined from experimental measurements carried out on actual gases. Their values vary from gas to gas.

The van der Waals equation also has a microscopic interpretation. Molecules interact with one another. The interaction is strongly repulsive at very short distance, becomes mildly attractive at intermediate range, and vanishes at long distance. The ideal gas law must be corrected when attractive and repulsive forces are considered. For example, the mutual repulsion between molecules has the effect of excluding neighbours from a certain amount of territory around each molecule. Thus, a fraction of the total space becomes unavailable to each molecule as it executes random motion. In the equation of state, this volume of exclusion (nb) should be subtracted from the volume of the container (V), thus: (V - nb). The other term that is introduced in the van der Waals equation, , describes a weak attractive force among molecules, which increases when n increases or V decreases and molecules become more crowded together.

Contents [hide]
1 Van der Waals volume
2 See also
3 References
4 External links

[edit] Van der Waals volume
The van der Waals volume of an atom is the volume of a sphere with the Van der Waals radius of the atom.

Two atoms which are not chemically bonded have a minimum distance between their centers, which is equal to the sum of their Van der Waals radii. However, if the atoms are bonded by a covalent bond, the distance between their centers is smaller.

Therefore, the van der Waals volume of a molecule with covalent bonds is smaller than the sum of the van der Waals volumes of the atoms.

The van der Waals volume of a system of molecules is the sum of their van der Waals volumes.

[edit] See also
van der Waals constant
van der Waals equation
van der Waals force
van der Waals potential

[edit] References
L. Pauling, The Nature of the Chemical Bond, Cornell University Press, USA, 1945.

[edit] External links
van der Waals radii at Webelements
Structural Biology Glossary: van der Waals radii
Retrieved from "http://en.wikipedia.org/wiki/Van_der_Waals_radius"
Category: Chemical properties

Covalent radius

2007-08-14T10:38:00.000-07:00

Covalent radius
From Wikipedia, the free encyclopedia
Jump to: navigation, search
Atomic radius:
Ionic radius
Covalent radius
Metallic radius
van der Waals radius
edit
The covalent radius, rcov, is a measure of the size of atom which forms part of a covalent bond. It is measured either in picometres (pm) or ångströms (Å), with 1 Å = 100 pm.

In principle, the sum of the two covalent radii should equal the covalent bond length between two atoms. This relationship does not hold exactly because the size of an atom is not constant but depends on its chemical environment. In particular, polar covalent bonds tend to be shorter than would be expected on the basis of the sum of covalent radii. Tabulated values of covalent radii are either average or idealized values, which nevertheless show a certain transferability between different situations.

Covalent radii are measured by X-ray diffraction (more rarely, neutron diffraction on molecular crystals). Rotational spectroscopy can also give extremely accurate values of bond lengths. One method takes the covalent radius to be half the single bond length in the element, e.g. d(H–H, in H2) = 74.14 pm so rcov(H) = 37.07 pm: in practice, it is usual to obtain an average value from a variety of covalent compounds, although the difference is usually small. Sanderson has published a recent set of non-polar covalent radii for the main-group elements,[1] but the availabilty of large collections of bond lengths, which are more transferable, from the Cambridge Crystallographic Database[2] has rendered covalent radii obsolete in many situations.

[edit] References
^ Sanderson, R. T. (1983). "Electronegativity and Bond Energy." J. Am. Chem. Soc. 105:2259–61.
^ Allen, F. H.; Kennard, O.; Watson, D. G.; Brammer, L.; Orpen, A. G.; Taylor, R. (1987). "Table of Bond Lengths Determined by X-Ray and Neutron Diffraction." J. Chem. Soc., Perkin Trans. 2 S1–S19.

[edit] External links
WebElements
Retrieved from "http://en.wikipedia.org/wiki/Covalent_radius"
Categories: Chemical properties | Chemical bonding

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
العربية
Asturianu
Català
Česky
Deutsch
Español
Esperanto
Euskara
Français
Italiano
한국어
Lietuvių
Magyar
Македонски
Nederlands
日本語
‪Norsk (bokmål)‬
‪Norsk (nynorsk)‬
Polski
Português
Русский
Slovenščina
Српски / Srpski
Srpskohrvatski / Српскохрватски
Suomi
Svenska
ไทย
Українська
O'zbek
中文

This page was last modified 03:53, 29 June 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Atomic radius

2007-08-14T10:36:00.000-07:00

Atomic radius
From Wikipedia, the free encyclopedia
Jump to: navigation, search
Atomic radius:
Ionic radius
Covalent radius
Metallic radius
van der Waals radius
edit

Atomic radius, and more generally the size of an atom, is not a precisely defined physical quantity, nor is it constant in all circumstances.[1] The value assigned to the radius of a particular atom will always depend on the definition chosen for "atomic radius", and different definitions are more appropriate for different situations.

The term "atomic radius" itself is problematic: it may be restricted to the size of free atoms, or it may be used as a general term for the different measures of the size of atoms, both bound in molecules and free. In the latter case, which is the approach adopted here, it should also include ionic radius, as the distinction between covalent and ionic bonding is itself somewhat arbitrary.[2]

The atomic radius is determined entirely by the electrons: The size of the atomic nucleus is measured in femtometres, 100,000 times smaller than the cloud of electrons. However the electrons do not have definite positions—although they are more likely to be in certain regions than others—and the electron cloud does not have a sharp edge.

Despite (or maybe because of) these difficulties, many different attempts have been made to quantify the size of atoms (and ions), based both on experimental measurements and calculational methods. It is undeniable that atoms do behave as if they were spheres with a radius of 30–300 pm, that atomic size varies in a predictable and explicable manner across the periodic table and that this variation has important consequences for the chemistry of the elements.

Contents [hide]
1 Periodic trends in atomic radius
1.1 Lanthanide contraction
1.2 d-Block contraction
2 Empirically measured atomic radius
3 Calculated atomic radius
4 See also
5 References
6 External links

[edit] Periodic trends in atomic radius
Atomic radius tends to decrease on passing along a period of the periodic table from left to right, and to increase on descending a group. This is, in part, because the distribution of electrons is not completely random. The electrons in an atom are arranged in shells which are, on average, further and further from the nucleus, and which can only hold a certain number of electrons.[3] Each new period of the periodic table corresponds to a new shell which starts to be filled up, and so the outermost electrons are further and further from the nucleus as a group is descended.

The second major effect which determines trends in atomic radius is the charge of the nucleus, which increases with the atomic number, Z. The nucleus is positively charged, and tends to attract the negatively-charged electrons. Passing along a period from left to right, the nuclear charge increases while the electrons are still entering the same shell: the effect is that the physical size of the shell (and hence of the atom) decreases in response.

The increasing nuclear charge is partly counterbalanced by the increasing number of electrons in a phenomenon known as shielding, which is why the size of atoms usually increases as a group is descended. However, there are two occasions where shielding is less effective: in these cases, the atoms are smaller than would otherwise be expected.

[edit] Lanthanide contraction
Main article: Lanthanide contraction
The electrons in the 4f-subshell, which is progressively filled from cerium (Z = 58) to lutetium (Z = 71), are not particularly effective at shielding the increasing nuclear charge from the sub-shells further out. The elements immediately following the lanthanides have atomic radii which are smaller than would be expected and which are almost identical to the atomic radii of the elements immediately above them.[4] Hence hafnium has virtually the same atomic radius (and chemistry) as zirconium, tantalum as niobium etc. The effect of the lanthanide contraction is noticeable up to platinum (Z = 78), after which it is masked by a relativistic effect known as the inert pair effect.

[edit] d-Block contraction
Main article: d-Block contraction
The d-block contraction is less pronounced than the lanthanide contraction but arises from a similar cause. In this case, it is the poor shielding capacity of the 3d-electrons which affects the atomic radii and chemistries of the elements immediately following the first row of the transition metals, from gallium (Z = 31) to bromine (Z = 35).[4]

[edit] Empirically measured atomic radius
Empirically measured atomic radius in picometres (pm) to an accuracy of about 5 pm

Group (vertical) 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Period (horizontal)
1 H
25 He

2 Li
145 Be
105 B
85 C
70 N
65 O
60 F
50 Ne

3 Na
180 Mg
150 Al
125 Si
110 P
100 S
100 Cl
100 Ar
71
4 K
220 Ca
180 Sc
160 Ti
140 V
135 Cr
140 Mn
140 Fe
140 Co
135 Ni
135 Cu
135 Zn
135 Ga
130 Ge
125 As
115 Se
115 Br
115 Kr

5 Rb
235 Sr
200 Y
180 Zr
155 Nb
145 Mo
145 Tc
135 Ru
130 Rh
135 Pd
140 Ag
160 Cd
155 In
155 Sn
145 Sb
145 Te
140 I
140 Xe

6 Cs
260 Ba
215 *
Hf
155 Ta
145 W
135 Re
135 Os
130 Ir
135 Pt
135 Au
135 Hg
150 Tl
190 Pb
180 Bi
160 Po
190 At
Rn

7 Fr
Ra
215 **
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
Uub
Uut
Uuq
Uup
Uuh
Uus
Uuo

Lanthanides *
La
195 Ce
185 Pr
185 Nd
185 Pm
185 Sm
185 Eu
185 Gd
180 Tb
175 Dy
175 Ho
175 Er
175 Tm
175 Yb
175 Lu
175
Actinides **
Ac
195 Th
180 Pa
180 U
175 Np
175 Pu
175 Am
175 Cm
Bk
Cf
Es
Fm
Md
No
Lr

Periodic table of empirically measured atomic radius in picometres (pm) to an accuracy of about 5 pm
See also Periodic table
Reference: J.C. Slater, J. Chem. Phys. 1964, 41, 3199.

[edit] Calculated atomic radius
Calculated atomic radius in picometres (pm)

Group (vertical) 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Period (horizontal)
1 H
53 He
31
2 Li
167 Be
112 B
87 C
67 N
56 O
48 F
42 Ne
38
3 Na
190 Mg
145 Al
118 Si
111 P
98 S
88 Cl
79 Ar
71
4 K
243 Ca
194 Sc
184 Ti
176 V
171 Cr
166 Mn
161 Fe
156 Co
152 Ni
149 Cu
145 Zn
142 Ga
136 Ge
125 As
114 Se
103 Br
94 Kr
88
5 Rb
265 Sr
219 Y
212 Zr
206 Nb
198 Mo
190 Tc
183 Ru
178 Rh
173 Pd
169 Ag
165 Cd
161 In
156 Sn
145 Sb
133 Te
123 I
115 Xe
108
6 Cs
298 Ba
253 *
Hf
208 Ta
200 W
193 Re
188 Os
185 Ir
180 Pt
177 Au
174 Hg
171 Tl
156 Pb
154 Bi
143 Po
135 At
Rn
120
7 Fr
Ra
**
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
Uub
Uut
Uuq
Uup
Uuh
Uus
Uuo

Lanthanides *
La
Ce
Pr
247 Nd
206 Pm
205 Sm
238 Eu
231 Gd
233 Tb
225 Dy
228 Ho
Er
226 Tm
222 Yb
222 Lu
217
Actinides **
Ac
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr

Periodic table of calculated atomic radius in picometres (pm)
See also Periodic table
Reference: E. Clementi, D.L.Raimondi, and W.P. Reinhardt, J. Chem. Phys. 1963, 38, 2686.

[edit] See also
Atomic radii of the elements (data page)
Chemical bond
Bond length
Steric hindrance

[edit] References
^ Cotton, F. A.; Wilkinson, G. (1988). Advanced Inorganic Chemistry (5th Edn). New York: Wiley. ISBN 0-471-84997-9. p. 1385.
^ See also the definition of an atom as "the smallest unit quantity of an element that is capable of existence whether alone or in chemical combination with other atoms of the same or other elements." IUPAC Commission on the Nomenclature of Inorganic Chemistry (1990). Nomenclature of Inorganic Chemistry. Oxford: Blackwell Scientific. ISBN 0-632-02494-1. p. 35.
^ The nth electron shell can hold 2n2 electrons.
^ a b Jolly, William L. (1991). Modern Inorganic Chemistry (2nd Edn.). New York: McGraw-Hill. ISBN 0-07-112651-1. p. 22.

[edit] External links
WebElements
Retrieved from "http://en.wikipedia.org/wiki/Atomic_radius"
Categories: Chemical properties | Chemical bonding

Ionization potential

2007-08-14T10:34:00.000-07:00

Ionization potential
From Wikipedia, the free encyclopedia
(Redirected from Ionization energy)
Jump to: navigation, search
The ionization potential, ionization energy or EI of an atom or molecule is the energy required to remove one mole of electrons from one mole of isolated gaseous atoms or ions. More generally, the nth ionization energy is the energy required to strip it of an nth mole of electrons after the first n − 1 mole of electrons have already been removed. It is considered in physical chemistry as a measure of the "reluctance" of an atom or ion to surrender an electron, or the "strength" by which the electron is bounded; the greater the ionization energy, the more difficult it is to remove an electron.

Contents [hide]
1 Values and trends
2 Reactivity
3 Electrostatic explanation
4 Quantum-mechanical explanation
5 See also

[edit] Values and trends
Main article: Ionization energies of the elements
Generally speaking, atomic ionization energies decrease down a group (a.k.a column) of the periodic table, and increase left-to-right across a period. Ionization energy exhibits a strong negative correlation with atomic radius. Successive ionization energies of any given element increase markedly. Particularly dramatic increases occur after any given block of atomic orbitals is exhausted, except when progressing to the next s orbital. This is because, after all the electrons are removed from an orbital, the next ionization energy involves removing an electron from an orbital closer to the nucleus. Electrons in the closer orbital experience greater forces of electrostatic attraction, and thus, require more energy to be removed.

Some values for elements of the third period are given in the following table:

Successive ionization energies in kJ/mol Element First Second Third Fourth Fifth Sixth Seventh
Na 496 4,560
Mg 738 1,450 7,730
Al 577 1,816 2,881 11,600
Si 786 1,577 3,228 4,354 16,100
P 1,060 1,890 2,905 4,950 6,270 21,200
S 999.6 2,260 3,375 4,565 6,950 8,490 27,107
Cl 1,256 2,295 3,850 5,160 6,560 9,360 11,000
Ar 1,520 2,665 3,945 5,770 7,230 8,780 12,000

In order to determine how many electrons are in the outermost shell of an element, one can use the ionization energy. If, for example, it required 1,500 kJ/mol to remove one electron and required 6,000 kJ/mol to remove another electron and then 5,000 kJ/mol, etc. this means that the element had one electron in its outermost shell. This means that the element is a metal and in order for this element to achieve a stable octet, it looks to lose one electron. Thus, the first electron is easy to remove and consequently the ionization energy is low. Notice, however, that once the stable octet has been formed, it becomes much more difficult to remove the next electron. If that electron can be removed the consequent one can be removed a bit more easily.

[edit] Reactivity
The ionization potential is basically the amount of force that the nucleus applies pulling the electrons towards itself.

The ionization potential is what determines the reactivity of an element. In non-metals, the ionization potential is proportional to the reactivity. In metals, the relationship is inverse, so as the ionization potential goes up the reactivity goes down.

[edit] Electrostatic explanation
Atomic ionization energy can be predicted by an analysis using electrostatic potential and the Bohr model of the atom, as follows.

Consider an electron of charge -e, and an ion with charge +ne, where n is the number of electrons missing from the ion. According to the Bohr model, if the electron were to approach and bind with the atom, it would come to rest at a certain radius a. The electrostatic potential V at distance a from the ionic nucleus, referenced to a point infinitely far away, is:

Since the electron is negatively charged, it is drawn to this positive potential. (The value of this potential is called the ionization potential). The energy required for it to "climb out" and leave the atom is:

This analysis is incomplete, as it leaves the distance a as an unknown variable. It can be made more rigorous by assigning to each electron of every chemical element a characteristic distance, chosen so that this relation agrees with experimental data.

It is possible to expand this model considerably by taking a semi-classical approach, in which momentum is quantized. This approach works very well for the hydrogen atom, which only has one electron. The magnitude of the angular momentum for a circular orbit is:

The total energy of the atom is the sum of the kinetic and potential energies, that is:

Velocity can be eliminated from the kinetic energy term by setting the Coulomb attraction equal to the centripetal force, giving:

Now the energy can be found in terms of k, e, and r. Using the new value for the kinetic energy in the total energy equation above, it is found that:

Solving the angular momentum for v and substituting this into the expression for kinetic energy, we have:

This establishes the dependence of the radius on n. That is:

At its smallest value, n is equal to 1 and r is the Bohr radius a0. Now, the equation for the energy can be established in terms of the Bohr radius. Doing so gives the result:

This can be expanded to larger nuclei by incorporating the atomic number into the equation.

[edit] Quantum-mechanical explanation
According to the more sophisticated theory of quantum mechanics, the location of an electron is best described as a "cloud" of likely locations that ranges near and far from the nucleus. The energy can be calculated by integrating over this cloud. This cloud corresponds to a wavefunction or, more specifically, to a linear combination of Slater determinants, i.e., according to Pauli exclusion principle, antisymmetrized products of atomic or molecular orbitals. This linear combination is called a configuration interaction expansion of the electronic wavefunction.

In general, calculating the nth ionization energy requires subtracting the energy of a Z − n + 1 electron system from the energy of a Z − n electron system. Calculating these energies is not simple, but is a well-studied problem and is routinely done in computational chemistry. At the lowest level of approximation, the ionization energy is provided by Koopmans' theorem.

[edit] See also
Bragg-Gray Cavity Theory
Electronegativity
Ionization
The ionization potential is equal to the ionization energy divided by the charge of an electron.
The work function is the energy required to strip an electron from a solid.
Ion
Koopmans' theorem
Di-tungsten tetra(hpp) has the lowest recorded ionization energy for a stable chemical compound.
Electron affinity
Retrieved from "http://en.wikipedia.org/wiki/Ionization_potential"
Categories: Ions | Molecular physics | Atomic physics | Chemical properties | Quantum chemistry

Electronegativity

2007-08-14T10:29:00.000-07:00

Electronegativity
From Wikipedia, the free encyclopedia
Jump to: navigation, search
Electronegativity, symbol χ, is a chemical property which describes the power of an atom (or, more rarely, a functional group) to attract electrons towards itself.[1] First proposed by Linus Pauling in 1932 as a development of valence bond theory,[2] it has been shown to correlate with a number of other chemical properties. Electronegativity cannot be directly measured and must be calculated from other atomic or molecular properties. Several methods of calculation have been proposed and, although there may be small differences in the numerical values of the electronegativity, all methods show the same periodic trends between elements.

The most commonly used method of calculation is that originally proposed by Pauling. This gives a dimensionless quantity, commonly referred to as Pauling electronegativity, on a relative scale running from 0.7 to 4.0 (hydrogen = 2.2). When other methods of calculation are used, it is conventional (although not obligatory) to quote the results on a scale which covers the same range of numerical values: this is known as an electronegativity in Pauling units.

Electronegativity, as it is usually calculated, is not strictly an atomic property, but rather a property of an atom in a molecule:[3] the equivalent property of a free atom is its electron affinity. It is to be expected that the electronegativity of an element will vary with its chemical environment,[4] but it is usually considered to be a transferable property, that is to say that similar values will be valid in a variety of situations.

Contents [hide]
1 Electronegativities of the elements
2 Methods of calculation
2.1 Pauling electronegativity
2.2 Mulliken electronegativity
2.3 Allred-Rochow electronegativity
2.4 Sanderson electronegativity
2.5 Allen electronegativity
3 Correlation of electronegativity with other properties
4 Trends in electronegativity
4.1 Periodic trends
4.2 Variation of electronegativity with oxidation number
5 Group electronegativity
6 See also
7 References
8 External links

[edit] Electronegativities of the elements
→ Atomic radius decreases → Ionization energy increases → Electronegativity increases →
Group (vertical) 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Period (horizontal)
1 H
2.20 He

2 Li
0.98 Be
1.57 B
2.04 C
2.55 N
3.04 O
3.44 F
3.98 Ne

3 Na
0.93 Mg
1.31 Al
1.61 Si
1.90 P
2.19 S
2.58 Cl
3.16 Ar

4 K
0.82 Ca
1.00 Sc
1.36 Ti
1.54 V
1.63 Cr
1.66 Mn
1.55 Fe
1.83 Co
1.88 Ni
1.91 Cu
1.90 Zn
1.65 Ga
1.81 Ge
2.01 As
2.18 Se
2.55 Br
2.96 Kr
3.00
5 Rb
0.82 Sr
0.95 Y
1.22 Zr
1.33 Nb
1.6 Mo
2.16 Tc
1.9 Ru
2.2 Rh
2.28 Pd
2.20 Ag
1.93 Cd
1.69 In
1.78 Sn
1.96 Sb
2.05 Te
2.1 I
2.66 Xe
2.6
6 Cs
0.79 Ba
0.89 *
Hf
1.3 Ta
1.5 W
2.36 Re
1.9 Os
2.2 Ir
2.20 Pt
2.28 Au
2.54 Hg
2.00 Tl
1.62 Pb
2.33 Bi
2.02 Po
2.0 At
2.2 Rn

7 Fr
0.7 Ra
0.9 **
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
Uub
Uut
Uuq
Uup
Uuh
Uus
Uuo

Lanthanides *
La
1.1 Ce
1.12 Pr
1.13 Nd
1.14 Pm
1.13 Sm
1.17 Eu
1.2 Gd
1.2 Tb
1.1 Dy
1.22 Ho
1.23 Er
1.24 Tm
1.25 Yb
1.1 Lu
1.27
Actinides **
Ac
1.1 Th
1.3 Pa
1.5 U
1.38 Np
1.36 Pu
1.28 Am
1.13 Cm
1.28 Bk
1.3 Cf
1.3 Es
1.3 Fm
1.3 Md
1.3 No
1.3 Lr

Periodic table of electronegativity using the Pauling scale
See also Periodic table

[edit] Methods of calculation

[edit] Pauling electronegativity
Pauling first proposed[2] the concept of electronegativity in 1932 as an explanation of the fact that the covalent bond between two different atoms (A–B) is stronger than would be expected by taking the average of the strengths of the A–A and B–B bonds. According to valence bond theory, of which Pauling was a notable proponent, this "additional stabilization" of the heteronuclear bond is due to the contribution of ionic canonical forms to the bonding.

The difference in electronegativity between atoms A and B is given by:

where the dissociation energies, Ed, of the A–B, A–A and B–B bonds are expressed in electronvolts, the factor (eV)−½ being included to ensure a dimensionless result. Hence, the difference in Pauling electronegativity between hydrogen and bromine is 0.73 (dissociation energies: H–Br, 3.79 eV; H–H, 4.52 eV; Br–Br 2.00 eV)

As only differences in electronegativity are defined, it is necessary to choose an arbitrary reference point in order to construct a scale. Hydrogen was chosen as the reference, as it forms covalent bonds with a large variety of elements: its electronegativity was fixed first[2] at 2.1, later revised[5] to 2.20. It is also necessary to decide which of the two elements is the more electronegative (equivalent to choosing one of the two possible signs for the square root). This is done by "chemical intuition": in the above example, hydrogen bromide dissolves in water to form H+ and Br− ions, so it may be assumed that bromine is more electronegative than hydrogen.

To calculate a Pauling electronegativity for an element, it is necessary to have data on the dissociation energies of at least two types of covalent bond formed by that element. Allred updated Pauling's original values in 1961 to take account of the greater availability of thermodynamic data,[5] and it is these "revised Pauling" values of the electronegativity which are most usually used.

[edit] Mulliken electronegativity

The correlation between Mulliken electronegativities (x-axis, in kJ/mol) and Pauling electronegativities (y-axis).Mulliken proposed that the arithmetic mean of the first ionization energy and the electron affinity should be a measure of the tendency of an atom to attract electrons.[6] As this definition is not dependent on an arbitrary relative scale, it has also been termed absolute electronegativity,[7] with the units of kilojoules per mole or electronvolts.

However, it is more usual to make use of a linear transformation to transform these absolute values into values which resemble the more familiar Pauling values. For ionization energies and electron affinities in electronvolts,[8]

χ = 0.187(Ei + Eea) + 0.17
and for energies in kilojoules per mole,[9]

The Mulliken electronegativity can only be calculated for an element for which the electron affinity is known, fifty-seven elements as of 2006.

[edit] Allred-Rochow electronegativity

The correlation between Allred-Rochow electronegativities (x-axis, in Å−2) and Pauling electronegativities (y-axis).Allred and Rochow considered[10] that electronegativity should be related to the charge experienced by an electron on the "surface" of an atom: the higher the charge per unit area of atomic surface, the greater the tendency of that atom to attract electrons. The effective nuclear charge, Z* experienced by valence electrons can be estimated using Slater's rules, while the surface area of an atom in a molecule can be taken to be proportional to the square of the covalent radius, rcov. When rcov is expressed in ångströms,

.

[edit] Sanderson electronegativity

The correlation between Sanderson electronegativities (x-axis, arbitrary units) and Pauling electronegativities (y-axis).Sanderson has also noted the relationship between electronegativity and atomic size, and has proposed a method of calculation based on the reciprocal of the atomic volume.[11] With a knowledge of bond lengths, Sanderson electronegativities allow the estimation of bond energies in a wide range of compounds.[12]

[edit] Allen electronegativity

The correlation between Allen electronegativities (x-axis, in kJ/mol) and Pauling electronegativities (y-axis).Perhaps the simplest definition of electronegativity is that of Allen, who has proposed that it is related to the average energy of the valence electrons in a free atom,[13]

where εs,p are the one-electron energies of s- and p-electrons in the free atom and ns,p are the number of s- and p-electrons in the valence shell. It is usual to apply a scaling factor, 1.75×10−3 for energies expressed in kilojoules per mole or 0.169 for energies measured in electronvolts, to give values which are numerically similar to Pauling electronegativities.

The one-electron energies can be determined directly from spectroscopic data, and so electronegativities calculated by this method are sometimes referred to as spectroscopic electronegativities. The necessary data are available for almost all elements, and this method allows the estimation of electronegativities for elements which cannot be treated by the other methods, e.g. francium, which has an Allen electronegativity of 0.67.[14] However, it is not clear what should be considered to be valence electrons for the d- and f-block elements, which leads to an ambiguity for their electronegativities calculated by the Allen method.

[edit] Correlation of electronegativity with other properties

The variation of the isomer shift (y-axis, in mm/s) of [SnX6]2− anions, as measured by 119Sn Mössbauer spectroscopy, against the sum of the Pauling electronegativities of the halide substituents (x-axis).The wide variety of methods of calculation of electronegativities, which all give results which correlate well with one another, is one indication of the number of chemical properties which might be affected by electronegativity. The most obvious application of electronegativities is in the discussion of bond polarity, for which the concept was introduced by Pauling. In general, the greater the difference in electronegativity between two atoms, the more polar the bond that will be formed between them, with the atom having the higher electronegativity being at the negative end of the dipole. Pauling proposed an equation to relate "ionic character" of a bond to the difference in electronegativity of the two atoms,[3] although this has fallen somewhat into disuse.

Several correlations have been shown between infrared stretching frequencies of certain bonds and the electronegativities of the atoms involved:[15] however, this is not surprising as such strectching frquencies depend in part on bond strength, which enters into the calculation of Pauling electronegativities. More convincing are the correlations between electronegativity and chemical shifts in NMR spectroscopy[16] or isomer shifts in Mössbauer spectroscopy[17] (see figure). Both these measurements depend on the s-electron density at the nucleus, and so are a good indication that the different measures of electronegativity really are describity "the ability of an atom in a molecule to attract electrons to itself".[1][3]

[edit] Trends in electronegativity

[edit] Periodic trends

The variation of Pauling electronegativity (y-axis) as one descends the main groups of the Periodic table from the second period to the sixth period.In general, electronegativity increases on passing from left to right along a period, and decreases on descending a group. Hence, fluorine is undoubtedly the most electronegative of the elements while caesium is the least electronegative, at least of those elements for which substantial data are available.[14]

There are some exceptions to this general rule. Gallium and germanium have higher electronegativities than aluminium and silicon respectively because of the d-block contraction. Elements of the fourth period immediately after the first row of the transition metals have unusually small atomic radii because the 3d-electrons are not effective at shielding the increased nuclear charge, and smaller atomic size correlates with higher electronegativity (see Allred-Rochow electronegativity, Sanderson electronegativity above). The anomalously high electronegativity of lead, particularly when compared to thallium and bismuth, appears to be an artefact of data selection (and data availability): methods of calculation other than the Pauling method show the normal periodic trends for these elements.

[edit] Variation of electronegativity with oxidation number
In inorganic chemistry it is common to consider a single value of the electronegativity to be valid for most "normal" situations. While this approach has the advantage of simplicity, it is clear that the electronegativity of an element is not an invariable atomic property and, in particular, increases with the oxidation state of the element.

Allred used the Pauling method to calculate separate electronegativities for different oxidation states of the handful of elements (including tin and lead) for which sufficient data was available.[5] However, for most elements, there are not enough different covalent compounds for which bond dissociation energies are known to make this approach feasible. This is particularly true of the transition elements, where quoted electronegativity values are usually, of necessity, averages over several different oxidation states and where trends in electronegativity are harder to see as a result.

Acid Formula Chlorine
oxidation
state pKa
Hypochlorous acid HClO +1 +7.5
Chlorous acid HClO2 +3 +2.0
Chloric acid HClO3 +5 −1.0
Perchloric acid HClO4 +7 −10
The chemical effects of this increase in electronegativity can be seen both in the structures of oxides and halides and in the acidity of oxides and oxoacids. Hence CrO3 and Mn2O7 are acidic oxides with low melting points, while Cr2O3 is amphoteric and Mn2O3 is a completely basic oxide.

The effect can also be clearly seen in the dissociation constants of the oxoacids of chlorine. The effect is much larger than could be explained by the negative charge being shared among a larger number of oxygen atoms, which would lead to a difference in pKa of log10(¼) = −0.6 between hypochlorous acid and perchloric acid. As the oxidation state of the central chlorine atom increases, more electron density is drawn from the oxygen atoms onto the chlorine, reducing the partial negative charge on the oxygen atoms and increasing the acidity.

[edit] Group electronegativity
Main article: Electronic effect of substituents
In organic chemistry, electronegativity is associated more with different functional groups than with individual atoms. The terms group electronegativity and substituent electronegativity are used synonymously. However, it is common to distinguish between the inductive effect and the resonance effect, which might be described as σ- and &pi-electronegativities respectively. There are a number of linear free energy relationships which have been used to quantify these effects, of which the Hammett equation is the best known. Kabachnik parameters are group electronegativities for use in organophosphorus chemistry.

[edit] See also
Electronegativities of the elements (data page)

[edit] References
Jolly, William L. (1991). Modern Inorganic Chemistry (2nd Edn.). New York: McGraw-Hill. ISBN 0-07-112651-1. pp. 71–76.
Mullay, J. (1987). "Estimation of atomic and group electronegativities." Struct. Bond. 66:1–25.
^ a b "Electronegativity.", Compendium of Chemical Terminology
^ a b c Pauling, Linus (1932). J. Am. Chem. Soc. 54:3570.
^ a b c Pauling, Linus (1960). Nature of the Chemical Bond (3rd Edn.). Ithaca, NY: Cornell University Press. pp. 88–107.
^ Greenwood, N. N.; Earnshaw, A. (1984). Chemistry of the Elements. Oxford: Pergamon. ISBN 0-08-022057-6. p. 30.
^ a b c Allred, A. L. (1961). J. Inorg. Nucl. Chem. 17:215.
^ Mulliken, R. S. (1934). J. Chem. Phys. 2:782. Mulliken, R. S. (1935). J. Chem. Phys. 3:573.
^ Pearson, R. G. (1985). J. Am. Chem. Soc. 107:6801.
^ Huheey, J. E. (1978). Inorganic Chemistry (2nd Edn.). New York: Harper & Row. p. 167.
^ This second relation has been recalculated using the best values of the first ionization energies and electron affinities available in 2006.
^ Allred, A. L.; Rochow, E. G. (1958). J. Inorg. Nucl. Chem. 5:264.
^ Sanderson, R. T. (1983). "Electronegativity and bond energy." J. Am. Chem. Soc. 105:2259.
^ Sanderson, R. T. (1983). Polar Covalence. New York: Academic Press.
^ Allen, L. C. (1989). J. Am. Chem. Soc. 111:9003.
^ a b The widely quoted Pauling electronegativity of 0.7 for francium is an extrapolated value of uncertain provenance. The Allen electronegativity of caesium is 0.66.
^ See, e.g., Bellamy, L. J. (1958). The Infra-Red Spectra of Complex Molecules (2nd Edn.). New York: Wiley. p. 392.
^ Spieseke, H.; Schneider, W. G. (1961). J. Chem. Phys. 35:722.
^ Clasen, C. A.; Good, M. L. (1970). Inorg. Chem. 9:817.

[edit] External links
Wikimedia Commons has media related to:
ElectronegativityWebElements, lists values of electronegativities by a number of different methods of calculation
[hide]v • d • ePeriodic tables
Layouts Standard · Vertical · Full names · Names and atomic masses · Text for last · Huge table · Metals and nonmetals · Blocks · Valences · Inline f-block · 218 elements · Electron configurations · Atomic masses · Electronegativities · Alternatives
Lists of elements Name · Atomic symbol · Atomic number · Boiling point · Melting point · Density · Atomic mass
Groups 1 · 2 · 3 · 4 · 5 · 6 · 7 · 8 · 9 · 10 · 11 · 12 · 13 · 14 · 15 · 16 · 17 · 18
Periods: 1 · 2 · 3 · 4 · 5 · 6 · 7 · 8
Series Alkalis · Alkaline earths · Lanthanides · Actinides · Transition metals · Poor metals · Metalloids · Nonmetals · Halogens · Noble gases
Blocks s-block · p-block · d-block · f-block · g-block

Retrieved from "http://en.wikipedia.org/wiki/Electronegativity"
Category: Chemical properties

Oxidation number

2007-08-14T10:28:00.000-07:00

Oxidation number
From Wikipedia, the free encyclopedia
Jump to: navigation, search
The oxidation number of an element in a molecule or complex is the charge that it would have if all the ligands (basically, atoms that donate electrons) were removed along with the electron pairs that were shared with the central atom[1]. It means that the oxidation number is the charge an atom had if it was in a compound composed of ions. It's used in the inorganic nomenclature of inorganic compounds. It is represented by a Roman numeral; the plus sign is omitted for positive oxidation numbers. The oxidation number is placed either as a right superscript to the element symbol, e.g. FeIII, or in parentheses after the name of the element, e.g. iron(III): in the latter case, there is no space between the element name and the oxidation number. The oxidation number can also be written with a number and either a + or - sign after it. If the element creates a positively charged ion, the oxidation number will have a + sign after it, (example-hydrogen 1+). If the element creates a negatively charged ion, the oxidation number will have a - sign after it, (example-oxygen 2-). The change in the oxidation number represents the number of electrons gained or lost in a chemical reaction.

The oxidation number is usually numerically equal to the oxidation state of the central atom. However, for a variety of reasons, the oxidation state of transition metals can be difficult to determine[2]. The most-accepted answer is that the electron pairs forming the coordination bonds are mostly associated with the ligands: this is a good approximation for most Werner-type complexes, but much less true for organometallic compounds as well as for certain hydrido complexes, dithiolene complexes and nitrosyl complexes.

[edit] Rules for the assigning of oxidation numbers
All species in their elemental form are given the oxidation number of zero.
All monoatomic ions have the same oxidation number as the charge on the ion. e.g. Mg2+ has the oxidation number of +2.
All combined hydrogen has an oxidation number of +1 (except metal hydrides where its oxidation number is -1).
All combined oxygen has an oxidation number of -2 (except peroxides where the oxidation number is -1, and compounds with fluorine where it can be positive).
Fluorine always has an oxidation number of -1.
In polyatomic species, the sum of the oxidation numbers of the element in the ion equals the charge on that species (we can use this to find the oxidation number of elements in polyatomic species).
Group 1 elements such as K and Na and Group 2 elements such as Mg always have a +1 and +2 oxidation state in compounds, respectively.
Cl has a range of oxidation states when bonded to O. However, its oxidation number is always -1 when bonded to ionic compounds.

[edit] Example
The oxidation number of sulfur in sulfuric acid (H2SO4) can be calculated from the rules above. Because this is a polyatomic species, the individual oxidation numbers must sum to equal the overall charge, which in this case is zero. Hydrogen has an oxidation number of +1, so the sum of the oxidation numbers of H2 = 2. Oxygen has an oxidation number of -2, so the sum of oxidation numbers of O4 = -8. Since the overall sum must equal zero, the oxidation state of sulfur can be calculated as +6 (8-2).

[edit] References
^ International Union of Pure and Applied Chemistry. "oxidation number". Compendium of Chemical Terminology Internet edition.
^ Oxidation numbers of transition metals
Assigning oxidation numbers
Retrieved from "http://en.wikipedia.org/wiki/Oxidation_number"
Category: Chemical nomenclature

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
Deutsch
Italiano
‪Norsk (bokmål)‬

This page was last modified 01:32, 12 August 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Crystal structure

2007-08-14T10:26:00.000-07:00

Crystal structure
From Wikipedia, the free encyclopedia
Jump to: navigation, search

Enargite crystalsIn mineralogy and crystallography, a crystal structure is a unique arrangement of atoms in a crystal. A crystal structure is composed of a motif, a set of atoms arranged in a particular way, and a lattice. Motifs are located upon the points of a lattice, which is an array of points repeating periodically in three dimensions. The points can be thought of as forming identical tiny boxes, called unit cells, that fill the space of the lattice. The lengths of the edges of a unit cell and the angles between them are called the lattice parameters. The symmetry properties of the crystal are embodied in its space group. A crystal's structure and symmetry play a role in determining many of its properties, such as cleavage, electronic band structure, and optical properties.

Contents [hide]
1 Unit cell
2 Classification of crystals by symmetry
2.1 Crystal system
2.2 The Bravais lattices
2.3 Point and space groups
3 Physical properties
3.1 Defects in crystals
3.2 Crystal symmetry and physical properties
4 See also
5 External links

[edit] Unit cell
The crystal structure of a material is often discussed in terms of its unit cell. The unit cell is a tiny box containing one or more motifs, a spatial arrangement of atoms. The units cells are tiled in three-dimensional space to describe the crystal. The unit cell is given by its lattice parameters, the length of the cell edges and the angles between them, while the positions of the atoms inside the unit cell are described by the set of atomic positions (xi,yi,zi) measured from a lattice point.

Although there are an infinite number of ways to specify a unit cell, for each crystal structure there is a conventional unit cell, which is chosen to display the full symmetry of the crystal (see below). However, the conventional unit cell is not always the smallest possible choice. A primitive unit cell of a particular crystal structure is the smallest possible unit cell one can construct such that, when tiled, completely fills space. This primitive unit cell does not always display all the symmetries inherent in the crystal. A Wigner-Seitz cell is a particular kind of primitive cell which has the same symmetry as the lattice.

There are only seven possible crystal systems that atoms can pack together to produce an infinite 3D space lattice in such a way that each lattice point has an identical environment to that around every other lattice point.

[edit] Classification of crystals by symmetry
The defining property of a crystal is its inherent symmetry, by which we mean that under certain operations the crystal remains unchanged. For example, rotating the crystal 180 degrees about a certain axis may result in an atomic configuration which is identical to the original configuration. The crystal is then said to have a twofold rotational symmetry about this axis. In addition to rotational symmetries like this, a crystal may have symmetries in the form of mirror planes and translational symmetries, and also the so-called compound symmetries which are a combination of translation and rotation/mirror symmetries. A full classification of a crystal is achieved when all of these inherent symmetries of the crystal are identified.

[edit] Crystal system
The crystal systems are a grouping of crystal structures according to the axial system used to describe their lattice. Each crystal system consists of a set of three axes in a particular geometrical arrangement. There are seven unique crystal systems. The simplest and most symmetric, the cubic (or isometric) system, has the symmetry of a cube, that is, it exhibits four threefold rotational axes oriented at 109.5 degrees (the tetrahedral angle) with respect to each other. These threefold axes lie along the body diagonals of the cube. This definition of a cubic is correct, although many textbooks incorrectly state that a cube is defined by three mutually perpendicular axes of equal length - if this were true there would be far more than 14 Bravais lattices. The other six systems, in order of decreasing symmetry, are hexagonal, tetragonal, rhombohedral (also known as trigonal), orthorhombic, monoclinic and triclinic. Some crystallographers consider the hexagonal crystal system not to be its own crystal system, but instead a part of the trigonal crystal system. The crystal system and Bravais lattice of a crystal describe the (purely) translational symmetry of the crystal.

[edit] The Bravais lattices
Crystal system Lattices:
triclinic
monoclinic simple base-centered

orthorhombic simple base-centered body-centered face-centered

hexagonal
rhombohedral
(trigonal)
tetragonal simple body-centered

cubic
(isometric) simple body-centered face-centered

When the crystal systems are combined with the various possible lattice centerings, we arrive at the Bravais lattices. They describe the geometric arrangement of the lattice points, and thereby the translational symmetry of the crystal. In three dimensions, there are 14 unique Bravais lattices which are distinct from one another in the translational symmetry they contain. All crystalline materials recognized until now (not including quasicrystals) fit in one of these arrangements. The fourteen three-dimensional lattices, classified by crystal system, are shown to the right. The Bravais lattices are sometimes referred to as space lattices.

The crystal structure consists of the same group of atoms, the basis, positioned around each and every lattice point. This group of atoms therefore repeats indefinitely in three dimensions according to the arrangement of one of the 14 Bravais lattices. The characteristic rotation and mirror symmetries of the group of atoms, or unit cell, is described by its crystallographic point group.

[edit] Point and space groups
The crystallographic point group or crystal class is the mathematical group comprising the symmetry operations that leave at least one point unmoved and that leave the appearance of the crystal structure unchanged. These symmetry operations can include reflection, which reflects the structure across a reflection plane, rotation, which rotates the structure a specified portion of a circle about a rotation axis, inversion which changes the sign of the coordinate of each point with respect to a center of symmetry or inversion point and improper rotation, which consists of a rotation about an axis followed by an inversion. Rotation axes (proper and improper), reflection planes, and centers of symmetry are collectively called symmetry elements. There are 32 possible crystal classes. Each one can be classified into one of the seven crystal systems.

The space group of the crystal structure is composed of the translational symmetry operations in addition to the operations of the point group. These include pure translations which move a point along a vector, screw axis, which rotate a point around an axis while translating parallel to the axis, and glide planes, which reflect a point through a plane while translating it parallel to the plane. There are 230 distinct space groups.

[edit] Physical properties

[edit] Defects in crystals
Real crystals feature defects or irregularities in the ideal arrangements described above and it is these defects that critically determine many of the electrical and mechanical properties of real materials. In particular dislocations in the crystal lattice allow shear at much lower stress than that needed for a perfect crystal structure[citation needed].

[edit] Crystal symmetry and physical properties
Twenty of the 32 crystal classes are so-called piezoelectric, and crystals belonging to one of these classes (point groups) display piezoelectricity. All 20 piezoelectric classes lack a center of symmetry. Any material develops a dielectric polarization when an electric field is applied, but a substance which has such a natural charge separation even in the absence of a field is called a polar material. Whether or not a material is polar is determined solely by its crystal structure. Only 10 of the 32 point groups are polar. All polar crystals are pyroelectric, so the 10 polar crystal classes are sometimes referred to as the pyroelectric classes.

There are a few crystal structures, notably the perovskite structure, which exhibit ferroelectric behaviour. This is analogous to ferromagnetism, in that, in the absence of an electric field during production, the ferroelectric crystal does not exhibit a polarisation. Upon the application of an electric field of sufficient magnitude, the crystal becomes permanently polarised. This polarisation can be reversed by a sufficiently large counter-charge, in the same way that a ferromagnet can be reversed. However, it is important to note that, although they are called ferroelectrics, the effect is due to the crystal structure, not the presence of a ferrous metal.

Incommensurate crystals have period-varying translational symmetry. The period between nodes of symmetry is constant in most crystals. The distance between nodes in an incommensurate crystal is dependent on the number of nodes between it and the base node.

[edit] See also
Cleavage (crystal)
Crystal
Crystal engineering
Crystallography
Crystallographic point group
Crystallographic defect
Crystal growth
Liquid crystal
Miller Index
Patterson function
Quasicrystals
Seed crystal
For more detailed information in specific technology applications see materials science, ceramic, or metallurgy.

[edit] External links
Appendix A from the manual for Atoms, software for XAFS
Intro to Minerals: Crystal Class and System
Introduction to Crystallography and Mineral Crystal Systems
Crystal planes and Miller indices
Interactive 3D Crystal models
Crystal Lattice Structures: Other Crystal Structure Web Sites
Retrieved from "http://en.wikipedia.org/wiki/Crystal_structure"
Categories: All articles with unsourced statements | Articles with unsourced statements since February 2007 | Chemical properties | Condensed matter physics | Crystallography | Materials science | Mineralogy

Heat capacity

2007-08-14T10:24:00.000-07:00

[edit] Heat capacity
Heat capacity (symbol: C) — as distinct from specific heat capacity — is the measure of the heat energy required to increase the temperature of material object by a certain temperature interval. Heat capacity is an extensive property because its value is proportional to the amount of material in the object; for example, a bathtub of water has a greater heat capacity than a cup of water.

Heat capacity is usually expressed in units of J K–1 (or J/K), subject to the caveats and exceptions detailed in both Basic metrics of specific heat capacity and Symbols and standards, above. For instance, one could write that the gasoline in a 55-gallon drum has an average heat capacity of 347 kJ/K.

The uncertainty of an object’s measured quantity is rarely better than one percent and this places an upper limit on the accuracy and precision of most stated values of heat capacity. Accordingly, it is usually unnecessary as a practical matter, to specify the defined state at which the measurement was made; e.g. “(25 °C, 100 kPa).” It most cases, it is assumed that the substance’s specific heat capacity is a published value and the object’s quantity is subject to such a sizable relative uncertainty that it renders this detail moot. An exception would be when an object has an accurately known or precisely defined quantity; e.g. “The heat capacity of the International Prototype Kilogram is 133 J/K (25 °C).” Another exception would be when the defined state varies significantly from standard conditions.

[edit] Factors that affect specific heat capacity

Molecules have internal structure because they are composed of atoms that have different ways of moving within molecules. Kinetic energy stored in these internal degrees of freedom contributes to a substance’s specific heat capacity and not to its temperature.Degrees of freedom: Molecules are quite different from the monatomic gases like helium and argon. With monatomic gases, heat energy comprises only translational motions. Translational motions are ordinary, whole-body movements in 3D space whereby particles move about and exchange energy in collisions (like rubber balls in a vigorously shaken container). These simple movements in the three X, Y, and Z–axis dimensions of space means monatomic atoms have three translational degrees of freedom. Molecules, however, have various internal vibrational and rotational degrees of freedom because they are complex objects; they are a population of atoms that can move about within a molecule in different ways (see animation at right). Heat energy is stored in these internal motions. For instance, nitrogen, which is a diatomic molecule, has five active degrees of freedom: the three comprising translational motion plus two rotational degrees of freedom internally. Not surprisingly, nitrogen has five-thirds the constant-volume molar heat capacity as do the monatomic gases.[2] See Thermodynamic temperature for more information on translational motions, kinetic (heat) energy, and their relationship to temperature.
Molar mass: When the specific heat capacity, c, of a material is measured (lowercase c means the unit quantity is in terms of mass), different values arise because different substances have different molar masses (essentially, the weight of the individual atoms or molecules). Heat energy arises, in part, due to the number of atoms or molecules that are vibrating. If a substance has a lighter molar mass, then each gram of it has more atoms or molecules available to store heat energy. This is why hydrogen—the lightest substance there is—has such a high specific heat capacity on a gram basis; one gram of it contains a relatively great many molecules. If specific heat capacity is measured on a molar basis (uppercase C), the differences between substances is less pronounced and hydrogen’s molar heat capacity is quite unremarkable. Conversely, for molecular-based substances (which also absorb heat into their internal degrees of freedom), massive, complex molecules with high atomic count — like gasoline — can store a great deal of energy per mole and yet, be quite unremarkable on a mass basis.
Since the bulk density of a solid chemical element is strongly related to its molar mass, generally speaking, there is a strong, inverse correlation between a solid’s density and its cp (constant-pressure specific heat capacity on a mass basis). Large ingots of low-density solids tend to absorb more heat energy than a small, dense ingot of the same mass because the former comprises more atoms. Thus, generally speaking, there a close correlation between the size of a solid chemical element and its total heat capacity (see Volumetric heat capacity). There are however, many departures from the general trend. For instance, arsenic, which is only 14.5% less dense than antimony, has nearly 59% more specific heat capacity on a mass basis. In other words; even though an ingot of arsenic is only about 17% larger than an antimony one of the same mass, it absorbs about 59% more heat energy for a given temperature rise.

Hydrogen bonds: Hydrogen-containing polar molecules like ethanol, ammonia, and water have powerful, intermolecular hydrogen bonds when in their liquid phase. These bonds provide yet another place where kinetic (heat) energy is stored.

[edit] Basic Equations
The equation relating heat energy to specific heat capacity, where the unit quantity is in terms of mass is:
Q = m c ΔT
where Q is the heat energy put into or taken out of the substance, m is the mass of the substance, c is the specific heat capacity, and ΔT is the temperature differential.
Where the unit quantity is in terms of moles, the equation relating heat energy to specific heat capacity (also known as molar heat capacity) is
Q = n C ΔT
where Q is the heat energy put into or taken out of the substance, n is the number of moles, C is the specific heat capacity, and ΔT is the temperature differential.

[edit] Table of specific heat capacities
Substance Phase cp
J g−1 K−1 Cp
J mol−1 K−1 Cv
J mol−1 K−1
Air (Sea level, dry, 0 °C) gas 1.0035 29.07
Air (typical room conditionsA) gas 1.012 29.19
Aluminium solid 0.897 24.2
Ammonia liquid 4.700 80.08
Antimony solid 0.207 25.2
Argon gas 0.5203 20.7862 12.4717
Arsenic solid 0.328 24.6
Beryllium solid 1.82 16.4
Copper solid 0.385 24.47
Diamond solid 0.5091 6.115
Ethanol liquid 2.44 112
Gasoline liquid 2.22 228
Gold solid 0.1291 25.42
Graphite solid 0.710 8.53
Helium gas 5.1932 20.7862 12.4717
Hydrogen gas 14.30 28.82
Iron solid 0.450 25.1
Lead solid 0.127 26.4
Lithium solid 3.58 24.8
Magnesium solid 1.02 24.9
Mercury liquid 0.1395 27.98
Nitrogen gas 1.040 29.12 20.8
Neon gas 1.0301 20.7862 12.4717
Oxygen gas 0.918 29.38
Silica (fused) solid 0.703 42.2
Uranium solid 0.116 27.7
Water gas (100 °C) 2.080 37.47 28.03
liquid (25 °C) 4.1813 75.327 74.53
solid (0 °C) 2.114 38.09
All measurements are at 25 °C unless otherwise noted.
Notable minima and maxima are shown in maroon.

A Assuming an altitude of 194 meters above mean sea level (the world–wide median altitude of human habitation), an indoor temperature of 23 °C, a dewpoint of 9 °C (40.85% relative humidity), and 760 mm–Hg sea level–corrected barometric pressure (molar water vapor content = 1.16%).

[edit] Specific heat capacity of building materials
Usually of interest to builders and solar designers

Substance Phase cp
J g−1 K−1
Asphalt solid 0.92
Brick solid 0.84
Concrete solid 0.88
Glass, silica solid 0.84
Glass, crown solid 0.67
Glass, flint solid 0.503
Glass, pyrex solid 0.753
Granite solid 0.790
Gypsum solid 1.09
Marble, mica solid 0.880
Sand solid 0.835
Soil solid 0.80
Wood solid 0.42

[edit] Derivations of heat capacity and specific heat capacity

[edit] Definition of heat capacity
Heat capacity is mathematically defined as the ratio of a small amount of heat δQ added to the body, to the corresponding small increase in its temperature dT:

For thermodynamic systems with more than one physical dimension, the above definition does not give a single, unique quantity unless a particular infinitesimal path through the system’s phase space has been defined (this means that one needs to know at all times where all parts of the system are, how much mass they have, and how fast they are moving). This information is used to account for different ways that heat can be stored as kinetic energy (energy of motion) and potential energy (energy stored in force fields), as an object expands or contracts. For all real systems, the path though these changes must be explicitly defined, since the value of heat capacity depends on which path from one temperature to another, is chosen. Of particular usefulness in this context are the values of heat capacity for constant volume, CV, and constant pressure, CP. These will be defined below.

[edit] Heat capacity of compressible bodies
The state of a simple compressible body with fixed mass is described by two thermodynamic parameters such as temperature T and pressure p. Therefore as mentioned above, one may distinguish between heat capacity at constant volume, CV, and heat capacity at constant pressure, Cp:

where

δQ is the infinitesimal amount of heat added,
dT is the subsequent rise in temperature.
The increment of internal energy is the heat added and the work added:

So the heat capacity at constant volume is

The enthalpy is defined by H = U + PV. The increment of enthalpy is

which, after replacing dU with the equation above and cancelling the PdV terms reduces to:

So the heat capacity at constant pressure is

Note that this last “definition” is a bit circular, since the concept of “enthalpy” itself was invented to be a measure of heat absorbed or produced at constant pressures (the conditions in which chemists usually work). As such, enthalpy merely accounts for the extra heat which is produced or absorbed by pressure-volume work at constant pressure. Thus, it is not surprising that constant-pressure heat capacities may be defined in terms of enthalpy, since “enthalpy” was defined in the first place to make this so.

[edit] Specific heat capacity
The specific heat capacity of a material is,

which in the absence of phase transitions is equivalent to,

where,

C is the heat capacity of a body made of the material in question,
m is the mass of the body,
V is the volume of the body, and
is the density of the material.
For gases, and also for other materials under high pressures, there is need to distinguish between different boundary conditions for the processes under consideration (since values differ significantly between different conditions). Typical processes for which a heat capacity may be defined include isobaric (constant pressure, dp = 0) or isochoric (constant volume, dV = 0) processes. The corresponding specific heat capacities are expressed as:

A related parameter to c is , the volumetric heat capacity. In engineering practice, for solids or liquids often signifies a volumetric heat capacity, rather than a constant-volume one. In such cases, the mass-specific heat capacity (specific heat) is often explicitly written with the subscript m, as . Of course, from the above relationships, for solids one writes:

[edit] Dimensionless heat capacity
The dimensionless heat capacity of a material is

where

C is the heat capacity of a body made of the material in question (J·K−1)
n is the amount of matter in the body (mol)
R is the gas constant (J·K−1·mol−1)
nR=Nk is the amount of matter in the body (J·K−1)
N is the number of molecules in the body. (dimensionless)
k is Boltzmann’s constant (J·K−1·molecule−1)
Again, SI units shown for example.

[edit] Theoretical models

[edit] Gas phase
The specific heat of the gas is best conceptualized in terms of the degrees of freedom of an individual molecule. The different degrees of freedom correspond to the different ways in which the molecule may store energy. The molecule may store energy in its translational motion according to the familiar formula

where m is the mass of the molecule and [vx,vy,vz] is velocity of the center of mass of the molecule. Each direction of motion constitutes a degree of freedom, so that there are three translational degrees of freedom.

In addition, a molecule may have rotational motion. The kinetic energy of rotational motion is generally expressed as

where I is the moment of inertia tensor of the molecule, and [ω1,ω2,ω3] is the angular velocity pseudovector (in a coordinate system aligned with the principle axes of the molecule). In general, then, there will be three additional degrees of freedom corresponding to the rotational motion of the molecule, (For linear molecules one of the inertia tensor terms vanishes and there are only two rotational degrees of freedom). The degrees of freedom corresponding to translations and rotations are called the “rigid” degrees of freedom, since they do not involve any deformation of the molecule.

The motions of the atoms in a molecule which are not part of its gross translational motion or rotation may be classified as vibrational motions. It can be shown that if there are n atoms in the molecule, there will be as many as 3n − 3 − nr vibrational degrees of freedom, where nr is the number of rotational degrees of freedom. The actual number may be less due to various symmetries.

If the molecule could be entirely described using classical mechanics, then we could use the theorem of equipartition of energy to predict that each degree of freedom would have an average energy in the amount of (1/2)kT where k is Boltzmann’s constant and T is the temperature. Our calculation of the heat content would be straightforward. Each molecule would be holding, on average, an energy of (f/2)kT where f is the total number of degrees of freedom in the molecule. The total internal energy of the gas would be (f/2)NkT where N is the total number of molecules. The heat capacity (at constant volume) would then be a constant (f/2)Nk , the specific heat capacity would be (f/2)k and the dimensionless heat capacity would be just f/2.

The various degrees of freedom cannot generally be considered to obey classical mechanics. Classically, the energy residing in each degree of freedom is assumed to be continuous - it can take on any positive value, depending on the temperature. In reality, the amount of energy that may reside in a particular degree of freedom is quantized: It may only be increased and decreased in finite amounts. A good estimate of the size of this minimum amount is the energy of the first excited state of that degree of freedom above its ground state. For example, the first vibrational state of the HCl molecule has an energy of about 5.74 × 10–20 joule. If this amount of energy were deposited in a classical degree of freedom, it would correspond to a temperature of about 4156 K.

If the temperature of the substance is so low that the equipartition energy of (1/2)kT is much smaller than this excitation energy, then there will be little or no energy in this degree of freedom. This degree of freedom is then said to be “frozen out". As mentioned above, the temperature corresponding to the first excited vibrational state of HCl is about 4156 K. For temperatures well below this value, the vibrational degrees of freedom of the HCL molecule will be frozen out. They will contain little energy and will not contribute to the heat content of the HCl gas.

It can be seen that for each degree of freedom there is a critical temperature at which the degree of freedom “unfreezes” and begins to accept energy in a classical way. In the case of translational degrees of freedom, this temperature is that temperature at which the thermal wavelength of the molecules is roughly equal to the size of the container. For a container of macroscopic size (e.g. 10 cm) this temperature is extremely small and has no significance, since the gas will certainly liquify or freeze before this low temperature is reached. For any real gas we may consider translational degrees of freedom to always be classical and contain an average energy of (3/2)kT per molecule.

The rotational degrees of freedom are the next to “unfreeze". In a diatomic gas, for example, the critical temperature for this transition is usually a few tens of kelvins. Finally, the vibrational degrees of freedom are generally the last to unfreeze. As an example, for diatomic gases, the critical temperature for the vibrational motion is usually a few thousands of kelvins.

It should be noted that it has been assumed that atoms have no rotational or internal degrees of freedom. This is in fact untrue. For example, atomic electrons can exist in excited states and even the atomic nucleus can have excited states as well. Each of these internal degrees of freedom are assumed to be frozen out due to their relatively high excitation energy. Nevertheless, for sufficiently high temperatures, these degrees of freedom cannot be ignored.

[edit] Monatomic gas
In the case of a monatomic gas such as helium under constant volume, if it assumed that no electronic or nuclear quantum excitations occur, each atom in the gas has only 3 degrees of freedom, all of a translational type. No energy dependence is associated with the degrees of freedom which define the position of the atoms. While, in fact, the degrees of freedom corresponding to the momenta of the atoms are quadratic, and thus contribute to the heat capacity. There are N atoms, each of which has 3 components of momentum, which leads to 3N total degrees of freedom. This gives:

where

CV is the heat capacity at constant volume of the gas
CV,m is the molar heat capacity at constant volume of the gas
N is the total number of atoms present in the container
n is the number of moles of atoms present in the container (n is the ratio of N and Avogadro’s number)
R is the ideal gas constant, (8.314570[70] J K−1mol−1). R is equal to the product of Boltzmann’s constant kB and Avogadro’s number
The following table shows experimental molar constant volume heat capacity measurements taken for each noble monatomic gas (at 1 atm and 25 °C):

Monatomic gas CV, m (J K−1 mol−1) CV, m/R
He 12.5 1.50
Ne 12.5 1.50
Ar 12.5 1.50
Kr 12.5 1.50
Xe 12.5 1.50

It is apparent from the table that the experimental heat capacities of the monatomic noble gases agrees with this simple application of statistical mechanics to a very high degree.

[edit] Diatomic gas
In the somewhat more complex case of an ideal gas of diatomic molecules, the presence of internal degrees of freedom are apparent. In addition to the three translational degrees of freedom, there are rotational and vibrational degrees of freedom. In general, the number of degrees of freedom, f, in a molecule with na atoms is 3na:

Mathematically, there are a total of three rotational degrees of freedom, one corresponding to rotation about each of the axes of three dimensional space. However, in practice we shall only consider the existence of two degrees of rotational freedom for linear molecules. This approximation is valid because the moment of inertia about the internuclear axis is vanishingly small with respect other moments of inertia in the molecule (this is due to the extremely small radii of the atomic nuclei, compared to the distance between them in a molecule). Quantum mechanically, it can be shown that the interval between successive rotational energy eigenstates is inversely proportional to the moment of inertia about that axis. Because the moment of inertia about the internuclear axis is vanishingly small relative to the other two rotational axes, the energy spacing can be considered so high that no excitations of the rotational state can possibly occur unless the temperature is extremely high. We can easily calculate the expected number of vibrational degrees of freedom (or vibrational modes). There are three degrees of translational freedom, and two degrees of rotational freedom, therefore

Each rotational and translational degree of freedom will contribute R/2 in the total molar heat capacity of the gas. Each vibrational mode will contribute R to the total molar heat capacity, however. This is because for each vibrational mode, there is a potential and kinetic energy component. Both the potential and kinetic components will contribute R/2 to the total molar heat capacity of the gas. Therefore, we expect that a diatomic molecule would have a molar constant-volume heat capacity of

where the terms originate from the translational, rotational, and vibrational degrees of freedom, respectively.

The following is a table of some molar constant-volume heat capacities of various diatomic gasses

Diatomic gas CV, m (J K−1 mol−1) CV, m / R
H2 20.18 2.427
CO 20.2 2.43
N2 19.9 2.39
Cl2 24.1 2.90
Br2 32.0 3.84

From the above table, clearly there is a problem with the above theory. All of the diatomics examined have heat capacities that are lower than those predicted by the Equipartition Theorem, except Br2. However, as the atoms composing the molecules become heavier, the heat capacities move closer to their expected values. One of the reasons for this phenomenon is the quantization of vibrational, and to a lesser extent, rotational states. In fact, if it is assumed that the molecules remain in their lowest energy vibrational state because the inter-level energy spacings are large, the predicted molar constant volume heat capacity for a diatomic molecule becomes

which is a fairly close approximation of the heat capacities of the lighter molecules in the above table. If the quantum harmonic oscillator approximation is made, it turns out that the quantum vibrational energy level spacings are actually inversely proportional to the square root of the reduced mass of the atoms composing the diatomic molecule. Therefore, in the case of the heavier diatomic molecules, the quantum vibrational energy level spacings become finer, which allows more excitations into higher vibrational levels at a fixed temperature.

[edit] Solid phase

The dimensionless heat capacity divided by three, as a function of temperature as predicted by the Debye model and by Einstein’s earlier model. The horizontal axis is the temperature divided by the Debye temperature. Note that, as expected, the dimensionless heat capacity is zero at absolute zero, and rises to a value of three as the temperature becomes much larger than the Debye temperature. The red line corresponds to the classical limit of the Dulong-Petit lawFor matter in a crystalline solid phase, the Dulong-Petit law, which was discovered empirically, states that the dimensionless specific heat capacity assumes the value 3R. Indeed, for solid metallic chemical elements at room temperature, heat capacities range from about 2.8 to 3.4 (beryllium being a notable exception at 2.0).

The theoretical maximum heat capacity for larger and larger multi-atomic gases at higher temperatures, also approaches the Dulong-Petit limit of 3R, so long as this is calculated per mole of atoms, not molecules. The reason is that gases with very large molecules, in theory have almost the same high-temperature heat capacity as solids, lacking only the (small) heat capacity contribution that comes from potential energy that cannot be stored between separate molecules in a gas.

The Dulong-Petit “limit” results from the equipartition theorem, and as such is only valid in the classical limit of a microstate continuum, which is a high temperature limit. For light and non-metallic elements, as well as most of the common molecular solids based on carbon compounds at standard ambient temperature, quantum effects may also play an important role, as they do in multi-atomic gases. These effects usually combine to give heat capacities lower than 3 R per mole of atoms in the solid, although heat capacities calculated per mole of molecules in molecular solids may be more than 3 R. For example, the heat capacity of water ice at the melting point is about 4.6 R per mole of molecules, but only 1.5 R per mole of atoms. The lower number results from the “freezing out” of possible vibration modes for light atoms at suitably low temperatures, just as in many gases. These effects are seen in solids more often than liquids: for example the heat capacity of liquid water is again close to the theoretical 3 R per mole of atoms of the Dulong-Petit theoretical maximum.

For a more modern and precise analysis of the heat capacities of solids, especially at low temperatures, it is useful to use the idea of phonons. See Debye model.

[edit] Heat capacity at absolute zero
From the definition of entropy

we can calculate the absolute entropy by integrating from zero temperature to the final temperature Tf

The heat capacity must be zero at zero temperature in order for the above integral not to yield an infinite absolute entropy, thus violating the third law of thermodynamics. One of the strengths of the Debye model is that (unlike the preceding Einstein model) it predicts an approach of heat capacity toward zero as zero temperature is approached, and also predicts the proper mathematical form of this approach.

[edit] See also
Heat
Heat capacity ratio
Heat equation
Heat transfer coefficient
Latent heat
Specific melting heat
Specific heat of vaporization
Temperature
Thermodynamic (absolute) temperature
Volumetric heat capacity

[edit] References
^ IUPAC.org, Gold Book, Standard Pressure
^ The comparison must be made under constant-volume conditions — CvH — so that no work is performed. Nitrogen’s CvH (100 kPa, 20 °C) = 20.8 J mol–1 K–1 vs. the monatomic gases which equal 12.4717 J mol–1 K–1. Citations: W.H. Freeman’s Physical Chemistry, Part 3: Change (422 kB PDF, here), Exercise 21.20b, Pg. 787. Also Georgia State University’s Molar Specific Heats of Gases.
Retrieved from "http://en.wikipedia.org/wiki/Specific_heat_capacity#Heat_capacity"
Categories: Chemical properties | Physical quantity | Thermodynamics | Heat | Fundamental physics concepts

Enthalpy of vaporization

2007-08-14T10:21:00.000-07:00

Enthalpy of vaporization
From Wikipedia, the free encyclopedia
Jump to: navigation, search

Molar heat content of zinc above 298.15 K and at 1 atm pressure, showing discontinuities at the melting and boiling points. The enthalpy of melting (ΔH°m) of zinc is 7323 J/mol, and the enthalpy of vaporization (ΔH°v) is 115 330 J/mol.The enthalpy of vaporization, (symbol ΔvH), also known as the heat of vaporization or heat of evaporation, is the energy required to transform a given quantity of a substance into a gas. It is measured at the boiling point of the substance, although tabulated values are usually corrected to 298 K: the correction is small, and is often smaller than the uncertainty in the measured value. Values are usually quoted in kJ/mol, although kJ/kg, kcal/mol, cal/g and Btu/lb are also possible, among others.

The enthalpy of condensation (or heat of condensation) is numerically exactly equal to the enthalpy of vaporization, but has the opposite sign: enthalpy changes of vaporization are always positive (heat is absorbed by the substance), whereas enthalpy changes of condensation are always negative (heat is released by the substance).

The enthalpy of vaporization can be viewed as the energy required to overcome the intermolecular interactions in the liquid (or solid, in the case of sublimation). Hence helium has a particularly low enthalpy of vaporization, 0.0845 kJ/mol, as the van der Waals forces between helium atoms are particularly weak. On the other hand, the molecules in liquid water are held together by relatively strong hydrogen bonds, and its enthalpy of vaporization, 40.8 kJ/mol, is more than five times the energy required to heat the same quantity of water from 0 °C to 100 °C (cp = 75.3 J K−1 mol−1). Care must be taken, however, when using enthalpies of vaporization to measure the strength of intermolecular forces, as these forces may persist to an extent in the gas phase (as is the case with hydrogen fluoride), and so the calculated value of the bond strength will be too low. This is particularly true of metals, which often form covalently bonded molecules in the gas phase: in these cases, the enthalpy of atomization must be used to obtain a true value of the bond energy.

An alternative description is to view the enthalpy of condensation as the heat which must be released to the surroundings to compensate for the drop in entropy when a gas condenses to a liquid. As the liquid and gas are in equilibrium at the boiling point (Tb), ΔvG = 0, which leads to:

As neither entropy nor enthalpy vary greatly with temperature, it is normal to use the tabulated standard values without any correction for the difference in temperature from 298 K. A correction must be made if the pressure is different from 100 kPa, as the entropy of a gas is proportional to its pressure (or, more precisely, to its fugacity): the entropies of liquids vary little with pressure, as the compressibility of a liquid is small.

These two definitions are equivalent: the boiling point is the temperature at which the increased entropy of the gas phase overcomes the intermolecular forces. As a given quantity of matter always has a higher entropy in the gas phase than in a condensed phase ( is always positive), and from

,
the Gibbs free energy change falls with increasing temperature: gases are favored at higher temperatures, as is observed in practice.

Contents [hide]
1 Selected values
1.1 Elements
1.2 Other common substances
2 See also
3 References

[edit] Selected values

[edit] Elements
Enthalpies of vaporization of the elements in kJ/mol

Group → 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
↓ Period

1 H
0.44936
He
0.0845
2 Li
145.92 Be
292.40
B
489.7 C
355.8 N
2.7928 O
3.4099 F
3.2698 Ne
1.7326
3 Na
96.96 Mg
127.4
Al
293.4 Si
384.22 P
12.129 S
1.7175 Cl
10.2 Ar
6.447
4 K
79.87 Ca
153.6 Sc
314.2 Ti
421 V
452 Cr
344.3 Mn
226 Fe
349.6 Co
376.5 Ni
370.4 Cu
300.3 Zn
115.3 Ga
258.7 Ge
330.9 As
34.76 Se
26.3 Br
15.438 Kr
9.029
5 Rb
72.216 Sr
144 Y
363 Zr
581.6 Nb
696.6 Mo
598 Tc
660 Ru
595 Rh
493 Pd
357 Ag
250.58 Cd
100 In
231.5 Sn
295.8 Sb
77.14 Te
52.55 I
20.752 Xe
12.636
6 Cs
67.74 Ba
142 *
Hf
575 Ta
743 W
824 Re
715 Os
627.6 Ir
604 Pt
510 Au
334.4 Hg
59.229 Tl
164.1 Pb
177.7 Bi
104.8 Po
60.1 At
114 Rn
16.4
7 Fr
n/a Ra
37 **
Rf
n/a Db
n/a Sg
n/a Bh
n/a Hs
n/a Mt
n/a Ds
n/a Rg
n/a Uub
n/a Uut
n/a Uuq
n/a Uup
n/a Uuh
n/a Uus
n/a Uuo
n/a

* Lanthanides La
414 Ce
414 Pr
n/a Nd
n/a Pm
n/a Sm
n/a Eu
n/a Gd
n/a Tb
n/a Dy
n/a Ho
n/a Er
n/a Tm
n/a Yb
n/a Lu
n/a
** Actinides Ac
n/a Th
514.4 Pa
n/a U
n/a Np
n/a Pu
n/a Am
n/a Cm
n/a Bk
n/a Cf
n/a Es
n/a Fm
n/a Md
n/a No
n/a Lr
n/a
0–10 kJ/mol 10–100 kJ/mol 100–300 kJ/mol >300 kJ/mol

[edit] Other common substances
Common substances sorted by heat of vaporization:

Compound Heat of vaporization (kJ/mol)
Water 40.65 (540 calories per gram)
Ethanol 38.6
Methanol 37.4
Ammonia 23.35
Butane 21.0 (362 kJ/kg)
Propane 15.7 (356 kJ/kg)
Phosphine 14.6
Methane 8.19
Compound Heat of vaporization (kJ/mol)

[edit] See also
Enthalpy of fusion

[edit] References
Sears, Zemansky et al., University Physics, Addison-Wessley Publishing Company, Sixth ed., 1982, ISBN 0-201-07199-1

Retrieved from "http://en.wikipedia.org/wiki/Enthalpy_of_vaporization"
Categories: Chemical properties | Thermodynamics | Heat

Enthalpy of fusion

2007-08-14T10:19:00.000-07:00

Enthalpy of fusion
From Wikipedia, the free encyclopedia
Jump to: navigation, search

Molar heat content of zinc above 298.15 K and at 1 atm pressure, showing discontinuities at the melting and boiling points. The enthalpy of melting (ΔH°m) of zinc is 7323 J/mol, and the enthalpy of vaporization (ΔH°v) is 115 330 J/mol.The enthalpy of fusion (symbol: ΔfusH), also known as the heat of fusion or specific melting heat, is the amount of thermal energy which must be absorbed or evolved for 1 mole of a substance to change states from a solid to a liquid or vice versa. It is also called the latent heat of fusion or the enthalpy change of fusion, and the temperature at which it occurs is called the melting point.

When you withdraw thermal energy from a liquid or solid, the temperature falls. When you add heat energy the temperature rises. However, at the transition point between solid and liquid (the melting point), extra energy is required (the heat of fusion). To go from liquid to solid, the molecules of a substance must become more ordered. For them to maintain the order of a solid, extra heat must be withdrawn. In the other direction, to create the disorder from the solid crystal to liquid, extra heat must be added.

The heat of fusion can be observed if you measure the temperature of water as it freezes. If you plunge a closed container of room temperature water into a very cold environment (say −20 °C), you will see the temperature fall steadily until it drops just below the freezing point (0 °C). The temperature then rebounds and holds steady while the water crystallizes. Once completely frozen, the temperature will fall steadily again.

The temperature stops falling at (or just below) the freezing point due to the heat of fusion. The energy of the heat of fusion must be withdrawn (the liquid must turn to solid) before the temperature can continue to fall.

The units of heat of fusion are usually expressed as:

joules per mole (the SI units)
calories per gram (old metric units now little used, except for a different, larger calorie used in nutritional contexts)
British thermal units per pound or Btu per pound-mole
Note: These are not the calories found in food. The calories found in food are more properly known as kilocalories—equal to 1000 calories. 1000 calories = 1 kilocalorie = 1 food calorie. Food calories are sometimes abbreviated as kcal as if small calories were being used, while calories are abbreviated as cal. Another distinguishing method, though often confusing, uses capitalisation. A Calorie is a food calorie, or 1000 calories. So 1 Cal = 1000 cal.
Contents [hide]
1 Reference Values of Common Substances
2 Applications
3 Solubility prediction
3.1 Proof
4 See also
5 References

[edit] Reference Values of Common Substances

Standard enthalpy change of fusion of period three.
Standard enthalpy change of fusion of period two of the periodic table of elements.Substance Heat of fusion
(cal/g) Heat of fusion
(kJ/kg)
water 79.72 333.55
methane 13.96 58.41
ethane 22.73 95.10
propane 19.11 79.96
methanol 23.70 99.16
ethanol 26.05 108.99
glycerol 47.95 200.62
formic acid 66.05 276.35
acetic acid 45.91 192.09
acetone 23.42 97.99
benzene 30.45 127.40
myristic acid 47.49 198.70
palmitic acid 39.18 163.93
stearic acid 47.54 198.91

These values are from the CRC Handbook of Chemistry and Physics, 62nd edition. The conversion between cal/g and kJ/kg in the above table uses the thermochemical calorie (calth) = 4.184 Joules rather than the International Steam Table calorie (calINT) = 4.1868 Joules.

[edit] Applications
To heat one kilogram (about 1 litre) of water from 10 °C to 30 °C requires 20 kcal.
However, to melt ice and raise the resulting water temperature 20 °C requires extra energy. To heat ice from 0 °C to water at 20 °C requires:

(1) 80 Cal/g (heat of fusion of ice) = 80 kcal for 1 kg
PLUS
(2) 1 cal/(g·°C) = 20 kcal for 1 kg to go up 20 °C
= 100 kcal

[edit] Solubility prediction
The heat of fusion can also be used to predict solubility for solids in liquids. Provided an ideal solution is obtained the mole fraction of solute at saturation is a function of the heat of fusion, the melting point of the solid and the temperature of the solution:

For example the solubility of paracetamol in water at 298 K is predicted to be:

This equals to a solubility in grams per liter of:

which is a deviation from the real solubility (240 g/L) of 11%. This error can be reduced when an additional heat capacity parameter is taken into account [1]

[edit] Proof
At equilibrium the chemical potentials for the pure solvent and pure solid are identical:

or

with the gas constant and the temperature.

Rearranging gives:

and since

the heat of fusion being the difference in chemical potential between the pure liquid and the pure solid, it follows that

Application of the Gibbs-Helmholtz equation:

ultimately gives:

or:

and with integration:

the end result is obtained:

[edit] See also
Heat of vaporization
Heat capacity
Specific heat capacity
Thermodynamic databases for pure substances

[edit] References
^ Measurement and Prediction of Solubility of Paracetamol in Water-Isopropanol Solution. Part 2. Prediction H. Hojjati and S. Rohani Org. Process Res. Dev.; 2006; 10(6) pp 1110 - 1118; (Article) DOI:10.1021/op060074g
Retrieved from "http://en.wikipedia.org/wiki/Enthalpy_of_fusion"
Category: Chemical properties

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
Afrikaans
Asturianu
Català
Česky
Deutsch
Español
Français
한국어
Lietuvių
Lojban
Nederlands
Polski
Português
Slovenščina
Српски / Srpski
Svenska
ไทย
Tiếng Việt

This page was last modified 11:56, 14 August 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Critical point

2007-08-14T10:18:00.000-07:00

Critical point (thermodynamics)
From Wikipedia, the free encyclopedia
Jump to: navigation, search
In physical chemistry, thermodynamics, chemistry and condensed matter physics, a critical point, also called a critical state, specifies the conditions (temperature, pressure) at which the liquid state of the matter ceases to exist. As a liquid is heated, its density decreases while the pressure and density of the vapor being formed increases. The liquid and vapor densities become closer and closer to each other until the critical temperature is reached where the two densities are equal and the liquid-gas line or phase boundary disappears. Additionally, as the equilibrium between liquid and gas approaches the critical point, heat of vaporization approaches zero, becoming zero at and beyond the critical point. More generally, the critical point is the point of termination of a phase equilibrium curve, which separates two distinct phases. At this point, the phases are no longer distinguishable.

The critical point in a phase diagram is at the high-temperature extreme of the liquid-gas phase boundary.In the phase diagram shown, the phase boundary between liquid and gas does not continue indefinitely. Instead, it terminates at a point on the phase diagram called the critical point. This reflects the fact that, at extremely high temperatures and pressures, the liquid and gaseous phases become indistinguishable. In water, the critical point occurs at around 647 K (374 °C or 705 °F) and 22.064 MPa (3200 PSIA or 218atm).

Critical variables are useful for rewriting a varied equation of state into one that applies to all materials. The effect is similar to a normalizing constant.

According to renormalization group theory, the defining property of criticality is that the natural length scale characteristic of the structure of the physical system, the so-called correlation length ξ, becomes infinite. There are also lines in phase space along which this happens: these are critical lines.

In equilibrium systems the critical point is reached only by tuning a control parameter precisely. However, in some non-equilibrium systems the critical point is an attractor of the dynamics in a manner that is robust with respect to system parameters, a phenomenon referred to as self-organized criticality.

The critical point is described by a conformal field theory.

[edit] See also
Critical temperature
Phase transition
Scale invariance
Conformal field theory
Critical exponents
Percolation thresholds
Self-organized criticality
Triple point
Supercritical fluid, Supercritical drying, Supercritical water oxidation
Rushbrooke inequality
Widom scaling

[edit] External link
Critical points for some common solvents

[hide]v • d • eStates of Matter (list)
Solid • Liquid • Gas • Plasma • Supercritical fluid • Superfluid • Supersolid • Degenerate matter • Quark-gluon plasma • Fermionic condensate • Bose–Einstein condensate • Strange matter • Melting point • Boiling point • Triple point • Critical point • Equation of state • Cooling curve

Retrieved from "http://en.wikipedia.org/wiki/Critical_point_%28thermodynamics%29"
Categories: Phase changes | Statistical mechanics | Conformal field theory | Renormalization group | Condensed matter physics

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
Asturianu
Català
Česky
Deutsch
Español
Français
Italiano
Nederlands
日本語
‪Norsk (bokmål)‬
Plattdüütsch
Polski
Português
Русский
Slovenščina
Suomi
Svenska

This page was last modified 23:34, 17 June 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Boiling point

2007-08-14T10:17:00.000-07:00

Boiling point
From Wikipedia, the free encyclopedia
Jump to: navigation, search
This article is about the boiling point of liquids. For other uses, see Boiling point (disambiguation).
The boiling point of a substance is the maximum temperature at which a liquid can remain a liquid. Adding a small amount of heat energy (latent heat of vaporization) can convert the liquid into a gas. A pure liquid may change to a gas at temperatures below the boiling point through the process of evaporation. Any change of state from a liquid to a gas at boiling point is considered vaporization. However, evaporation is a surface phenomenon, in which only molecules located near the gas/liquid surface could evaporate. Boiling on the other hand is a bulk process, so at the boiling point molecules anywhere in the liquid may be vaporized, resulting in the formation of vapor bubbles.

A somewhat clearer (and perhaps more useful) definition of boiling point is "the temperature at which the vapor pressure of the liquid equals the atmospheric pressure."

Contents [hide]
1 Saturation temperature and pressure
2 Intermolecular interactions
3 Properties of other elements
4 See also

[edit] Saturation temperature and pressure
A saturated liquid contains as much thermal energy as it can without boiling (or conversely a saturated vapor contains as little thermal energy as it can without condensing).

Saturation temperature means boiling point. The saturation temperature is the temperature for a corresponding saturation pressure at which a liquid boils into its vapor phase. The liquid can be said to be saturated with thermal energy. Any addition of thermal energy results in a phase change.

If the pressure in a system remains constant (isobaric), a vapor at saturation temperature will begin to condense into its liquid phase as thermal energy (heat) is removed. Similarly, a liquid at saturation temperature and pressure will boil into its vapor phase as additional thermal energy is applied.

The boiling point corresponds to the temperature at which the vapor pressure of the substance equals the ambient pressure. Thus the boiling point is dependent on the pressure. Usually, boiling points are published with respect to standard pressure (101.325 kilopascals or 1 atm). At higher elevations, where the atmospheric pressure is much lower, the boiling point is also lower. The boiling point increases with increased ambient pressure up to the critical point, where the gas and liquid properties become identical. The boiling point cannot be increased beyond the critical point. Like wise, the boiling point decreases with decreasing ambient pressure until the triple point is reached. The boiling point cannot be reduced below the triple point.

If the Heat of Vaporization and the vapor pressure of a substance at a certain temperature is known, the normal boiling point (under standard pressure) can be calculated by:

Where TB is the boiling point under standard pressure, R is the ideal gas constant, P0 is the vapor pressure at a given temperature, T0 is that temperature, and ΔHvap is the heat of vaporization of the substance.

Saturation Pressure, or vapor point, is the pressure for a corresponding saturation temperature at which a liquid boils into its vapor phase. Saturation pressure and saturation temperature have a direct relationship: as saturation pressure is increased so is saturation temperature.

If the temperature in a system remains constant (an isothermal system), vapor at saturation pressure and temperature will begin to condense into its liquid phase as the system pressure is increased. Similarly, a liquid at saturation pressure and temperature will tend to flash into its vapor phase as system pressure is decreased.

[edit] Intermolecular interactions
In terms of intermolecular interactions, the boiling point represents the point at which the liquid molecules possess enough thermal energy to overcome the various intermolecular attractions binding the molecules into the liquid (eg. dipole-dipole attraction, instantaneous-dipole induced-dipole attractions, and hydrogen bonds). Therefore the boiling point is also an indicator of the strength of these attractive forces.

The boiling point of water is 100 °C (212 °F) at standard pressure. On top of Mount Everest the pressure is about 260 mbar (26 kPa) so the boiling point of water is 69 °C. (156.2 °F).

For purists with a knowledge of thermodynamics, the normal boiling point of water is 99.97 degrees Celsius (at a pressure of 1 atm, i.e. 101.325 kPa). Until 1982 this was also the standard boiling point of water, but the IUPAC now recommends a standard pressure of 1 bar (100 kPa). At this slightly reduced pressure, the standard boiling point of water is 99.61 degrees Celsius.

[edit] Properties of other elements
The element with the lowest boiling point is helium. Both the boiling points of rhenium and tungsten exceed 5000 K at standard pressure. Due to the experimental difficulty of precisely measuring extreme temperatures without bias, there is some discrepancy in the literature as to whether tungsten or rhenium has the higher boiling point. (Cf. DeVoe, Howard, Thermodynamics and Chemistry. Prentice-Hall, 2001)

[edit] See also
List of elements by boiling point
Leidenfrost effect
flash point
boiling delay
critical temperature
triple point
boiling-point elevation
Retrieved from "http://en.wikipedia.org/wiki/Boiling_point"
Categories: Thermodynamics | Fundamental physics concepts

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
Afrikaans
العربية
Asturianu
Български
Català
Česky
Dansk
Deutsch
Eesti
Español
Esperanto
Galego
हिन्दी
한국어
Bahasa Indonesia
Íslenska
Italiano
עברית
Latviešu
Lietuvių
Lojban
Magyar
Македонски
Bahasa Melayu
Nederlands
日本語
‪Norsk (bokmål)‬
‪Norsk (nynorsk)‬
Polski
Português
Română
Русский
Simple English
Slovenčina
Slovenščina
Српски / Srpski
Srpskohrvatski / Српскохрватски
Suomi
Svenska
தமிழ்
ไทย
Tiếng Việt
Türkçe
Українська
O'zbek
中文

This page was last modified 13:53, 13 August 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Melting point

2007-08-14T10:15:00.000-07:00

Melting point
From Wikipedia, the free encyclopedia
Jump to: navigation, search
It has been suggested that melting be merged into this article or section. (Discuss)
The melting point of a crystalline solid is the temperature range at which it changes state from solid to liquid. Although the phrase would suggest a specific temperature and is commonly and incorrectly used as such in most textbooks and literature, most crystalline compounds actually melt over a range of a few degrees or less. At the melting point the solid and liquid phase exist in equilibrium. When considered as the temperature of the reverse change from liquid to solid, it is referred to as the freezing point.

Contents [hide]
1 Fundamentals
2 Melting point measurements
3 Thermodynamics
4 Carnelley’s Rule
5 See also
6 References

[edit] Fundamentals

Melting points (in blue) and boiling points (in pink) of the first eight Carboxylic acids (°C)For most substances, melting and freezing points are equal. For example, the melting point and freezing point of the element mercury is 234.32 kelvin (−38.83 °C or −37.89 °F). However, certain substances possess differing solid-liquid transition temperatures. For example, agar melts at 85 °C (185 °F) and solidifies from 31 °C to 40 °C (89.6 °F to 104 °F); this process is known as hysteresis.

Certain materials, such as glass, may harden without crystallizing; these are called amorphous solids. Amorphous materials as well as some polymers do not have a true melting point as there is no abrupt phase change at any specific temperature. Instead, there is a gradual change in their viscoelastic properties over a range of temperatures. Such materials are characterized by a glass transition temperature which may be roughly defined as the "knee" point of the material's density vs. temperature graph.

The melting point of water at 1 atmosphere of pressure is very close [1] to 0 °C (32 °F, 273.15 K), this is also known as the ice point. In the presence of nucleating substances the freezing point of water is the same as the melting point, but in the absence of nucleators water can supercool to −42 °C (−43.6 °F, 231 K) before freezing.

Unlike the boiling point, the melting point is relatively insensitive to pressure. Melting points are often used to characterize organic compounds and to ascertain the purity. The melting point of a pure substance is always higher and has a smaller range than the melting point of an impure substance. The more impurity is present, the lower the melting point and the broader the range. Eventually, a minimum melting point will be reached. The mixing ratio that results in the lowest possible melting point is known as the eutectic point.

The chemical element with the highest melting point is tungsten, at 3695 K (3422 °C, 6192 °F) making it excellent for use as filaments in light bulbs. The often-cited carbon does not melt at ambient pressure but sublimates at about 4000 K; a liquid phase only exists above pressures of 10 MPa and estimated 4300–4700 K. Tantalum hafnium carbide (Ta4HfC5) is a refractory compound with a very high melting point of 4488 K (4215 °C, 7619 °F).[2] At the other end of the scale, helium does not freeze at all at normal pressure, even at temperatures infinitesimally close to absolute zero; pressures over 20 times normal atmospheric pressure are necessary.

[edit] Melting point measurements
Many Laboratory techniques exist for the determination of melting points. A Kofler bench is a metal strip with a temperature gradient (range room temperature to 300°C). Any substance can be place on a section of the strip revealing its thermal behaviour at the temperature at that point. Differential scanning calorimetry gives information on melting point together with its Enthalpy of fusion.

[edit] Thermodynamics

Pressure dependence of water melting point (MPa/K)Not only is heat required to raise the temperature of the solid to the melting point, but the melting itself requires heat called the heat of fusion.

From a thermodynamics point of view, at the melting point the change in Gibbs free energy (ΔG) of the material is zero, because the enthalpy (H) and the entropy (S) of the material are increasing (ΔH,ΔS > 0). Melting phenomenon happens when the Gibbs free energy of the liquid becomes lower than the solid for that material. At various pressures this happens at a specific temperature. It can also be shown that:

The "T","ΔS", and "ΔH" in the above are respectively the temperature at the melting point, change of entropy of melting, and the change of enthalpy of melting.

[edit] Carnelley’s Rule
In organic chemistry Carnelley’s Rule established in 1882 by Thomas Carnelley, states that high molecular symmetry is associated with high melting point [3]. Carnelley based his rule on examination of 15,000 chemical compounds. For example for three structural isomers with molecular formula C5H12 the melting point increases in the series isopentane −160 °C (113 K) n-pentane −129.8 °C (143 K) and neopentane −18 °C (255 K). Likewise in xylenes and also dichlorobenzenes the melting point increases in the order meta, ortho and then para. Pyridine has a lower symmetry than benzene hence its lower melting point but the melting point again increases with diazine and triazines. Many cage-like compounds like adamantane and cubane with high symmetry have very high melting points.

A high melting point results from a high heat of fusion or a low entropy of fusion or a combination. In highly symmetrical molecules the crystal phase is densely packed with many efficient intermolecular interactions resulting in a higher enthalpy change on melting.

[edit] See also
Melting
Phases of matter
Triple point
Freezing-point depression
Boiling point
Melting Points for various elements

[edit] References
^ The ice point of purified water has been measured to be 0.000089 +/- 0.00001 degrees Celsius - see Magnum, B.W. (June 1995). "Reproducibility of the Temperature of the Ice Point in Routine Measurements" (PDF). Nist Technical Note 1411. Retrieved on 2007-02-11.
^ hafnium entry at Britannica.com
^ Melting Point and Molecular Symmetry R. J. C. Brown, R. F. C. Brown Journal of Chemical Education 724 Vol. 77 No. 6 June 2000
Retrieved from "http://en.wikipedia.org/wiki/Melting_point"
Categories: Articles to be merged since July 2007 | Fundamental physics concepts | Chemical properties | Atmospheric thermodynamics | Phase changes

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
Afrikaans
العربية
Asturianu
Bosanski
Български
Català
Česky
Dansk
Deutsch
Eesti
Ελληνικά
Español
Esperanto
فارسی
Français
한국어
Hrvatski
Bahasa Indonesia
Íslenska
Italiano
עברית
Latviešu
Lietuvių
Lojban
Lumbaart
Magyar
Македонски
Bahasa Melayu
Nederlands
日本語
‪Norsk (bokmål)‬
‪Norsk (nynorsk)‬
Polski
Português
Română
Русский
Simple English
Slovenščina
Српски / Srpski
Srpskohrvatski / Српскохрватски
Suomi
Svenska
ไทย
Türkçe
Українська
O'zbek
粵語
中文

This page was last modified 17:43, 4 August 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Density

2007-08-14T10:13:00.000-07:00

Density
From Wikipedia, the free encyclopedia
Jump to: navigation, search
For other uses, see Density (disambiguation).
In physics, density is mass m per unit volume V. For the common case of a homogeneous substance, it is expressed as:

where, in SI units:

ρ (rho) is the density of the substance, measured in kg·m-3
m is the mass of the substance, measured in kg
V is the volume of the substance, measured in m3
Contents [hide]
1 History
2 Measurement of density
3 Common units
4 Changes of density
5 Density of water
6 Density of air
7 References
8 Books
9 See also
10 External links

[edit] History
In a famous problem, Archimedes was given the task of determining if King Hiero's goldsmith was embezzling gold during the manufacture of the king's crown and replacing it with another, cheaper alloy.[1]

Archimedes knew that the crown could be smashed into a cube or sphere, where the volume could be calculated more easily when compared with the weight; the king did not approve of this.

Baffled, Archimedes went to take a bath and observed from the rise of the water upon entering that he could calculate the volume of the crown by the displacement of the water. Allegedly, Archimedes went running though the streets naked shouting, "Eureka! Eureka!"

It is not known what happened to the goldsmith.

[edit] Measurement of density
For a homogeneous object, the formula mass/volume may be used. The mass is normally measured with an appropriate scale; the volume may be measured directly (from the geometry of the object) or by the displacement of a liquid. A very common instrument for the direct measurement of the density of a liquid is the hydrometer. A less common device for measuring fluid density is a pycnometer, a similar device for measuring the absolute density of a solid is a gas pycnometer.

The density of a solid material can be ambiguous, depending on exactly how it is defined, and this may cause confusion in measurement. A common example is sand: if gently filled into a container, the density will be small; when the same sand is compacted into the same container, it will occupy less volume and consequently carry a greater density. This is because "sand" contains a lot of air space in between individual grains; this overall density is called the bulk density, which differs significantly from the density of an individual grain of sand.

[edit] Common units
In U.S. customary units or Imperial units, the units of density include;

ounces per cubic inch (oz/in³)
pounds per cubic inch (lb/in³)
pounds per cubic foot (lb/ft³)
pounds per cubic yard (lb/yd³)
pounds per gallon (for U.S. or imperial gallons) (lb/gal)
pounds per U.S. bushel (lb/bu)
slugs per cubic foot.

[edit] Changes of density
In general density can be changed by changing either the pressure or the temperature. Increasing the pressure will always increase the density of a material. Increasing the temperature generally decreases the density, but there are notable exceptions to this generalisation. For example, the density of water increases between its melting point at 0 °C and 4 °C and similar behaviour is observed in silicon at low temperatures.

The effect of pressure and temperature on the densities of liquids and solids is small so that a typical compressibility for a liquid or solid is 10-6 bar-1 (1 bar=0.1 MPa) and a typical thermal expansivity is 10-5 K-1.

In contrast, the density of gases is strongly affected by pressure. Boyle's law says that the density of an ideal gas is given by

where R is the universal gas constant, P is the pressure, m the molar mass, and T the absolute temperature.

This means that a gas at 300 K and 1 bar will have its density doubled by increasing the pressure to 2 bar or by reducing the temperature to 150 K.

[edit] Density of water
Temperature Density[2] (at 1 atm)
°C °F kg/m³
0.0 32.0 999.8425
4.0 39.2 999.9750
15.0 59.0 999.1026
20.0 68.0 998.2071
25.0 77.0 997.0479
37.0 98.6 993.3316
50.0 122.0 988.04
100.0 212.0 958.3665

[edit] Density of air
T in °C ρ in kg/m³ (at 1 atm)
−10 1.341
−5 1.316
0 1.293
5 1.269
10 1.247
15 1.225
20 1.204
25 1.184
30 1.164

[edit] References
^ Archimedes, A Gold Thief and Buoyancy by Larry "Harris" Taylor, Ph.D.[1]
^ Density of water, as reported by Daniel Harris in Quantitative Chemical Analysis, 4th ed., p. 36, W. H. Freeman and Company, New York, 1995.

[edit] Books
Fundamentals of Aerodynamics Second Edition, McGraw-Hill, John D. Anderson, Jr.
Fundamentals of Fluid Mechanics Wiley, B.R. Munson, D.F. Young & T.H. Okishi
Introduction to Fluid Mechanics Fourth Edition, Wiley, SI Version, R.W. Fox & A.T. McDonald
Thermodynamics: An Engineering Approach Second Edition, McGraw-Hill, International Edition, Y.A. Cengel & M.A. Boles

[edit] See also
Charge density
Buoyancy
Bulk density
Dord
Energy density
Lighter than air
Number density
Population density
Specific weight
Standard temperature and pressure

[edit] External links
Glass Density Calculation Calculation of the density of glass at room temperature and of glass melts at 1000-1400°C
List of Elements of the Periodic Table - Sorted by Density
Retrieved from "http://en.wikipedia.org/wiki/Density"
Categories: Continuum mechanics | Introductory physics | Fundamental physics concepts | Physical quantity | Physical chemistry

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
Afrikaans
Alemannisch
العربية
Aragonés
Български
Català
Česky
Dansk
Deutsch
Eesti
Español
Esperanto
فارسی
Français
한국어
Ido
Bahasa Indonesia
Íslenska
Italiano
עברית
Latina
Latviešu
Lietuvių
Lojban
Македонски
മലയാളം
Bahasa Melayu
Nederlands
日本語
‪Norsk (bokmål)‬
‪Norsk (nynorsk)‬
O'zbek
Plattdüütsch
Polski
Português
Română
Русский
Shqip
Slovenščina
Српски / Srpski
Suomi
Svenska
ไทย
Tiếng Việt
Türkçe
Українська
中文

This page was last modified 09:00, 11 August 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Gas

2007-08-14T10:11:00.000-07:00

Gas
From Wikipedia, the free encyclopedia
Jump to: navigation, search
This article does not cite any references or sources.
Please help improve this article by adding citations to reliable sources. (help, get involved!)
Unverifiable material may be challenged and removed.
This article has been tagged since February 2007.
For other uses, see Gas (disambiguation).

Gas phase particles (atoms, molecules, or ions.)A gas is one of the four major states of matter, consisting of freely moving atoms or molecules without a definite shape. Compaired to the solid and liquid states of matter a gas has lower density and a lower viscosity. The volume of a gas will change with changes in temperature or pressure, as described by the ideal gas law. A gas also has the characteristic that it will diffuse readily, spreading appart in order to uniformly fill the space of any container.

Gas can also refer to gasoline, natural gas, and other various meanings.
Contents [hide]
1 Physics
2 Some types of gases
3 Etymology
4 See also
5 References

[edit] Physics
In a gas phase, the atoms or molecules constituting the matter basically move independently, (more freely than those in a solid or liquid) with no forces keeping them together or pushing them apart. Their only interactions are rare and random collisions. The particles move in random directions, at high speed. The range in speed is dependent on the temperature and defined by the Maxwell-Boltzmann distribution. Therefore, the gas phase is a completely disordered state. Following the second law of thermodynamics, when no work is being done on or by a gas, the gas particles will immediately diffuse to homogeneously fill any shape or volume of space that is made available to them.

The thermodynamic state of a gas is characterized by its volume, its temperature, and its pressure. These variables are related by the fundamental gas laws, which state that the pressure in an ideal gas is proportional to its temperature and number of molecules, but inversely proportional to its volume.

Like liquids and plasmas, gases are flowing and free moving fluids: they have the ability to flow and do not tend to return to their former configuration after deformation, although they do have viscosity. Unlike liquids, unconstrained gases in a vacuum environment do not occupy a fixed volume, but expand to fill the entire space. Note that this is true in the case of empty, vacuum environments. If one sprays carbon dioxide from a fire extinguisher, for example, the gas will not expand to fill the room. Instead, the gas will pour out like a fluid and pool on the floor. This is due to the fact that it is denser than the air surrounding it.[1]

The kinetic energy per molecule in a gas is the second greatest of the states of matter (after plasma). Because of this high kinetic energy, gas atoms and molecules tend to bounce off of any containing surface and off one another, the more powerfully as the kinetic energy is increased. A common misconception is that the collisions of the molecules with each other is essential to explain gas pressure, but in fact their random velocities are sufficient to define that quantity. Mutual collisions are important only for establishing the Maxwell-Boltzmann distribution.

Gas particles are normally well separated, as opposed to liquid particles, which are in contact. A material particle (say a dust mote) in a gas substrate moves in Brownian Motion. Since it is at the limit of (or beyond) current technology to observe individual gas particles (atoms or molecules), only theoretical calculations give suggestions as to how they move, but their motion is different from Brownian Motion. The reason is that Brownian Motion involves a smooth drag due to the frictional force of many gas molecules, punctuated by violent collisions of an individual (or several) gas molecule(s) with the particle. The particle (generally consisting of millions or billions of atoms) thus moves in a jagged course, yet not so jagged as we would expect to find if we could examine an individual gas molecule.

[edit] Some types of gases
Ideal gas, in physics, chemistry, and thermodynamics
Various hydrocarbon gases used for heating, lighting, and energy transmission:
Natural gas
Liquefied Petroleum Gas (LPG), including propane and butane
Syngas: various synthetic fuel gases: names include coal gas, water gas, illuminating gas, wood gas, producer gas, holzgas, air gas, blue gas, manufactured gas, town gas, hygas
Gas (chemical warfare), various poison gases used in warfare
Inhalational anaesthetic, including nitrous oxide or laughing gas.
Trace gas
Toxic gases
Noble gases

[edit] Etymology
The word "gas" was apparently proposed by the 17th century Flemish chemist Jan Baptist van Helmont, as a phonetic spelling of his Dutch pronunciation of the Greek word "chaos", which was used since 1538 after Paracelsus for "air". [2]

[edit] See also
Look up Gas in
Wiktionary, the free dictionary.Cooling curve
Fuel gas
Gas laws
Gas metal arc welding
Ideal gas, in physics
Kinetic theory of gases
Liquefied Petroleum Gas, including propane and butane
List of phases of matter
Natural gas
Vapor

[edit] References
^ Beatty, William J. "Recurring science misconceptions in K-6 textbooks". Retrieved 2007-06-08.
^ http://www.etymonline.com/index.php?term=gas
[hide]v • d • eStates of Matter (list)
Solid • Liquid • Gas • Plasma • Supercritical fluid • Superfluid • Supersolid • Degenerate matter • Quark-gluon plasma • Fermionic condensate • Bose–Einstein condensate • Strange matter • Melting point • Boiling point • Triple point • Critical point • Equation of state • Cooling curve

Retrieved from "http://en.wikipedia.org/wiki/Gas"
Categories: Articles lacking sources from February 2007 | All articles lacking sources | Fundamental physics concepts | Dutch loanwords | Gases | Phases of matter

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
العربية
Asturianu
Bosanski
Български
Català
Česky
Dansk
Deutsch
Eesti
Ελληνικά
Español
Esperanto
فارسی
Français
Gàidhlig
Galego
한국어
Hrvatski
Ido
Bahasa Indonesia
Italiano
עברית
ಕನ್ನಡ
Kurdî / كوردي
Latviešu
Lietuvių
Lojban
Magyar
Македонски
Bahasa Melayu
Nederlands
日本語
‪Norsk (bokmål)‬
‪Norsk (nynorsk)‬
Plattdüütsch
Polski
Português
Русский
Simple English
Slovenčina
Slovenščina
Српски / Srpski
Srpskohrvatski / Српскохрватски
Suomi
Svenska
தமிழ்
ไทย
Tiếng Việt
Türkçe
ייִדיש
Українська
中文

This page was last modified 07:19, 14 August 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Phase

2007-08-14T10:07:00.000-07:00

Phase (matter)
From Wikipedia, the free encyclopedia
Jump to: navigation, search
This article or section appears to have been copied and pasted from a source, possibly in violation of a copyright.
Please edit this article to remove any copyrighted text and to be an original source, following the Guide to layout and the Manual of Style. Remove this template after editing.

This article describes the form of a substance. For other uses, see Phase.
In the physical sciences, a phase is a set of states of a macroscopic physical system that have relatively uniform chemical composition and physical properties (i.e. density, crystal structure, index of refraction, and so forth).

Contents [hide]
1 Phases and states of matter
2 General definition of phases
3 Phase diagrams
4 Phase separation
5 Phase equilibrium
6 Phase transition
7 External links
8 See also

[edit] Phases and states of matter
Phases are sometimes confused with states of matter, but there are significant differences. States of matter refers to the differences between gases, liquids, solids, etc. If there are two regions in a chemical system that are in different states of matter, then they must be different phases. However, the reverse is not true -- a system can have multiple phases which are in equilibrium with each other and also in the same state of matter. For example, diamond and graphite are both solids but they are different phases, even though their composition may be identical. A system with oil and water at room temperature will be two different phases of differing composition, but both will be the liquid state of matter. This difference is especially important when considering the Gibbs' phase rule, which governs the number of allowed phases.

[edit] General definition of phases
In general, two different states of a system are in different phases if there is an abrupt change in their physical properties while transforming from one state to the other. Conversely, two states are in the same phase if they can be transformed into one another without any abrupt changes. There are, however, exceptions to this statement -- for example the liquid-gas critical point discussed below in the Phase Diagrams section.

An important point is that different types of phases are associated with different physical qualities. When discussing the solid, liquid, and gaseous phases, we talked about rigidity and compressibility, and the effects of varying the pressure and volume, because those are the relevant properties that distinguish a solid, a liquid, and a gas. On the other hand, when discussing paramagnetism and ferromagnetism, we look at the magnetization, because that is what distinguishes the ferromagnetic phase from the paramagnetic phase. Several more examples of phases will be given in the following section.

In more technical language, a phase is a region in the parameter space of thermodynamic variables in which the free energy is analytic; between such regions there are abrupt changes in the properties of the system, which correspond to discontinuities in the derivatives of the free energy function. As long as the free energy is analytic, all thermodynamic properties (such as entropy, heat capacity, magnetization, and compressibility) will be well-behaved, because they can be expressed in terms of the free energy and its derivatives. For example, the entropy is the first derivative of the free energy with temperature.

When a system goes from one phase to another, there will generally be a stage where the free energy is non-analytic. This is a phase transition. Due to this non-analyticity, the free energies on either side of the transition are two different functions, so one or more thermodynamic properties will behave very differently after the transition. The property most commonly examined in this context is the heat capacity. During a transition, the heat capacity may become infinite, jump abruptly to a different value, or exhibit a "kink" or discontinuity in its derivative. See also differential scanning calorimetry.

Possible graphs of heat capacity (C) against temperature (T) at a phase transition

[edit] Phase diagrams
Main article: Phase diagram
The different phases of a system may be represented using a phase diagram. The axes of the diagrams are the relevant thermodynamic variables. For simple mechanical systems, we generally use the pressure and temperature.

A phase diagram for a typical material exhibiting solid, liquid and gaseous phasesThe markings on the phase diagram show the points where the free energy is non-analytic. The open spaces, where the free energy is analytic, correspond to the phases. The phases are separated by lines of non-analyticity, where phase transitions occur, which are called phase boundaries.

In the diagram, the phase boundary between liquid and gas does not continue indefinitely. Instead, it terminates at a point on the phase diagram called the critical point. At temperatures and pressure above the critical point, the physical property differences that differentiate the liquid phase from the gas phase become less defined. This reflects the fact that, at extremely high temperatures and pressures, the liquid and gaseous phases become indistinguishable. In water, the critical point occurs at around 647 K (374 °C or 705 °F) and 22.064 MPa.

The existence of the liquid-gas critical point reveals a slight ambiguity in our above definitions. When going from the liquid to the gaseous phase, one usually crosses the phase boundary, but it is possible to choose a path that never crosses the boundary by going to the right of the critical point. Thus, phases can sometimes blend continuously into each other. This new phase which has some properties that are similar to a liquid and some properties that are similar to a gas is called a supercritical fluid. We should note, however, that this does not always happen. For example, it is impossible for the solid-liquid phase boundary to end in a critical point in the same way as the liquid-gas boundary, because the solid and liquid phases have different symmetry.

An interesting thing to note is that the solid-liquid phase boundary in the phase diagram of most substances, such as the one shown above, has a positive slope. This is due to the solid phase having a higher density than the liquid, so that increasing the pressure increases the melting temperature. However, in the phase diagram for water the solid-liquid phase boundary has a negative slope. This reflects the fact that ice has a lower density than water, which is an unusual property for a material.

[edit] Phase separation
Phase separation is transformation of a homogenous system in two (or more) phases and commonly encountered in many branches of science and technology. One example is the crystallization of a solid from a solution. A universal mathematical model of phase separation is provided by the Cahn-Hilliard Equation.

[edit] Phase equilibrium
The distribution of kinetic energy among molecules is not uniform, and it changes randomly. This means that at, say, the surface of a liquid, there may be an individual molecule with enough kinetic energy to jump into the gas phase. Likewise, individual gas molecules may have low enough kinetic energy to join other molecules in the liquid phase. This phenomenon means that at any given temperature and pressure, multiple phases may co-exist.

For example, under standard conditions for temperature and pressure, a bowl of liquid water in dry air will evaporate until the partial pressure of gaseous water equals the vapor pressure of water. At this point, the rate of molecules leaving and entering the liquid phase becomes the same (due to the increased number of gaseous water molecules available to re-condense). The fact that liquid molecules with above-average kinetic energy have been removed from the bowl results in evaporative cooling. Similar processes may occur on other types of phase boundaries.

Gibbs' phase rule relates the number of possible phases, variables such as temperature and pressure, and whether or not an equilibrium will be reached.

[edit] Phase transition
A phase transition or, phase change, describes when a substance changes its state of matter - ex. ice melting to water is a phase change because a solid changed to a liquid. For a phase change to occur, energy must be added or removed from the substance. Normally adding or removing energy will change the temperature of the substance as the kinetic energy of the particles will increase or decrease. During a phase change however, the potential energy of the substance changes as the particles are moved further apart or closer together. There is no change in kinetic energy of the particles and therefore no resulting change in temperature.

[edit] External links
French physicists find a solution that reversibly solidifies with a rise in temperature - α-cyclodextrin, water, and 4-methylpyridine

[edit] See also
State of matter
Condensed matter physics
Cooling curve
Supercooling
Superheating
Multiphasic liquid
Retrieved from "http://en.wikipedia.org/wiki/Phase_%28matter%29"
Categories: Copied and pasted articles and sections | Fundamental physics concepts | Condensed matter physics | Chemical engineering

Electron shell

2007-08-14T10:05:00.000-07:00

Electron shell
From Wikipedia, the free encyclopedia
Jump to: navigation, search

Example of a sodium electron shell modelAn electron shell, also known as a main energy level, is a group of atomic orbitals with the same value of the principal quantum number n. Electron shells are made up of one or more electron subshells, or sublevels, which have two or more orbitals with the same angular momentum quantum number l. Electron shells make up the electron configuration of an atom. It can be shown that the number of electrons that can reside in a shell is equal to 2n2.

Contents [hide]
1 History
2 Valence shell
3 Subshells
4 See also
5 References

[edit] History
The existence of electron shells was first observed experimentally in Charles Barkla's and Henry Moseley's X-ray absorption studies. Barkla labelled them with the letters K, L, M, etc. (The origin of this terminology was alphabetic. K and L were originally called B and A, but were later renamed to leave room for hypothetical spectral lines that were never discovered.) These letters were later found to correspond to the n-values 1, 2, 3, etc. They are used in the spectroscopic Siegbahn notation.

The name for electron shells originates from the Bohr model, in which groups of electrons were believed to orbit the nucleus at certain distances, so that their orbits formed "shells".

[edit] Valence shell
The valence shell is the outermost shell of an atom in its uncombined state, which contains the electrons most likely to account for the nature of any reactions involving the atom and of the bonding interactions it has with other atoms. Care must be taken to note that the outermost shell of an ion is not commonly termed valence shell. Electrons in the valence shell are referred to as valence electrons. The physical chemist Gilbert Lewis was responsible for much of the early development of the theory of the participation of valence shell electrons in chemical bonding. Linus Pauling later generalized and extended the theory while applying insights from quantum mechanics.

In a noble gas, an atom tends to have 8 electrons in its outer shell (except helium, which is only able to fill its shell with 2 electrons). This serves as the model for the octet rule which is mostly applicable to main group elements of the second and third periods. In terms of atomic orbitals, the electrons in the valence shell are distributed 2 in the single s orbital and 2 each in the three p orbitals.

For coordination complexes containing transition metals, the valence shell consists of electrons in these s and p orbitals, as well as up to 10 additional electrons, distributed as 2 into each of 5 d orbitals, to make a total of 18 electrons in a complete valence shell for such a compound. This is referred to as the eighteen electron rule.

Possible Number of Electrons in shells 1-7 Shell Electrons
K 2
L 8
M 18
N 32
O 32
P 18
Q 8

[edit] Subshells
Electron subshells are identified by the letters s, p, d, f, g, h, i, etc., corresponding to the azimuthal quantum numbers (l-values) 0, 1, 2, 3, 4, 5, 6, etc. Each shell can hold up to 2, 6, 10, 14, 18, 22 and 26 electrons respectively, or 2(2l + 1) electrons in each subshell. The notation 's', 'p', 'd', and 'f' originate from a now-discredited system of categorizing spectral lines as "sharp", "principal", "diffuse", or "fundamental", based on their observed fine structure. When the first four types of orbitals were described, they were associated with these spectral line types, but there were no other names. The designations 'g', 'h', and so on, were derived by following alphabetical order.

[edit] See also
Atomic orbital
Electron configuration
Molecular orbital

[edit] References
Tipler, Paul & Ralph Llewellyn (2003). Modern Physics (4th ed.). New York: W. H. Freeman and Company. ISBN 0-7167-4345-0
Retrieved from "http://en.wikipedia.org/wiki/Electron_shell"
Categories: Atomic physics | Quantum chemistry | Chemical bonding

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
العربية
العربية
Deutsch
Deutsch
Français
עברית
Magyar
日本語
Polski
Português
Русский
Suomi
Українська
中文

This page was last modified 15:47, 24 July 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Electron

2007-08-14T10:03:00.000-07:00

Electron
From Wikipedia, the free encyclopedia
Jump to: navigation, search
For other uses, see Electron (disambiguation).
"e-" redirects here. For the Internet-related prefix e-, see Wiktionary's entry e-.
Electron

Theoretical estimates of the electron density for the first few hydrogen atom electron orbitals shown as cross-sections with color-coded probability density
Composition: Elementary particle
Family: Fermion
Group: Lepton
Generation: First
Interaction: Gravity, Electromagnetic, Weak
Antiparticle: Positron
Theorized: G. Johnstone Stoney (1874)
Discovered: J.J. Thomson (1897)
Symbol: e−, β−
Mass: 9.109 3826(16) × 10–31 kg[1]

5.485 899 0945(24) × 10–4 u

1⁄1822.888 4849(8) u

0.510 998 918(44) MeV/c2
Electric charge: –1.602 176 487(40) × 10–19 C[2]
Spin: ½
The electron is a fundamental subatomic particle that carries a negative electric charge. It is a spin-½ lepton that participates in electromagnetic interactions, and its mass is less than one thousandth of that of the smallest atom. Its electric charge is defined by convention to be negative, with a value of −1 in atomic units. Together with atomic nuclei, electrons make up atoms; their interaction with adjacent nuclei is the main cause of chemical bonding.

Contents [hide]
1 History
2 Classification
3 Properties and behavior
4 In practice
4.1 In the universe
4.2 In industry
4.3 In the laboratory
4.4 In medicine
5 In theory
6 Notes
7 See also
8 External links

[edit] History
The name "electron" comes from the Greek word for amber, ήλεκτρον. This material played an essential role in the discovery of electrical phenomena. The ancient Greeks knew, for example, that rubbing a piece of amber with fur left an electric charge on its surface, which could then create sparks. For more about the history of the term electricity, see History of electricity.

The electron as a unit of charge in electrochemistry was posited by G. Johnstone Stoney in 1874, who also coined the term electron in 1894. During the late 1890s a number of physicists posited that electricity could be conceived of as being made of discrete units, which were given a variety of names, but their reality had not been confirmed in a compelling way.

The discovery that the electron was a subatomic particle was made in 1897 by J.J. Thomson at the Cavendish Laboratory at Cambridge University, while he was studying cathode ray tubes. A cathode ray tube is a sealed glass cylinder in which two electrodes are separated by a vacuum. When a voltage is applied across the electrodes, cathode rays are generated, causing the tube to glow. Through experimentation, Thomson discovered that the negative charge could not be separated from the rays (by the application of magnetism), and that the rays could be deflected by an electric field. He concluded that these rays, rather than being waves, were composed of negatively charged particles he called "corpuscles". He measured their mass-to-charge ratio and found it to be over a thousand times smaller than that of a hydrogen ion, suggesting that they were either very highly charged or very small in mass. Later experiments by other scientists upheld the latter conclusion. Their mass-to-charge ratio was also independent of the choice of cathode material and the gas originally on vacuum tube. This led Thomson to conclude that they were universal among all materials.

The electron's charge was carefully measured by Robert Millikan in his oil-drop experiment of 1909.

The periodic law states that the chemical properties of elements largely repeat themselves periodically and is the foundation of the periodic table of elements. The law itself was initially explained by the atomic mass of the element. However, as there were anomalies in the periodic table, efforts were made to find a better explanation for it. In 1913, Henry Moseley introduced the concept of the atomic number and explained the periodic law in terms of the number of protons each element has. In the same year, Niels Bohr showed that electrons are the actual foundation of the table. In 1916, Gilbert Newton Lewis explained the chemical bonding of elements by electronic interactions. hgg

[edit] Classification
The electron is in the class of subatomic particles called leptons, which are believed to be fundamental particles.

As with all particles, electrons can also act as waves. This is called the wave-particle duality, also known by the term complementarity coined by Niels Bohr and can be demonstrated using the double-slit experiment.

The antiparticle of an electron is the positron, which has the same mass but positive rather than negative charge. The discoverer of the positron, Carl D. Anderson, proposed calling standard electrons negatrons, and using electron as a generic term to describe both the positively and negatively charged variants. This usage never caught on and is rarely, if ever, encountered today.

[edit] Properties and behavior
Electrons have an electric charge of −1.6022 × 10−19 coulomb, a mass of 9.11 × 10−31 kg based on charge/mass measurements and a relativistic rest mass of about 0.511 MeV/c2. The mass of the electron is approximately 1/1836 of the mass of the proton. The common electron symbol is e−.[1]

According to quantum mechanics, electrons can be represented by wavefunctions, from which a calculated probabilistic electron density can be determined. The orbital of each electron in an atom can be described by a wavefunction. Based on the Heisenberg uncertainty principle, the exact momentum and position of the actual electron cannot be simultaneously determined. This is a limitation which, in this instance, simply states that the more accurately we know a particle's position, the less accurately we can know its momentum, and vice versa.

The electron has spin ½ and is a fermion (it follows Fermi-Dirac statistics). In addition to its intrinsic angular momentum, an electron has an intrinsic magnetic moment along its spin axis.

Electrons in an atom are bound to that atom; electrons moving freely in vacuum, space or certain media are free electrons that can be focused into an electron beam. When free electrons move, there is a net flow of charge, this flow is called an electric current. The drift velocity of electrons in metal wires is on the order of mm/hour. However, the speed at which a current at one point in a wire causes a current in other parts of the wire is typically 75% of light speed.

In some superconductors, pairs of electrons move as Cooper pairs in which their motion is coupled to nearby matter via lattice vibrations called phonons. The distance of separation between Cooper pairs is roughly 100 nm. (Rohlf, J.W.)

A body has an electric charge when that body has more or fewer electrons than are required to balance the positive charge of the nuclei. When there is an excess of electrons, the object is said to be negatively charged. When there are fewer electrons than protons, the object is said to be positively charged. When the number of electrons and the number of protons are equal, their charges cancel each other and the object is said to be electrically neutral. A macroscopic body can develop an electric charge through rubbing, by the phenomenon of triboelectricity.

When electrons and positrons collide, they annihilate each other and produce pairs of high energy photons or other particles. On the other hand, high-energy photons may transform into an electron and a positron by a process called pair production, but only in the presence of a nearby charged particle, such as a nucleus.

The electron is currently described as a fundamental particle or an elementary particle. It has no substructure. Hence, for convenience, it is usually defined or assumed to be a point-like mathematical point charge, with no spatial extension. However, when a test particle is forced to approach an electron, we measure changes in its properties (charge and mass). This effect is common to all elementary particles: Current theory suggests that this effect is due to the influence of vacuum fluctuations in its local space, so that the properties measured from a significant distance are considered to be the sum of the bare properties and the vacuum effects (see renormalization).

The classical electron radius is 2.8179 × 10−15 m. This is the radius that is inferred from the electron's electric charge, by using the classical theory of electrodynamics alone, ignoring quantum mechanics. Classical electrodynamics (Maxwell's electrodynamics) is the older concept that is widely used for practical applications of electricity, electrical engineering, semiconductor physics, and electromagnetics; quantum electrodynamics, on the other hand, is useful for applications involving modern particle physics and some aspects of optical, laser and quantum physics.

Based on current theory, the speed of an electron can approach, but never reach, c (the speed of light in a vacuum). This limitation is attributed to Einstein's theory of special relativity which defines the speed of light as a constant within all inertial frames. However, when relativistic electrons are injected into a dielectric medium, such as water, where the local speed of light is significantly less than c, the electrons will (temporarily) be traveling faster than light in the medium. As they interact with the medium, they generate a faint bluish light, called Cherenkov radiation.

The effects of special relativity are based on a quantity known as γ or the Lorentz factor. γ is a function of v, the velocity of the particle. It is defined as:

The energy necessary to accelerate a particle is γ minus one times the rest mass. For example, the Stanford linear accelerator can accelerate an electron to roughly 51 GeV [1]. This gives a gamma of 100,000, since the rest mass of an electron is 0.51 MeV/c² (the relativistic mass of this electron is 100,000 times its rest mass). Solving the equation above for the speed of the electron (and using an approximation for large γ) gives:

[edit] In practice

[edit] In the universe
Scientists believe that the number of electrons existing in the known universe is at least 1079. This number amounts to an average density of about one electron per cubic metre of space. Astronomers have estimated that 90% of the mass of atoms in the universe is hydrogen, which is made of one electron and one proton.

Based on the classical electron radius and assuming a dense sphere packing, it can be calculated that the number of electrons that would fit in the observable universe is on the order of 10130.

[edit] In industry
Electron beams are used in welding, lithography, scanning electron microscopes and transmission electron microscopes. LEED and RHEED are also important tools where electrons are used.

They are also at the heart of cathode ray tubes, which are used extensively as display devices in laboratory instruments, computer monitors and television sets. In photomultiplier tubes, one photon strikes the photocathode, initiating an avalanche of electrons that produces a detectable current.

[edit] In the laboratory
Electron microscopes are used to magnify details up to 500,000 times. Quantum effects of electrons are used in Scanning tunneling microscope to study features at the atomic scale.

[edit] In medicine
In radiation therapy, electron beams are used for treatment of superficial tumours.

[edit] In theory
In relativistic quantum mechanics, the electron is described by the Dirac Equation which defines the electron as a (mathematical) point. In quantum field theory, the behavior of the electron is described by quantum electrodynamics (QED), a U(1) gauge theory. In Dirac's model, an electron is defined to be a mathematical point, a point-like, charged "bare" particle surrounded by a sea of interacting pairs of virtual particles and antiparticles. These provide a correction of just over 0.1% to the predicted value of the electron's gyromagnetic ratio from exactly 2 (as predicted by Dirac's single-particle model). The extraordinarily precise agreement of this prediction with the experimentally determined value is viewed as one of the great achievements of modern physics.[3]

In the Standard Model of particle physics, the electron is the first-generation charged lepton. It forms a weak isospin doublet with the electron neutrino; these two particles interact with each other through both the charged and neutral current weak interaction. The electron is very similar to the two more massive particles of higher generations, the muon and the tau lepton, which are identical in charge, spin, and interaction but differ in mass.

The antimatter counterpart of the electron is the positron. The positron has the same amount of electrical charge as the electron, except that the charge is positive. It has the same mass and spin as the electron. When an electron and a positron meet, they may annihilate each other, giving rise to two gamma-ray photons emitted at roughly 180° to each other. If the electron and positron had negligible momentum, each gamma ray will have an energy of 0.511 MeV. See also Electron-positron annihilation.

Electrons are a key element in electromagnetism, a theory that is accurate for macroscopic systems, and for classical modelling of microscopic systems.

[show]v • d • eQuantum Electrodynamics
electron • positron • photon • self-energy • vacuum polarization • vertex function • Gupta-Bleuler formalism • ξ gauge • Ward-Takahashi identity • Compton scattering • Bhabha scattering • Møller scattering • anomalous magnetic dipole moment • bremsstrahlung • positronium

[edit] Notes
^ a b All masses are CODATA values accessed via the NIST’s electron mass page. The fractional version’s denominator is the inverse of the decimal value (along with its relative standard uncertainty of 4.4 × 10–10)
^ The electron’s charge is the negative of elementary charge (which is a positive value for the proton). CODATA value accessed via the NIST’s elementary charge page.
^ *Griffiths, David J. (2004). Introduction to Quantum Mechanics (2nd ed.). Prentice Hall. ISBN 0-13-805326-X.

Electron configuration

2007-08-14T10:01:00.000-07:00

Electron configuration
From Wikipedia, the free encyclopedia
Jump to: navigation, search

Electron atomic and molecular orbitalsIn atomic physics and quantum chemistry, the electron configuration is the arrangement of electrons in an atom, molecule, or other physical structure (eg, a crystal).

Like other elementary particles, the electron is subject to the laws of quantum mechanics, and exhibits both particle-like and wave-like properties. Formally, the quantum state of a particular electron is defined by its wavefunction, a complex-valued function of space and time. According to the Copenhagen interpretation of quantum mechanics, the position of a particular electron is not well defined until an act of measurement causes it to be detected. The probability that the act of measurement will detect the electron at a particular point in space is proportional to the square of the absolute value of the wavefunction at that point.

Electrons are able to move from one energy level to another by emission or absorption of a quantum of energy, in the form of a photon. Because of the Pauli exclusion principle, no more than two electrons may exist in a given atomic orbital; therefore an electron may only leap to another orbital if there is a vacancy there.

Knowledge of the electron configuration of different atoms is useful in understanding the structure of the periodic table of elements. The concept is also useful for describing the chemical bonds that hold molecules together. In bulk materials this same idea helps explain the peculiar properties of lasers and semiconductors.

Contents [hide]
1 Electron configuration in atoms
1.1 Summary of the quantum numbers
1.2 Shells and subshells
1.2.1 Worked example
1.3 Notation
1.4 Aufbau principle
1.4.1 Orbitals table
1.4.2 Exceptions in 3d, 4d, 5d
1.5 Relation to the structure of the periodic table
2 Electron configuration in molecules
3 Electron configuration in solids
4 See also
5 Notes

[edit] Electron configuration in atoms
The discussion below presumes knowledge of material contained at Atomic orbital.

[edit] Summary of the quantum numbers
The state of an electron in an atom is given by four quantum numbers. Three of these are integers and are properties of the atomic orbital in which it sits (a more thorough explanation is given in that article).

number denoted allowed range represents
principal quantum number n integer, 1 or more Partly the overall energy of the orbital, and by extension its general distance from the nucleus. In short, the energy level it is in. (1+)
azimuthal quantum number l integer, 0 to n-1 The orbital's angular momentum, also seen as the number of nodes in the density plot. Otherwise known as its orbital. (s=0, p=1...)
magnetic quantum number m integer, -l to +l, including zero. Determines energy shift of an atomic orbital due to external magnetic field (Zeeman effect). Indicates spatial orientation.
spin quantum number ms +½ or -½ (sometimes called "up" and "down") Spin is an intrinsic property of the electron and independent of the other numbers. s and l in part determine the electron's magnetic dipole moment.

No two electrons in one atom can have the same set of these four quantum numbers (Pauli exclusion principle).

[edit] Shells and subshells
Shells and subshells (also called energy levels and sublevels) are defined by the quantum numbers, not by the distance of its electrons from the nucleus, or even their overall energy. In large atoms, shells above the second shell overlap (see Aufbau principle).

States with the same value of n are related, and said to lie within the same electron shell.
States with the same value of n and also l are said to lie within the same electron subshell, and those electrons having the same n and l are called equivalent electrons.
If the states also share the same value of m, they are said to lie in the same atomic orbital.
Because electrons have only two possible spin states, an atomic orbital cannot contain more than two electrons (Pauli exclusion principle).

A subshell can contain up to 4l+2 electrons; a shell can contain up to 2n² electrons; where n equals the shell number.

[edit] Worked example
Here is the electron configuration for a filled fifth shell:

Shell Subshell Orbitals Electrons
n = 5 l = 0 m = 0 → 1 type s orbital → max 2 electrons
l = 1 m = -1, 0, +1 → 3 type p orbitals → max 6 electrons
l = 2 m = -2, -1, 0, +1, +2 → 5 type d orbitals → max 10 electrons
l = 3 m = -3, -2, -1, 0, +1, +2, +3 → 7 type f orbitals → max 14 electrons
l = 4 m = -4, -3 -2, -1, 0, +1, +2, +3, +4 → 9 type g orbitals → max 18 electrons
Total: max 50 electrons

This information can be written as 5s2 5p6 5d10 5f14 5g18 (see below for more details on notation).

[edit] Notation
Physicists and chemists use a standard notation to describe atomic electron configurations. In this notation, a subshell is written in the form nxy, where n is the shell number, x is the subshell label and y is the number of electrons in the subshell. An atom's subshells are written in order of increasing energy - in other words, the sequence in which they are filled (see Aufbau principle below).

For instance, ground-state hydrogen has one electron in the s orbital of the first shell, so its configuration is written 1s1. Lithium has two electrons in the 1s subshell and one in the (higher-energy) 2s subshell, so its ground-state configuration is written 1s2 2s1. Phosphorus (atomic number 15), is as follows: 1s2 2s2 2p6 3s2 3p3.

For atoms with many electrons, this notation can become lengthy. It is often abbreviated by noting that the first few subshells are identical to those of one or another noble gas. Phosphorus, for instance, differs from neon (1s2 2s2 2p6) only by the presence of a third shell. Thus, the electron configuration of neon is pulled out, and phosphorus is written as follows: [Ne]3s2 3p3.

An even simpler version is simply to quote the number of electrons in each shell, e.g. (again for phosphorus): 2-8-5.

The orbital labels s, p, d, and f originate from a now-discredited system of categorizing spectral lines as sharp, principal, diffuse, and fundamental, based on their observed fine structure. When the first four types of orbitals were described, they were associated with these spectral line types, but there were no other names. The designation g was derived by following alphabetical order. Shells with more than five subshells are theoretically permissible, but this covers all discovered elements. For mnemonic reasons, some call the s and p orbitals spherical and peripheral.

[edit] Aufbau principle
In the ground state of an atom (the condition in which it is ordinarily found), the electron configuration generally follows the Aufbau principle. According to this principle, electrons enter into states in order of the states' increasing energy; i.e., the first electron goes into the lowest-energy state, the second into the next lowest, and so on. The order in which the states are filled is as follows:

s p d f g
1 1
2 2 3
3 4 5 7
4 6 8 10 13
5 9 11 14 17 21
6 12 15 18 22
7 16 19 23
8 20 24

The order of increasing energy of the subshells can be constructed by going through downward-leftward diagonals of the table above (also see the diagram at the top of the page), going from the topmost diagonals to the bottom. The first (topmost) diagonal goes through 1s; the second diagonal goes through 2s; the third goes through 2p and 3s; the fourth goes through 3p and 4s; the fifth goes through 3d, 4p, and 5s; and so on. In general, a subshell that is not "s" is always followed by a "lower" subshell of the next shell; e.g. 2p is followed by 3s; 3d is followed by 4p, which is followed by 5s, 4f is followed by 5d, which is followed by 6p, and then 7s. This explains the ordering of the blocks in the periodic table.

A pair of electrons with identical spins has slightly less energy than a pair of electrons with opposite spins. Since two electrons in the same orbital must have opposite spins, this causes electrons to prefer to occupy different orbitals. This preference manifests itself if a subshell with l > 0 (one that contains more than one orbital) is less than full. For instance, if a p subshell contains four electrons, two electrons will be forced to occupy one orbital, but the other two electrons will occupy both of the other orbitals, and their spins will be equal. This phenomenon is called Hund's rule.

The Aufbau principle can be applied, in a modified form, to the protons and neutrons in the atomic nucleus (see the shell model of nuclear physics).

[edit] Orbitals table
This table shows all orbital configurations up to 7s, therefore it covers the simple electronic configuration for all elements from the periodic table up to Ununbium (element 112) with the exception of Lawrencium (element 103), which would require a 7p orbital.

s (l=0) p (l=1) d (l=2) f (l=3)
n=1
n=2
n=3
n=4
n=5
n=6
n=7

[edit] Exceptions in 3d, 4d, 5d
A d subshell that is half-filled or full (ie 5 or 10 electrons) is more stable than the s subshell of the next shell. This is the case because it takes less energy to maintain an electron in a half-filled d subshell than a filled s subshell. For instance, copper (atomic number 29) has a configuration of [Ar]4s1 3d10, not [Ar]4s2 3d9 as one would expect by the Aufbau principle. Likewise, chromium (atomic number 24) has a configuration of [Ar]4s1 3d5, not [Ar]4s2 3d4 where [Ar] represents the configuration for argon.

Exceptions in Period 4:[1]

Element Z Electron configuration Short electron conf.
Scandium 21 1s2 2s2 2p6 3s2 3p6 4s2 3d1 [Ar] 4s2 3d1
Titanium 22 1s2 2s2 2p6 3s2 3p6 4s2 3d2 [Ar] 4s2 3d2
Vanadium 23 1s2 2s2 2p6 3s2 3p6 4s2 3d3 [Ar] 4s2 3d3
Chromium 24 1s2 2s2 2p6 3s2 3p6 4s1 3d5 [Ar] 4s1 3d5
Manganese 25 1s2 2s2 2p6 3s2 3p6 4s2 3d5 [Ar] 4s2 3d5
Iron 26 1s2 2s2 2p6 3s2 3p6 4s2 3d6 [Ar] 4s2 3d6
Cobalt 27 1s2 2s2 2p6 3s2 3p6 4s2 3d7 [Ar] 4s2 3d7
Nickel 28 1s2 2s2 2p6 3s2 3p6 4s2 3d8 [Ar] 4s2 3d8
Copper 29 1s2 2s2 2p6 3s2 3p6 4s1 3d10 [Ar] 4s1 3d10
Zinc 30 1s2 2s2 2p6 3s2 3p6 4s2 3d10 [Ar] 4s2 3d10
Gallium 31 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p1 [Ar] 3d10 4s2 4p1

Exceptions in Period 5:[2]

Element Z Electron configuration Short electron conf.
Yttrium 39 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d1 [Kr] 5s2 4d1
Zirconium 40 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d2 [Kr] 5s2 4d2
Niobium 41 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d4 [Kr] 5s1 4d4
Molybdenum 42 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d5 [Kr] 5s1 4d5
Technetium 43 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d5 [Kr] 5s2 4d5
Ruthenium 44 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d7 [Kr] 5s1 4d7
Rhodium 45 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d8 [Kr] 5s1 4d8
Palladium 46 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 4d10 [Kr] 4d10
Silver 47 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 4d10 [Kr] 5s1 4d10
Cadmium 48 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 [Kr] 5s2 4d10
Indium 49 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p1 [Kr] 5s2 4d10 5p1

Exceptions in Period 6:[3]

Element Z Short electron conf.
Iridium 77 [Xe] 6s2 4f14 5d7
Platinum 78 [Xe] 6s1 4f14 5d9
Gold 79 [Xe] 6s1 4f14 5d10
Mercury 80 [Xe] 6s2 4f14 5d10
Thallium 81 [Xe] 6s2 4f14 5d10 6p1

[edit] Relation to the structure of the periodic table
Main article: Periodic_table#Structure_of_the_periodic_table
Electron configuration is intimately related to the structure of the periodic table. The chemical properties of an atom are largely determined by the arrangement of the electrons in its outermost "valence" shell (although other factors, such as atomic radius, atomic mass, and increased accessibility of additional electronic states also contribute to the chemistry of the elements as atomic size increases) therefore elements in the same table group are chemically similar because they contain the same number of "valence" electrons.

[edit] Electron configuration in molecules
In molecules, the situation becomes more complex, as each molecule has a different orbital structure. See the molecular orbital article and the linear combination of atomic orbitals method for an introduction and the computational chemistry article for more advanced discussions.

[edit] Electron configuration in solids
In a solid, the electron states become very numerous. They cease to be discrete, and effectively blend together into continuous ranges of possible states (an electron band). The notion of electron configuration ceases to be relevant, and yields to band theory.

[edit] See also
Atomic electron configuration table
Periodic table (electron configurations)
Atomic orbital
Energy level
Molecular term symbol
HOMO/LUMO
Atompaw Software package for electron configuration calculations
Pwpaw Software package for electron configuration calculations
Periodic Table Group

[edit] Notes
^ This can be most easily understood with the electron configuration diagram of Scandium at Webelements.
^ This can be most easily understood with the electron configuration diagram of Yttrium at Webelements.
^ This can be most easily understood with the electron configuration diagram of Iridium at Webelements.
Retrieved from "http://en.wikipedia.org/wiki/Electron_configuration"
Categories: Chemical properties | Atomic physics | Molecular physics | Quantum chemistry | Theoretical chemistry

Atomic mass

2007-08-14T09:59:00.000-07:00

Atomic mass
From Wikipedia, the free encyclopedia
(Redirected from Standard atomic weight)
Jump to: navigation, search
The atomic mass (ma) is the mass of an atom at rest, most often expressed in unified atomic mass units.[1] The atomic mass is sometimes incorrectly used as a synonym of relative atomic mass, average atomic mass and atomic weight; however, these differ subtly from the atomic mass. The atomic mass is defined as the mass of an atom, which can only be one isotope at a time and is not an abundance-weighted average. The actual numerical difference is usually very small such that it does not affect most bulk calculations but such an error can be critical when considering individual atoms.

The relative atomic mass (Ar) (also known as atomic weight and average atomic mass) is the average of the atomic masses of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance.[2] This is frequently used as a synonym for the standard atomic weight and is not incorrect to do so since the standard atomic weights are relative atomic masses, although it is less specific to do so. Relative atomic mass also refers to non-terrestrial environments and highly specific terrestrial environments that deviate from the average or have different certainties (number of significant figures) than the standard atomic weights.

The standard atomic weight refers to the mean relative atomic mass of an element in the local environment of the Earth's crust and atmosphere as determined by the IUPAC Commission on Atomic Weights and Isotopic Abundances.[3] These are what are included in a standard periodic table and is what is used in most bulk calculations. An uncertainty in brackets is included which often reflects natural variability in isotopic distribution rather than uncertainty in measurement.[4] For synthetic elements the isotope formed depends on the means of synthesis, so the concept of natural isotope abundance has no meaning. Therefore, for synthetic elements the total nucleon count of the most stable isotope (ie, the isotope with the longest half-life) is listed in brackets in place of the standard atomic weight. Lithium represents a unique case where the natural abundances of the isotopes have been perturbed by human activities to the point of affecting the uncertainty in its standard atomic weight, even in samples obtained from natural sources such as rivers.

The relative isotopic mass is the relative mass of the isotope, scaled with carbon-12 as exactly 12. No other isotopes have whole number masses due to the different mass of neutrons and protons, as well as loss/gain of mass to binding energy. However, since mass defect due to binding energy is minimal compared to the mass of a nucleon, rounding the atomic mass of an isotope tells you the total nucleon count. Neutron count can then be derived by subtracting the atomic number.

Contents [hide]
1 Mass Defects in atomic masses
2 Measurement of atomic masses
3 Conversion factor between atomic mass units and grams
4 The relationship between atomic and molecular masses
5 History
6 Table of standard atomic weights
7 See also
8 External links
9 References

[edit] Mass Defects in atomic masses
The pattern in the amounts the atomic masses deviate from their mass numbers is as follows: the deviation starts positive at hydrogen-1, becomes negative until a minimum is reached at iron-56, iron-58 and nickel-62, then increases to positive values in the heavy isotopes, with increasing atomic number. This corresponds to the following: nuclear fission in an element heavier than iron produces energy, and fission in any element lighter than iron requires energy. The opposite is true of nuclear fusion reactions: fusion in elements lighter than iron produces energy, and fusion in elements heavier than iron requires energy.

[edit] Measurement of atomic masses
Direct comparison and measurement of the masses of atoms is achieved with mass spectrometry.

[edit] Conversion factor between atomic mass units and grams
The standard scientific unit for dealing with atoms in macroscopic quantities is the mole (mol), which is defined arbitrarily as the amount of a substance with as many atoms or other units as there are in 12 grams of the carbon isotope C-12. The number of atoms in a mole is called Avogadro's number, the value of which is approximately 6.02 × 1023. One mole of a substance always contains almost exactly the relative atomic or molar mass of that substance (which is the concept of molar mass), expressed in grams. For example, the relative atomic mass of iron is 55.847, and therefore one mole of iron has a mass of 55.847 grams. The formulaic conversion between atomic mass and SI mass in grams for a single atom is

where u is the atomic mass unit and NA is Avogadro's number.

[edit] The relationship between atomic and molecular masses
Similar definitions apply to molecules. One can compute the molecular mass of a compound by adding the atomic masses of its constituent atoms (nuclides). One can compute the molar mass of a compound by adding the relative atomic masses of the elements given in the chemical formula. In both cases the multiplicity of the atoms (the number of times it occurs) must be taken into account, usually by multiplication of each unique mass by its multiplicity.

[edit] History
Before the 1960s, this was expressed so that the oxygen-16 isotope received the atomic weight 16, however, the proportions of oxygen-17 and oxygen-18 present in natural oxygen, which were also used to calculate atomic mass led to two different tables of atomic mass.

Formerly chemists and physicists used two different atomic mass scales. The chemists used a scale such that the natural mixture of oxygen isotopes had an atomic mass 16, while the physicists assigned the same number 16 to the atomic mass of the most common oxygen isotope (containing eight protons and eight neutrons). The unified scale based on carbon-12, 12C, met the physicists' need to base the scale on a pure isotope, while being numerically close to the old chemists' scale.

The term atomic weight is being phased out slowly and being replaced by relative atomic mass, in most current usage. The history of this shift in nomenclature reaches back to the 1960's and has been the source of much debate in the scientific community. The debate was largely created by the adoption of the unified atomic mass unit and the realization that weight was in some ways an inappropriate term. The argument for keeping the term "atomic weight" was primarily that it was a well understood term to those in the field, that the term "atomic mass" was already in use (as it is currently defined) and that the term "relative atomic mass" was in some ways redundant. In 1979, in a compromise move, the definition was refined and the term "relative atomic mass" was introduced as a secondary synonym. Twenty years later the primacy of these synonyms was reversed and the term "relative atomic mass" is now the preferred term; however the "standard atomic weights" have maintained the same name. [5]

[edit] Table of standard atomic weights
For more accurate standard atomic weights, including uncertainties , see Periodic table (detailed).

Group → 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
↓ Period

1 H
1.008
He
4.003
2 Li
6.941 Be
9.012
B
10.81 C
12.01 N
14.01 O
16.00 F
19.00 Ne
20.18
3 Na
22.99 Mg
24.31
Al
26.98 Si
28.09 P
30.97 S
32.07 Cl
35.45 Ar
39.95
4 K
39.10 Ca
40.08 Sc
44.96 Ti
47.87 V
50.94 Cr
52.00 Mn
54.94 Fe
55.84 Co
58.93 Ni
58.69 Cu
63.55 Zn
65.39 Ga
69.72 Ge
72.61 As
74.92 Se
78.96 Br
79.90 Kr
83.80
5 Rb
85.47 Sr
87.62 Y
88.91 Zr
91.22 Nb
92.91 Mo
95.94 Tc
[98] Ru
101.07 Rh
102.91 Pd
106.42 Ag
107.87 Cd
112.41 In
114.82 Sn
118.71 Sb
121.76 Te
127.60 I
126.90 Xe
131.29
6 Cs
132.91 Ba
137.33 *
Hf
178.49 Ta
180.95 W
183.84 Re
186.21 Os
190.23 Ir
192.22 Pt
195.08 Au
196.97 Hg
200.59 Tl
204.38 Pb
207.2 Bi
208.98 Po
[209] At
[210] Rn
[222]
7 Fr
[223] Ra
[226] **
Rf
[263] Db
[262] Sg
[266] Bh
[264] Hs
[269] Mt
[268] Ds
[272] Rg
[272] Uub
[277] Uut
[284] Uuq
[289] Uup
[288] Uuh
[292] Uus
[291]‡ Uuo
[293]‡

* Lanthanides La
138.91 Ce
140.12 Pr
140.91 Nd
144.24 Pm
[145] Sm
150.36 Eu
151.96 Gd
157.25 Tb
158.93 Dy
162.50 Ho
164.93 Er
167.26 Tm
168.93 Yb
173.04 Lu
174.97
** Actinides Ac
[227] Th
232.04 Pa
231.04 U
238.03 Np
[237] Pu
[244] Am
[243] Cm
[247] Bk
[247] Cf
[251] Es
[252] Fm
[257] Md
[258] No
[259] Lr
[262]

Chemical series of the periodic table Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals
Poor metals Metalloids Nonmetals Halogens Noble gases

State at standard temperature and pressure (0 °C and 1 atm) Gases Liquids Solids
Natural occurrence Undiscovered Synthetic From decay Primordial

[edit] See also
Atomic mass unit
Isotope
Molecular mass
Jean Stas

[edit] External links
NIST relative atomic masses of all isotopes and the standard atomic weights of the elements
IUPAC goldbook definition of atomic mass
IUPAC goldbook definition of relative atomic mass & atomic weight
Tutorial on the concept and measurement of atomic mass
Atomic Weights and the International Committee — A Historical Review

[edit] References
^ IUPAC Definition of Atomic Mass
^ IUPAC Definition of Relative Atomic Mass
^ IUPAC Definition of Standard Atomic Weight
^ ATOMIC WEIGHTS OF THE ELEMENTS 2005 (IUPAC TECHNICAL REPORT), M. E. WIESER Pure Appl. Chem., V.78, pp. 2051, 2006
^ 'ATOMIC WEIGHT' -THE NAME, ITS HISTORY, DEFINITION, AND UNITS, P. DE BIEVRE and H. S. PEISER Pure&App. Chem., 64, 1535, 1992
Retrieved from "http://en.wikipedia.org/wiki/Atomic_mass"
Categories: Chemical properties | Mass | Stoichiometry

Color

2007-08-14T09:56:00.000-07:00

Color
From Wikipedia, the free encyclopedia
Jump to: navigation, search
This article is about the perceptual property. For usage of color in templates and Wikipedia pages, see Wikipedia:Color. For other uses, see Color (disambiguation).
"Coloration" redirects here. For the musical sense, see diatonic and chromatic. For coloration in animals, see animal colouration.

Color is an important part of the visual arts.Color or colour (see spelling differences) is the visual perceptual property corresponding in humans to the categories called red, yellow, white, etc. Color derives from the spectrum of light (distribution of light energy versus wavelength) interacting in the eye with the spectral sensitivities of the light receptors. Color categories and physical specifications of color are also associated with objects, materials, light sources, etc., based on their physical properties such as light absorption, reflection, or emission spectra.

Typically, only features of the composition of light that are detectable by humans (wavelength spectrum from 400 nm to 700 nm, roughly) are included, thereby objectively relating the psychological phenomenon of color to its physical specification. Because perception of color stems from the varying sensitivity of different types of cone cells in the retina to different parts of the spectrum, colors may be defined and quantified by the degree to which they stimulate these cells. These physical or physiological quantifications of color, however, do not fully explain the psychophysical perception of color appearance.

The science of color is sometimes called chromatics. It includes the perception of color by the human eye and brain, the origin of color in materials, color theory in art, and the physics of electromagnetic radiation in the visible range (that is, what we commonly refer to simply as light).

Periodic table block

2007-08-14T09:55:00.000-07:00

Periodic table block
From Wikipedia, the free encyclopedia
Jump to: navigation, search
A block of the periodic table of elements is a set of adjacent groups. The respective highest-energy electrons in each element in a block belong to the same atomic orbital type. Each block is named after its characteristic orbital; thus, the blocks are:

s-block
p-block
d-block
f-block
g-block (no elements belonging to the g-block have been observed)
The block names (s, p, d, f, g, h,...) are derived from the quality of the spectroscopic lines of the associated atomic orbitals: sharp, principal, diffuse and fundamental, the rest being named in alphabetical order.

[hide]v • d • ePeriodic tables
Layouts Standard · Vertical · Full names · Names and atomic masses · Text for last · Huge table · Metals and nonmetals · Blocks · Valences · Inline f-block · 218 elements · Electron configurations · Atomic masses · Electronegativities · Alternatives
Lists of elements Name · Atomic symbol · Atomic number · Boiling point · Melting point · Density · Atomic mass
Groups 1 · 2 · 3 · 4 · 5 · 6 · 7 · 8 · 9 · 10 · 11 · 12 · 13 · 14 · 15 · 16 · 17 · 18
Periods: 1 · 2 · 3 · 4 · 5 · 6 · 7 · 8
Series Alkalis · Alkaline earths · Lanthanides · Actinides · Transition metals · Poor metals · Metalloids · Nonmetals · Halogens · Noble gases
Blocks s-block · p-block · d-block · f-block · g-block

Retrieved from "http://en.wikipedia.org/wiki/Periodic_table_block"
Category: Periodic table

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
Afrikaans
العربية
Asturianu
Català
Deutsch
Español
Esperanto
Français
Hrvatski
Bahasa Indonesia
Íslenska
Lumbaart
Nederlands
日本語
‪Norsk (nynorsk)‬
Português
Română
Slovenčina
Suomi
Svenska
ไทย
中文

This page was last modified 22:12, 16 June 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Period

2007-08-14T09:52:00.000-07:00

Period (periodic table)
From Wikipedia, the free encyclopedia
Jump to: navigation, search
In the periodic table of the elements, a period is a horizontal row of the table.

The elements are laid out in a series rows so that those with similar properties line up in vertical columns: this reflects the periodic recurrence of similar properties as atomic weight increases.

Modern quantum mechanics explains these periodic trends in properties in terms of electron shells. As atomic number increases, electron shells are filled in roughly this order. The filling of each shell corresponds to a row in the table.

1s
2s 2p
3s 3p
4s 3d 4p
5s 4d 5p
6s 4f 5d 6p
7s 5f 6d 7p
8s 5g 6f 7d 8p
...
Hence the structure of the periodic table. Since the valence electrons determine chemical properties, those tend to be similar within periodic table groups.

Elements adjacent to one another within a group have similar physical properties, despite their significant differences in mass. Elements adjacent to one another within a period have similar mass but different properties.

[hide]v • d • ePeriodic tables
Layouts Standard · Vertical · Full names · Names and atomic masses · Text for last · Huge table · Metals and nonmetals · Blocks · Valences · Inline f-block · 218 elements · Electron configurations · Atomic masses · Electronegativities · Alternatives
Lists of elements Name · Atomic symbol · Atomic number · Boiling point · Melting point · Density · Atomic mass
Groups 1 · 2 · 3 · 4 · 5 · 6 · 7 · 8 · 9 · 10 · 11 · 12 · 13 · 14 · 15 · 16 · 17 · 18
Periods: 1 · 2 · 3 · 4 · 5 · 6 · 7 · 8
Series Alkalis · Alkaline earths · Lanthanides · Actinides · Transition metals · Poor metals · Metalloids · Nonmetals · Halogens · Noble gases
Blocks s-block · p-block · d-block · f-block · g-block

Retrieved from "http://en.wikipedia.org/wiki/Period_%28periodic_table%29"
Category: Periodic table

ViewsArticle Discussion Edit this page History Personal toolsSign in / create account Navigation
Main page
Contents
Featured content
Current events
Random article
interaction
About Wikipedia
Community portal
Recent changes
Contact Wikipedia
Make a donation
Help
Search
Toolbox
What links here
Related changes
Upload file
Special pages
Printable version
Permanent link
Cite this article
In other languages
Afrikaans
العربية
Asturianu
Български
Català
Deutsch
Español
Esperanto
Euskara
Français
한국어
Bahasa Indonesia
Íslenska
Italiano
Lumbaart
Magyar
Nederlands
日本語
‪Norsk (nynorsk)‬
Plattdüütsch
Polski
Português
Română
Русский
Slovenčina
Српски / Srpski
Srpskohrvatski / Српскохрватски
Basa Sunda
Suomi
Svenska
ไทย
Tiếng Việt
中文

This page was last modified 14:21, 21 June 2007. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.)
Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a US-registered 501(c)(3) tax-deductible nonprofit charity.
Privacy policy About Wikipedia Disclaimers

Group

2007-08-14T09:42:00.000-07:00

Group (periodic table)
From Wikipedia, the free encyclopedia
Jump to: navigation, search
A group, also known as a family, is a vertical column in the periodic table of the chemical elements. There are 18 groups in the standard periodic table.

The modern explanation of the pattern of the periodic table is that the elements in a group have similar configurations of the outermost electron shells of their atoms: as most chemical properties are dominated by the orbital location of the outermost electron. There are three ways of numbering the groups of the periodic table, one using Hindu-Arabic numerals and the other two using Roman numerals. The Roman numeral names are the original traditional names of the groups; the Arabic numeral names are those recommended by the International Union of Pure and Applied Chemistry (IUPAC) to replace the old names in an attempt to reduce the confusion generated by the two older, but mutually confusing, schemes.

There is considerable confusion surrounding the two old systems in use (old IUPAC and CAS) that combined the use of Roman numerals with letters. In the old IUPAC system the letters A and B were designated to the left (A) and right (B) part of the table, while in the CAS system the letters A and B were designated to main group elements (A) and transition elements (B). The former system was frequently used in Europe while the latter was most common in America. The new IUPAC scheme was developed to replace both systems as they confusingly used the same names to mean different things.

The periodic table groups are as follows (in the brackets are shown the old systems: European and American):

Group 1 (IA,IA): the alkali metals or hydrogen family/lithium family
Group 2 (IIA,IIA): the alkaline earth metals or beryllium family
Group 3 (IIIA,IIIB): the scandium family
Group 4 (IVA,IVB): the titanium family
Group 5 (VA,VB): the vanadium family
Group 6 (VIA,VIB): the chromium family
Group 7 (VIIA,VIIB): the manganese family
Group 8 (VIII): the iron family
Group 9 (VIII): the cobalt family
Group 10 (VIII): the nickel family
Group 11 (IB,IB): the coinage metals (not an IUPAC-recommended name) or copper family
Group 12 (IIB,IIB): the zinc family
Group 13 (IIIB,IIIA): the boron family
Group 14 (IVB,IVA): the carbon family
Group 15 (VB,VA): the pnictogens (not an IUPAC-recommended name) or nitrogen family
Group 16 (VIB,VIA): the chalcogens or oxygen family
Group 17 (VIIB,VIIA): the halogens or fluorine family
Group 18 (Group 0): the noble gases or helium family/neon family

--------------------------------------------------------------------------------

Note: Wikipedia style should be to replace the old names of the groups with the new IUPAC names throughout, with a historical mention of the old name where appropriate.

[hide]v • d • ePeriodic tables
Layouts Standard · Vertical · Full names · Names and atomic masses · Text for last · Huge table · Metals and nonmetals · Blocks · Valences · Inline f-block · 218 elements · Electron configurations · Atomic masses · Electronegativities · Alternatives
Lists of elements Name · Atomic symbol · Atomic number · Boiling point · Melting point · Density · Atomic mass
Groups 1 · 2 · 3 · 4 · 5 · 6 · 7 · 8 · 9 · 10 · 11 · 12 · 13 · 14 · 15 · 16 · 17 · 18
Periods: 1 · 2 · 3 · 4 · 5 · 6 · 7 · 8
Series Alkalis · Alkaline earths · Lanthanides · Actinides · Transition metals · Poor metals · Metalloids · Nonmetals · Halogens · Noble gases
Blocks s-block · p-block · d-block · f-block · g-block

Retrieved from "http://en.wikipedia.org/wiki/Group_%28periodic_table%29"
Categories: Chemical element groups | Periodic table

List of elements by symbol

2007-08-14T09:06:00.000-07:00

List of elements by symbol
From Wikipedia, the free encyclopedia
Jump to: navigation, search
Contents [hide]
1 Current symbols
2 Symbols not currently used
2.1 Pictographic symbols
3 Symbols for named isotopes
4 Other symbols that look like element symbols
5 Notes
6 References
7 External links

This is a list of chemical elements by symbol, including the current signification used to identify the chemical elements as recognized by the International Union of Pure and Applied Chemistry, as well as proposed and historical signs. Also given is each element's atomic number, atomic mass or most stable isotope, group and period numbers on the periodic table, and etymology of the symbol if not obvious from the current name of the element.

Chemical series of the periodic table Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals
Poor metals Metalloids Nonmetals Halogens Noble gases

[edit] Current symbols
Symbol Name Etymology of Symbol Atomic Number Relative Atomic Mass Group Period
Ac Actinium corruption of the Greek aktinos 89 [227][1] 7
Ag Silver Latin argentum 47 107.8682(2)[2] 11 5
Al Aluminium (Aluminum) Latin alumen 13 26.9815386(8) 13 3
Am Americium the Americas 95 [243][1] 7
Ar Argon Greek argon 18 39.948(1)[2] [3] 18 3
As Arsenic Greek arsenikos 33 74.92160(2) 15 4
At Astatine Greek astatos 85 [210][1] 17 6
Au Gold Latin aurum 79 196.966569(4) 11 6
B Boron borax 5 10.811(7)[2] [4] [3] 13 2
Ba Barium Greek barys 56 137.327(7) 2 6
Be Beryllium beryl 4 9.012182(3) 2 2
Bh Bohrium Neils Bohr 107 [264][1] 7 7
Bi Bismuth corruption of the German wissmuth 83 208.98040(1) 15 6
Bk Berkelium Berkeley, California 97 [247][1] 7
Br Bromine Greek bromos 35 79.904(1) 17 4
C Carbon Latin carbo 6 12.0107(8)[2] [3] 14 2
Ca Calcium Latin calx 20 40.078(4)[2] 2 4
Cd Cadmium corruption of the Greek kadmia 48 112.411(8)[2] 12 5
Ce Cerium Ceres 58 140.116(1)[2] 6
Cf Californium State and University of California 98 [251][1] 7
Cl Chlorine Greek chloros 17 35.453(2)[2] [4] [3] 17 3
Cm Curium Pierre and Marie Curie and the traditional -um ending 96 [247][1] 7
Co Cobalt corruption of the German kobold 27 58.933195(5) 9 4
Cr Chromium Greek chroma 24 51.9961(6) 6 4
Cs Caesium (Cesium) Latin caesius 55 132.9054519(2) 1 6
Cu Copper Latin Cuprum 29 63.546(3)[3] 11 4
Db Dubnium Dubna, Russia 105 [262][1] 5 7
Ds Darmstadtium Darmstadt, Germany 110 [271][1] 10 7
Dy Dysprosium Greek dysprositos 66 162.500(1)[2] 6
Er Erbium Ytterby, Sweden 68 167.259(3)[2] 6
Es Einsteinium Albert Einstein 99 [252][1] 7
Eu Europium Europe 63 151.964(1)[2] 6
F Fluorine Latin fluo 9 18.9984032(5) 17 2
Fe Iron Latin ferrum 26 55.845(2) 8 4
Fm Fermium Enrico Fermi 100 [257][1] 7
Fr Francium France 87 [223][1] 1 7
Ga Gallium Latin Gallia 31 69.723(1) 13 4
Gd Gadolinium gadolinite 64 157.25(3)[2] 6
Ge Germanium Germany 32 72.64(1) 14 4
H Hydrogen Greek hydror 1 1.00794(7)[2] [4] [3] 1 1
He Helium Greek helios 2 4.002602(2)[2] [3] 18 1
Hf Hafnium Latin Hafnia 72 178.49(2) 4 6
Hg Mercury Latin hydrargyrum 80 200.59(2) 12 6
Ho Holmium Latin Holmia 67 164.930 32(2) 6
Hs Hassium Hesse, Germany 108 [277][1] 8 7
I Iodine Greek ioeides 53 126.904 47(3) 17 5
In Indium indigo blue 49 114.818(3) 13 5
Ir Iridium Greek iris 77 192.217(3) 9 6
K Potassium (Kalium) Latin kalium 19 39.0983(1) 1 4
Kr Krypton Greek kryptos 36 83.798(2)[2] [4] 18 4
La Lanthanum Greek lanthanien 57 138.90547(7)[2] 6
Li Lithium Greek lithos 3 6.941(2)[2] [4] [3] [5] 1 2
Lr Lawrencium Ernest O. Lawrence 103 [262][1] 3 7
Lu Lutetium Latin Lutetia 71 174.967(1)[2] 3 6
Md Mendelevium Dmitri Mendeleyev 101 [258][1] 7
Mg Magnesium Magnesia, Greece 12 24.3050(6) 2 3
Mn Manganese Latin magnes 25 54.938045(5) 7 4
Mo Molybdenum Greek molybdos 42 95.94(2)[2] 6 5
Mt Meitnerium Lise Meitner 109 [268][1] 9 7
N Nitrogen Greek nitron 7 14.0067(2)[2] [3] 15 2
Na Sodium Latin natrium 11 22.98976928(2) 1 3
Nb Niobium Niobe 41 92.906 38(2) 5 5
Nd Neodymium Greek neos didymos 60 144.242(3)[2] 6
Ne Neon Greek neos 10 20.1797(6)[2] [4] 18 2
Ni Nickel German kupfernickel 28 58.6934(2) 10 4
No Nobelium Alfred Nobel 102 [259][1] 7
Np Neptunium Neptune 93 [237][1] 7
O Oxygen Greek oxys 8 15.9994(3)[2] [3] 16 2
Os Osmium Greek osme 76 190.23(3)[2] 8 6
P Phosphorus Greek phosphoros 15 30.973762(2) 15 3
Pa Protactinium Greek protos and actinium 91 231.03588(2)[1] 7
Pb Lead Latin plumbum 82 207.2(1)[2] [3] 14 6
Pd Palladium Pallas and the traditional -dium ending 46 106.42(1)[2] 10 5
Pm Promethium Prometheus 61 [145][1] 6
Po Polonium Poland 84 [210][1] 16 6
Pr Praseodymium Greek prasios 59 140.90765(2) 6
Pt Platinum Greek platina 78 195.084(9) 10 6
Pu Plutonium Pluto 94 [244][1] 7
Ra Radium Latin radius 88 [226][1] 2 7
Rb Rubidium Latin rubidus 37 85.4678(3)[2] 1 5
Re Rhenium German Rheinprovinz 75 186.207(1) 7 6
Rf Rutherfordium Ernest Rutherford 104 261[1] 4 7
Rg Roentgenium Wilhelm Conrad Röntgen 111 [272][1] 11 7
Rh Rhodium Greek rhodon 45 102.905 50(2) 9 5
Rn Radon radium and emanation 86 [220][1] 18 6
Ru Ruthenium Latin Ruthenia 44 101.07(2)[2] 8 5
S Sulfur (Sulphur) Latin sulfur 16 32.065(5)[2] [3] 16 3
Sb Antimony Latin stibium 51 121.760(1)[2] 15 5
Sc Scandium Scandinavia 21 44.955912(6) 3 4
Se Selenium Greek selene 34 78.96(3)[3] 16 4
Sg Seaborgium Glenn T. Seaborg 106 [266][1] 6 7
Si Silicon Latin silex 14 28.0855(3)[3] 14 3
Sm Samarium samarskite 62 150.36(2)[2] 6
Sn Tin Latin stannum 50 118.710(7)[2] 14 5
Sr Strontium Latin Strontian 38 87.62(1)[2] [3] 2 5
Ta Tantalum King Tantalus 73 180.94788(2) 5 6
Tb Terbium Ytterby, Sweden 65 158.92535(2) 6
Tc Technetium Greek technetos 43 [98][1] 7 5
Te Tellurium Greek tellus 52 127.60(3)[2] 16 5
Th Thorium Thor 90 232.03806(2)[1] [2] 7
Ti Titanium the Titans 22 47.867(1) 4 4
Tl Thallium Greek thallos 81 204.3833(2) 13 6
Tm Thulium Thule and the traditional -ium ending 69 168.93421(2) 6
U Uranium Uranus 92 238.02891(3)[1] [2] [4] 7
Uub Ununbium Latin uni, uni, and bi 112 [285][1] 12 7
Uuh Ununhexium Latin uni, uni, and Greek hex 116 [292][1] 16 7
Uuo Ununoctium Latin uni, uni, and oct 118 [294][1] 18 7
Uup Ununpentium Latin uni, uni, and Greek pent 115 [288][1] 15 7
Uuq Ununquadium Latin uni, uni, and quadr 114 [289][1] 14 7
Uut Ununtrium Latin uni, uni, and Greek tri 113 [284][1] 13 7
V Vanadium Vanadis 23 50.9415(1) 5 4
W Tungsten German wolfram 74 183.84(1) 6 6
Xe Xenon Greek xenos 54 131.293(6)[2] [4] 18 5
Y Yttrium Ytterby, Sweden 39 88.90585(2) 3 5
Yb Ytterbium Ytterby, Sweden 70 173.04(3)[2] 6
Zn Zinc German zin 30 65.409(4) 12 4
Zr Zirconium zircon 40 91.224(2)[2] 4 5
Chemical series of the periodic table Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals
Poor metals Metalloids Nonmetals Halogens Noble gases

[edit] Symbols not currently used
The following is a list of names formerly used or suggested for use in naming the elements listed above. Also included in this list are placeholder names and names given by discredited claimants for discovery. Names which retain local use (as with aluminum and jod) are not included in this table.

Chemical Symbol Name Atomic Number Origin of Symbol Why Not Used
A Argon 18 Current symbol is Ar. [6]
Ab Alabamine 85 Discredited claim to discovery of astatine. [7]
Am Alabamium 85 Discredited claim to discovery of astatine. [7]
An Athenium 99 Proposed name for einsteinium. [8]
Ao Ausonium 93 Discredited claim to discovery of neptunium. [7]
Az Azote 7 Former name for nitrogen. [6]
Bv Brevium 91 Former name for protactinium. [6]
Bz Berzelium 59 Suggested name for praseodymium. [8]
Cb Columbium 41 Former name of niobium. [6]
Cb Columbium 95 Suggested name for americium. [8]
Cp Cassiopeium 71 Former name for lutetium. [6]
Ct Centurium 100 Proposed name for fermium. [8]
Ct Celtium 72 Former name of hafnium. [6]
Da Danubium 43 Suggested name for technetium. [8]
Db Dubnium 104 Proposed name for rutherfordium. The symbol and name were used for element 105. [6] [8]
Eb Ekaboron 21 Name given by Mendeleev to an as of then undiscovered element. When discovered, scandium closely matched the prediction. [8] [9]
El Ekaaluminium 31 Name given by Mendeleev to an as of then undiscovered element. When discovered, gallium closely matched the prediction. [8] [9]
Em Emanation 86 Also called radium emanation, the name was originally given by Friedrich Ernst Dorn in 1900. In 1923, this element officially became radon (the name given at one time to 222Rn, an isotope identified in the decay chain of radium). [6]
Em Ekamangan 43 Name given by Mendeleev to an as of then undiscovered element. When discovered, technetium closely matched the prediction. [8] [9]
Es Ekasilicon 32 Name given by Mendeleev to an as of then undiscovered element. When discovered, germanium closely matched the prediction. [8] [9]
Es Esperium 94 Discredited claim to discovery of plutonium. [7]
Fa Francium 87 Current symbol is Fr. [6]
Fr Florentium 61 Discredited claim to discovery of promethium. [7]
Gl Glucinium 4 Former name of beryllium. [6]
Ha Hahnium 105 Proposed name for dubnium. [8]
Ha Hahnium 108 Proposed name for hassium. [8]
Il Illinium 61 Discredited claim to discovery of promethium. [7]
Jg Jargonium 72 Discredited claim to discovery of hafnium. [7]
Jo Joliotium 105 Proposed name for dubnium. [8]
Ku Kurchatovium 104 Proposed name for rutherfordium. [8]
Lw Lawrencium 103 Current symbol is Lr. [6]
M Muriaticum 17 Former name of chlorine. [6]
Ma Masurium 43 Disputed claim to discovery of technetium. [7]
Md Mendelevium 97 Proposed name for berkelium. The symbol and name were later used for element 101. [6] [8]
Me Mendelevium 68 Suggested name for erbium. [6]
Ms Masrium 49 Discredited claim of discovery of indium. [7]
Mt Meitnium 91 Suggested name for protactinium. [8]
Mv Mendelevium 101 Current symbol is Md. [6]
Ng Norwegium 72 Discredited claim to discovery of hafnium. [7]
Ni Niton 86 Former name for radon. [6]
No Norium 72 Discredited claim to discovery of hafnium. [7]
Ns Nielsbohrium 105 Proposed name for dubnium. [8]
Ns Nielsbohrium 107 Proposed name for bohrium. [8]
Nt Niton 86 Suggested name for radon. [8]
Nw Newtonium 67 Suggested name for holmium. [8]
Ny Neoytterbium 70 Former name of ytterbium. [6]
Od Odinium 62 Suggested name for samarium. [8]
Pc Policium 110 Proposed name for darmstadtium. [8]
Pe Pelopium 41 Former name for niobium. [6]
Po Potassium 19 Current symbol is K. [6]
Rf Rutherfordium 106 Proposed name for seaborgium. The symbol and name were instead used for element 104. [6] [8]
Sa Samarium 62 Current symbol is Sm. [6]
So Sodium 11 Current symbol is Na. [6]
Sp Spectrium 70 Suggested name for ytterbium. [8]
St Antimony 51 Current symbol is Sb. [6]
Tn Tungsten 74 Current symbol is W. [6]
Tu Thulium 69 Current symbol is Tm. [6]
Tu Tungsten 74 Current symbol is W. [6]
Ty Tyrium 60 Suggested name for neodymium. [8]
Unb Unnilbium 102 Temporary name given to nobelium until it was permanently named by IUPAC. [9]
Une Unnilennium 109 Temporary name given to meitnerium until it was permanently named by IUPAC. [9]
Unh Unnilhexium 106 Temporary name given to seaborgium until it was permanently named by IUPAC. [9]
Uno Unniloctium 108 Temporary name given to hassium until it was permanently named by IUPAC. [9]
Unp Unnilpentium 105 Temporary name given to dubnium until it was permanently named by IUPAC. [9]
Unq Unnilquadium 104 Temporary name given to rutherfordium until it was permanently named by IUPAC. [9]
Uns Unnilseptium 107 Temporary name given to bohrium until it was permanently named by IUPAC. [9]
Unt Unniltrium 103 Temporary name given to lawrencium until it was permanently named by IUPAC. [9]
Unu Unnilunium 101 Temporary name given to mendelevium until it was permanently named by IUPAC. [9]
Uun Ununnilium 110 Temporary name given to darmstadtium until it was permanently named by IUPAC. [9]
Uuu Unununium 111 Temporary name given to roentgenium until it was permanently named by IUPAC. [9]
Vi Virginium 87 Discredited claim to discovery of francium. [7]
Vm Virginium 87 Discredited claim to discovery of francium. [7]
Yt Yttrium 39 Current symbol is Y. [6]

[edit] Pictographic symbols
The following is a list of pictographic symbols employed to symbolize elements known since ancient times (for example to the alchemists). Not included in this list are symbolic representations of substances previously called elements (such as certain rare earth mineral blends and the classical elements fire and water of ancient philosophy) which are known today to be multi-atomic. Also not included are symbolic representations currently used for elements in other languages such as the Traditional Chinese elements. Modern alphabetic notation was introduced in 1814 by Jöns Jakob Berzelius.

Chemical Symbol Name Atomic Number Origin of Symbol Why Not Used
Hydrogen 1 Daltonian symbol circa 1808. [6]
Sulfur 16 Alchemical symbol. [6]
Pallas 16 Alchemical symbol. [6]
Sulfur 16 Daltonian symbol circa 1808. [6]
⊛ Magnesium 21 Alchemical symbol. [6]
♂ Mars 26 Alchemical symbol. [6]
Stellae Fixae 29 Pre-1500s alchemical symbol. [6]
♀ Venus 29 Alchemical symbol. [6]
Copper 29 Alchemical symbol. [6]
© Copper 29 Daltonian symbol circa 1808. [6]
Arsenic 33 Alchemical symbol. [6]
Luna 48 Alchemical symbol. [6]
Iupiter 50 Alchemical symbol. [6]
♁ Antimony 51 Alchemical symbol. [6]
Antimony 51 Alchemical symbol. [6]
☾ Platinum 78 Alchemical symbol. [6]
☉ Platinum 78 Alchemical symbol. [6]
Uranus 78 Alchemical symbol. [6]
☼ Sol 79 Alchemical symbol from the 1500s [6]
Sol 79 Alchemical symbol from 1700 through 1783. [6]
[1] Pisces 80 Pre-1500s alchemical symbol. [6]
Neptunus 80 Alchemical symbol from the 1600s. [6]
Mercurius 80 Alchemical symbol from 1700 through 1783. [6]
Saturnus 82 Alchemical symbol circa 1783. [6]
Taurus 83 Alchemical symbol. [6]

^ a b c d e f g h i j k l m n o p q r s t u v w x y z aa ab ac ad ae af ag ah ai aj The element does not have any stable nuclides, and a value in brackets, e.g. [209], indicates the mass number of the longest-lived isotope of the element. However, three elements, Thorium, Protactinium, and Uranium, have a characteristic terrestrial isotopic composition, and thus their atomic mass given.
^ a b c d e f g h i j k l m n o p q r s t u v w x y z aa ab ac ad ae af ag ah ai aj ak al am The isotopic composition of this element varies in some geological specimens, and the variation may exceed the uncertainty stated in the table.
^ a b c d e f g h i j k l m n o The isotopic composition varies in terrestrial material such that a more precise atomic weight can not be given.
^ a b c d e f g h The isotopic composition of the element can vary in commercial materials, which can cause the atomic weight to deviate significantly from the given value.
^ The atomic weight of commercial Lithium can vary between 6.939 and 6.996—analysis of the specific material is necessary to find a more accurate value.
^ a b c d e f g h i j k l m n o p q r s t u v w x y z aa ab ac ad ae af ag ah ai aj ak al am an ao ap aq ar as at au av aw ax ay az Name changed due to a standardization of, modernization of, or update to older previously used symbol.

^ a b c d e f g h i j k l m Name designated by discredited/disputed claimant.

^ a b c d e f g h i j k l m n o p q r s t u v w x y Name proposed prior to discovery/creation of element or prior to official re-naming of a placeholder name.

^ a b c d e f g h i j k l m n o Temporary placeholder name.

[edit] Symbols for named isotopes
The following is a list of isotopes of elements given in the previous tables which have been designated unique symbols. By this it is meant that a comprehensive list of current systematic symbols (in the uAtom form) are not included in the list and can instead be found in the Index to isotope pages chart. Some of the following symbols are no longer in use within the scientific community, however others (most notably some of the named isotopes of hydrogen) continue to be used today. Many of these symbols were designated during the early years of radiochemistry, and several isotopes (namely those in the actinium decay family, the radium decay family, and the thorium decay family) bear placeholder names using the early naming system devised by Ernest Rutherford. Although it is not an isotope, this is perhaps the most useful place to mention that H is the proposed symbol for the only created anti-element, antihydrogen.

Chemical Symbol Name Atomic Number Origin of Symbol
Ac Actinium 89 From the Greek aktinos. Name restricted at one time to 227Ac, an isotope of actinium. This named isotope later became the official name for element 89.
AcA Actinium A 84 From actinium and A. Placeholder name given at one time to 215Po, an isotope of polonium identified in the decay chain of actinium.
AcB Actinium B 82 From actinium and B. Placeholder name given at one time to 211Pb, an isotope of lead identified in the decay chain of actinium.
AcC Actinium C 83 From actinium and C. Placeholder name given at one time to 211Bi, an isotope of bismuth identified in the decay chain of actinium.
AcC' Actinium C' 84 From actinium and C'. Placeholder name given at one time to 211Po, an isotope of polonium identified in the decay chain of actinium.
AcC" Actinium C" 81 From actinium and C". Placeholder name given at one time to 207Tl, an isotope of thallium identified in the decay chain of actinium.
AcK Actinium K 87 Name given at one time to 223Fr, an isotope of francium identified in the decay chain of actinium.
AcU Actino-Uranium 92 Name given at one time to 235U, an isotope of uranium.
AcX Actinium X 88 Name given at one time to 223Ra, an isotope of radium identified in the decay chain of actinium.
An Actinon 86 From actinium and emanation. Name given at one time to 219Rn, an isotope of radon identified in the decay chain of actinium.
D Deuterium 1 From the Greek deuteros. Name given to 2H.
Io Ionium 90 Name given at one time to 230Th, an isotope of thorium identified in the decay chain of uranium.
MsTh1 Mesothorium 1 88 Name given at one time to 228Ra, an isotope of radium.
MsTh2 Mesothorium 2 89 Name given at one time to 228Ac, an isotope of actinium.
Pa Protactinium 91 From the Greek protos and actinium. Name restricted at one time to 231Pa, an isotope of protactinium. This named isotope later became the official name for element 91.
Ra Radium 88 From the Latin radius. Name restricted at one time to 226Ra, an isotope of radium. This named isotope later became the official name for element 88.
RaA Radium A 84 From radium and A. Placeholder name given at one time to 218Po, an isotope of polonium identified in the decay chain of radium.
RaB Radium B 82 From radium and B. Placeholder name given at one time to 214Pb, an isotope of lead identified in the decay chain of radium.
RaC Radium C 83 From radium and C. Placeholder name given at one time to 214Bi, an isotope of bismuth identified in the decay chain of radium.
RaC' Radium C' 84 From radium and C'. Placeholder name given at one time to 214Po, an isotope of polonium identified in the decay chain of radium.
RaC" Radium C" 81 From radium and C". Placeholder name given at one time to 210Tl, an isotope of thallium identified in the decay chain of radium.
RaD Radium D 82 From radium and D. Placeholder name given at one time to 210Pb, an isotope of lead identified in the decay chain of radium.
RaE Radium E 83 From radium and E. Placeholder name given at one time to 210Bi, an isotope of bismuth identified in the decay chain of radium.
RaE" Radium E" 81 From radium and E". Placeholder name given at one time to 206Tl, an isotope of thallium identified in the decay chain of radium.
RaF Radium F 84 From radium and F. Placeholder name given at one time to 210Po, an isotope of polonium identified in the decay chain of radium.
RdAc Radioactinium 90 Name given at one time to 227Th, an isotope of thorium.
RdTh Radiothorium 90 Name given at one time to 228Th, an isotope of thorium.
Rn Radon 86 From radium and emanation. Name restricted at one time to 222Rn, an isotope of radon identified in the decay chain of radium. This named isotope later became the official name for element 86 in 1923.
T Tritium 1 From the Greek tritos. Name given to 3H.
Th Thorium 90 After Thor. Name restricted at one time to 232Th, an isotope of thorium. This named isotope later became the official name for element 90.
ThA Thorium A 84 From thorium and A. Placeholder name given at one time to 216Po, an isotope of polonium identified in the decay chain of thorium.
ThB Thorium B 82 From thorium and B. Placeholder name given at one time to 212Pb, an isotope of lead identified in the decay chain of thorium.
ThC Thorium C 83 From thorium and C. Placeholder name given at one time to 212Bi, an isotope of bismuth identified in the decay chain of thorium.
ThC' Thorium C' 84 From thorium and C'. Placeholder name given at one time to 212Po, an isotope of polonium identified in the decay chain of thorium.
ThC" Thorium C" 81 From thorium and C". Placeholder name given at one time to 208Tl, an isotope of thallium identified in the decay chain of thorium.
ThX Thorium X 88 Name given at one time to 224Ra, an isotope of radium identified in the decay chain of thorium.
Tn Thoron 86 From thorium and emanation. Name given at one time to 220Rn, an isotope of radon identified in the decay chain of thorium.
UI Uranium I 92 Name given at one time to 238U, an isotope of uranium.
UII Uranium II 92 Name given at one time to 234U, an isotope of uranium.
UX1 Uranium X1 90 Name given at one time to 234Th, an isotope of thorium identified in the decay chain of uranium.
UX2 Uranium X2 91 Name given at one time to 234Pa, an isotope of protactinium identified in the decay chain of uranium.
UY Uranium Y 90 Name given at one time to 231Th, an isotope of thorium identified in the decay chain of uranium.
UZ Uranium Z 91 Name given at one time to 234Pa, an isotope of protactinium identified in the decay chain of uranium.

[edit] Other symbols that look like element symbols
Ab: albite
Ac: acetate - (also used for the element actinium: see above)
Ar: aryl - (also used for the element argon: see above)
Bn: benzyl
Bu: butyl
Bz: benzoyl - (also used for berzelium, an old suggested name for praseodymium).
Di: didymium - Rare earth metal that proved to be a mixture of the elements praseodymium and neodymium.
Dp: decipium - Rare earth metal that proved to be a mixture of the elements samarium, neodymium and praseodymium.
Et: ethyl
M: any metal atom
Me: methyl
Mu: muonium
Ph: phenyl
Pp: philippium - rare earth metal that proved to be a mixture of the elements holmium and samarium.
Pr: propyl - (also used for the element praseodymium: see above)
Ps: positronium
R: - some unspecified element or radical

[edit] Notes
^ a b c d e f g h i j k l m n o p q r s t u v w x y z aa ab ac ad ae af ag ah ai aj The element does not have any stable nuclides, and a value in brackets, e.g. [209], indicates the mass number of the longest-lived isotope of the element. However, three elements, Thorium, Protactinium, and Uranium, have a characteristic terrestrial isotopic composition, and thus their atomic mass given.
^ a b c d e f g h i j k l m n o p q r s t u v w x y z aa ab ac ad ae af ag ah ai aj ak al am The isotopic composition of this element varies in some geological specimens, and the variation may exceed the uncertainty stated in the table.
^ a b c d e f g h i j k l m n o The isotopic composition varies in terrestrial material such that a more precise atomic weight can not be given.
^ a b c d e f g h The isotopic composition of the element can vary in commercial materials, which can cause the atomic weight to deviate significantly from the given value.
^ The atomic weight of commercial Lithium can vary between 6.939 and 6.996—analysis of the specific material is necessary to find a more accurate value.
^ a b c d e f g h i j k l m n o p q r s t u v w x y z aa ab ac ad ae af ag ah ai aj ak al am an ao ap aq ar as at au av aw ax ay az Name changed due to a standardization of, modernization of, or update to older previously used symbol.

^ a b c d e f g h i j k l m Name designated by discredited/disputed claimant.

^ a b c d e f g h i j k l m n o p q r s t u v w x y Name proposed prior to discovery/creation of element or prior to official re-naming of a placeholder name.

^ a b c d e f g h i j k l m n o Temporary placeholder name.

[edit] References
Element name etymologies. Retrieved July 15, 2005.
Atomic Weights of the Elements 2001, Pure Appl. Chem. 75(8), 1107-1122, 2003. Retrieved June 30, 2005. Atomic weights of elements with atomic numbers from 1-109 taken from this source.
IUPAC Standard Atomic Weights Revised (2005).
WebElements Periodic Table. Retrieved June 30, 2005. Atomic weights of elements with atomic numbers 110-116 taken from this source.
Lapp, Ralph E. Matter. Life Science Library. New York: Time Incorporated. 1963.
Leighton, Robert B. Principles of Modern Physics. New York: McGraw-Hill. 1959.

[edit] External links
Berzelius’ List of Elements
History of IUPAC Atomic Weight Values (1883 to 1997)
[hide]v • d • ePeriodic tables
Layouts Standard · Vertical · Full names · Names and atomic masses · Text for last · Huge table · Metals and nonmetals · Blocks · Valences · Inline f-block · 218 elements · Electron configurations · Atomic masses · Electronegativities · Alternatives
Lists of elements Name · Atomic symbol · Atomic number · Boiling point · Melting point · Density · Atomic mass
Groups 1 · 2 · 3 · 4 · 5 · 6 · 7 · 8 · 9 · 10 · 11 · 12 · 13 · 14 · 15 · 16 · 17 · 18
Periods: 1 · 2 · 3 · 4 · 5 · 6 · 7 · 8
Series Alkalis · Alkaline earths · Lanthanides · Actinides · Transition metals · Poor metals · Metalloids · Nonmetals · Halogens · Noble gases
Blocks s-block · p-block · d-block · f-block · g-block
Retrieved from "http://en.wikipedia.org/wiki/List_of_elements_by_symbol"
Category: Chemical elements

Xenon

2007-08-14T08:56:00.000-07:00

Xenon
From Wikipedia, the free encyclopedia
Jump to: navigation, search
For other uses, see Xenon (disambiguation).
54 iodine ← xenon → caesium
Kr
↑
Xe
↓
Rn
Periodic Table - Extended Periodic Table

General
Name, Symbol, Number xenon, Xe, 54
Chemical series noble gases
Group, Period, Block 18, 5, p
Appearance colorless

Standard atomic weight 131.293(6) g·mol−1
Electron configuration [Kr] 4d10 5s2 5p6
Electrons per shell 2, 8, 18, 18, 8
Physical properties
Phase gas
Density (0 °C, 101.325 kPa)
5.894 g/L
Melting point 161.4 K
(-111.7 °C, -169.1 °F)
Boiling point 165.03 K
(-108.12 °C, -162.62 °F)
Triple point 161.405 K, 81.6 kPa[1]
Critical point 289.77 K, 5.841 MPa
Heat of fusion 2.27 kJ·mol−1
Heat of vaporization 12.64 kJ·mol−1
Heat capacity (25 °C) 20.786 J·mol−1·K−1
Vapor pressure P(Pa) 1 10 100 1 k 10 k 100 k
at T(K) 83 92 103 117 137 165

Atomic properties
Crystal structure cubic face centered
Oxidation states 0, +1, +2, +4, +6, +8
(rarely more than 0)
(weakly acidic oxide)
Electronegativity 2.6 (scale Pauling)
Ionization energies 1st: 1170.4 kJ/mol
2nd: 2046.4 kJ/mol
3rd: 3099.4 kJ/mol
Atomic radius (calc.) 108 pm
Covalent radius 130 pm
Van der Waals radius 216 pm
Miscellaneous
Magnetic ordering nonmagnetic
Thermal conductivity (300 K) 5.65 m W·m−1·K−1
Speed of sound (liquid) 1090 m/s
CAS registry number 7440-63-3
Selected isotopes
Main article: Isotopes of xenon iso NA half-life DM DE (MeV) DP
124Xe 0.1% 1.1×1017y ε ε no data 124Te
125Xe syn 16.9 h ε 1.652 125I
126Xe 0.09% Xe is stable with 72 neutrons
127Xe syn 36.4 d ε 0.662 127I
128Xe 1.91% Xe is stable with 74 neutrons
129Xe 26.4% Xe is stable with 75 neutrons
130Xe 4.1% Xe is stable with 76 neutrons
131Xe 21.29% Xe is stable with 77 neutrons
132Xe 26.9% Xe is stable with 78 neutrons
133Xe syn 5.243 d Beta- 0.427 133Cs
134Xe 10.4% Xe is stable with 80 neutrons
135Xe syn 9.10 h Beta- 1.16 135Cs
136Xe 8.9% 2.36×1021y Beta- no data 136Ba

References
Xenon (IPA: /ˈzɛnɒn, ˈziːnɒn/) is a chemical element that has the symbol Xe and atomic number 54. A colorless, heavy, odorless noble gas, xenon occurs in the earth's atmosphere in trace amounts and was part of the first noble gas compound synthesized.[2][3]

Contents [hide]
1 Notable characteristics
2 Applications
3 History
4 Occurrence
5 Compounds
6 Isotopes
7 Precautions
8 References
9 External links

[edit] Notable characteristics
Xenon is a member of the zero-valence elements that are called noble or inert gases; however, "inert" may not be an entirely accurate description of this chemical series since at least 80 compounds of this noble gas have been synthesized. In a gas filled tube, xenon emits a blue glow when the gas is excited by electrical discharge. Using gigapascals of pressure, xenon has been forced into a metallic phase.[4] Xenon can also form clathrates with water when atoms of it are trapped in a lattice of the water molecules.

[edit] Applications

Xenon in shaped Geissler tubes.This gas is most widely and most famously used in light-emitting devices called Xenon flash lamps, which are used in photographic flashes and stroboscopic lamps, to excite the active medium in lasers which then generate coherent light, to produce laser power for inertial confinement fusion, in bactericidal lamps (rarely), and in certain dermatological uses. Continuous, short-arc, high pressure Xenon arc lamps have a color temperature closely approximating noon sunlight and are used in solar simulators, typical 35mm and IMAX film projection systems, automotive HID headlights and other specialized uses. They are an excellent source of short wavelength ultraviolet radiation and they have intense emissions in the near infrared, which are used in some night vision systems. Other uses of Xenon:

Has been used as a general anesthetic, though it is expensive. Even so, anesthesia machines that can deliver Xenon are about to appear on the European market. [5] Two kinds of mechanism have been proposed. The first one involves the inhibition of the calcium ATPase pump in synaptic plasma membranes,[6] which results from a conformational change when xenon binds to nonpolar sites inside the protein.[7] The second mechanism focuses on the non-specific interactions between the anesthetic and the lipid membrane.[8]
In nuclear energy applications it is used in bubble chambers, probes, and in other areas where a high molecular weight and inert nature is a desirable quality.
Perxenates are used as oxidizing agents in analytical chemistry.
The isotope 133Xe is useful as a radioisotope.
Hyperpolarized MRI of the lungs and other tissues using 129Xe.[9]
Preferred fuel for Ion propulsion because of high atomic weight, ease of ionization, store as a liquid at near room temperature (but at high pressure) yet easily converts back into a gas to fuel the engine, inert nature makes it environmentally friendly and less corrosive to an ion engine than other fuels such as mercury or caesium. Europe's SMART-1 spacecraft utilized Xenon in its engines.[10]
Is used in protein crystallography. Applied at high pressure (~600 psi) to a protein crystal, xenon atoms bind in predominantly hydrophobic cavities, often creating a high quality, isomorphous, heavy-atom derivative..[11]
Xenon difluoride is used as an etchant for silicon, particularly in the production of microelectromechanical systems, or MEMS[12].

[edit] History
Xenon (from Greek ξένον meaning "strange one" or "stranger") was discovered in England by William Ramsay and Morris Travers on July 12, 1898, shortly after their discovery of the elements krypton and neon. They found it in the residue left over from evaporating components of liquid air.[13]

[edit] Occurrence
Xenon is a trace gas in Earth's atmosphere, occurring in one part in twenty million. The element is obtained commercially through extraction from the residues of liquefied air. This noble gas is naturally found in gases emitted from some mineral springs. Radioactive species of xenon, for example 133Xe and 135Xe are produced by neutron irradiation of fissionable material within nuclear reactors. Like the noble gas krypton, xenon can also be extracted by fractional distillation or liquefaction of liquid air and by selective adsorption on activated carbon.

[edit] Compounds

Xenon tetrafluorideXenon and the other noble gases had for a long time been considered to be completely chemically inert and not able to form compounds. However, in 1962 at the University of British Columbia, the first xenon compound, xenon hexafluoroplatinate, was synthesized by Neil Bartlett. Now, many compounds of xenon are known, including xenon difluoride, xenon tetrafluoride, xenon hexafluoride, xenon tetroxide, xenon hydrate, xenon deuterate, and sodium perxenate. A highly explosive compound xenon trioxide has also been made. There are at least 80 xenon compounds in which fluorine or oxygen is bonded to xenon. Some compounds of xenon are colored but most are colorless.

Recently at the University of Helsinki in Finland, a group of scientists (M. Räsänen et al.) prepared HXeH, HXeOH, and HXeCCH (xenon dihydride, xenon hydride-hydroxide, and hydroxenoacetylene). They are stable up to 40K.[14]

XeF4 crystals. 1962.
[edit] Isotopes
Main article: Isotopes of xenon
Naturally occurring xenon is made of seven stable and two slightly radioactive isotopes. Beyond these stable forms, there are 20 unstable isotopes that have been studied. 129Xe is produced by beta decay of 129I (half-life: 16 million years); 131Xe m, 133Xe, 133Xe m, and 135Xe are some of the fission products of both 235U and 239Pu, and therefore used as indicators of nuclear explosions.

The artificial isotope 135Xe is of considerable significance in the operation of nuclear fission reactors. 135Xe has a huge cross section for thermal neutrons, 2.65x106 barns, so it acts as a neutron absorber or "poison" that can slow or stop the chain reaction after a period of operation. This was discovered in the earliest nuclear reactors built by the American Manhattan Project for plutonium production. Fortunately the designers had made provisions in the design to increase the reactor's reactivity (the number of neutrons per fission that go on to fission other atoms of nuclear fuel).

Relatively high concentrations of radioactive xenon isotopes are also found emanating from nuclear reactors due to the release of this fission gas from cracked fuel rods or fissioning of uranium in cooling water. The concentrations of these isotopes are still usually low compared to naturally occurring radioactive noble gases such as 222Rn.

Because xenon is a tracer for two parent isotopes, Xe isotope ratios in meteorites are a powerful tool for studying the formation of the solar system. The I-Xe method of dating gives the time elapsed between nucleosynthesis and the condensation of a solid object from the solar nebula. Xenon isotopes are also a powerful tool for understanding terrestrial differentiation. Excess 129Xe found in carbon dioxide well gases from New Mexico was believed to be from the decay of mantle-derived gases soon after Earth's formation.[15]

[edit] Precautions
The gas can be safely kept in normal sealed glass containers at standard temperature and pressure. Xenon is non-toxic, but many of its compounds are toxic due to their strong oxidative properties.

The speed of sound in xenon is slower than that in air (due to the slower average speed of the heavy xenon atoms compared to nitrogen and oxygen molecules), so xenon lowers the resonant frequencies of the vocal tract when inhaled. This produces a characteristic lowered voice pitch, opposite the high-pitched voice caused by inhalation of helium. Like helium, xenon does not satisfy the body's need for oxygen and is a simple asphyxiant; consequently, many universities no longer allow the voice stunt as a general chemistry demonstration. As xenon is expensive, the gas sulfur hexafluoride, which is similar to xenon in molecular weight (146 vs 131), is generally used in this stunt, although it too is an asphyxiant.

A myth exists that xenon is too heavy for the lungs to expel unassisted, and that after inhaling xenon, it is necessary to bend over completely at the waist to allow the excess gas to "spill" out of the body. In fact, the lungs mix gases very effectively and rapidly, such that xenon would be purged from the lungs within a breath or two. There is, however, a danger associated with any heavy gas in large quantities: it may sit invisibly in a container, and if a person enters a container filled with an odorless, colorless gas, they may find themselves breathing it unknowingly. Xenon is rarely used in large enough quantities for this to be a concern, though the potential for danger exists any time a tank or container of xenon is kept in an unventilated space.

[edit] References
^ (2005) "Section 4, Properties of the Elements and Inorganic Compounds; Melting, boiling, triple, and critical temperatures of the elements", CRC Handbook of Chemistry and Physics, 85th edition, Boca Raton, Florida: CRC Press.
^ Los Alamos National Laboratory – Xenon
^ Thermophysical properties of neon, argon, krypton, and xenon / V. A. Rabinovich ... Theodore B. Selover, English-language edition ed, Washington [u.a.] Hemisphere Publ. Corp. [u.a.] , 1988. - XVIII (National standard reference data service of the USSR, You can now find Xenon at $60.00 per .077 pps
^ Caldwell, W. A.; Nguyen, J., Pfrommer, B., Louie, S., and Jeanloz, R. (1997). "Structure, bonding and geochemistry of xenon at high pressures". Science 277: 930-933.
^ Tonner PH. Xenon: one small step for anaesthesia... ? Current Opinion in Anaesthesiology. 2006 Aug;19(4):382-4.
^ Franks, John J. MD; Horn, Jean-Louis MD; Janicki, Piotr K. MD, PhD; Singh, Gurkeerat PhD. Halothane, Isoflurane, Xenon, and Nitrous Oxide Inhibit Calcium ATPase Pump Activity in Rat Brain Synaptic Plasma Membranes. Anesthesiology 1995, 82, 108-117.
^ Maria M. Lopez , Danuta Kosk-Kosicka. How Do Volatile Anesthetics Inhibit Ca2+-ATPases? Journal of Biological Chemistry 1995, 270, 28239-28245.
^ Thomas Heimburg and Andrew D. Jackson. The Thermodynamics of General Anesthesia. Biophysical Journal 2007, 92, 3159–3165. DOI:10.1529/biophysj.106.099754
^ Use of Xe in MRI
^ CNN Article regarding SMART-1 and Xenon
^ Xenon derivativation in protein crystllography at the Daresbury Synchrotron Radiation Source Laboratory, UK
^ Brazzle, J.D.; Dokmeci, M.R.; Mastrangelo, C.H.; Modeling and characterization of sacrificial polysilicon etching using vapor-phase xenon difluoride , 17th IEEE International Conference on Micro Electro Mechanical Systems (MEMS), 2004, pages 737-740.
^ Gagnon, Steve. It's Elemental - Xenon (English). Thomas Jefferson National Accelerator Facility. Retrieved on 16, 2007. Retrieved on June 2007.
^ See http://pubs.acs.org/cen/80th/noblegases.html in its paragraph starting "Many recent findings".
^ Boulos, M.S.; Manuel, O.K. (1971). "The xenon record of extinct radioactivities in the Earth.". Science 174: 1334-1336.

[edit] External links
Wikimedia Commons has media related to:
XenonLook up xenon in
Wiktionary, the free dictionary.WebElements.com – Xenon
Xenon as an anaesthetic
USGS Periodic Table - Xenon
Retrieved from "http://en.wikipedia.org/wiki/Xenon"
Categories: Chemical elements | Noble gases | Anesthetics | Xenon

Ununquadium

2007-08-14T08:55:00.000-07:00

Ununquadium
From Wikipedia, the free encyclopedia
Jump to: navigation, search
114 ununtrium ← ununquadium → ununpentium
Pb
↑
Uuq
↓
(Uhq)
Periodic Table - Extended Periodic Table

General
Name, Symbol, Number ununquadium, Uuq, 114
Chemical series presumably poor metals
Group, Period, Block 14, 7, p
Appearance unknown, probably silvery
white or metallic gray
Standard atomic weight (298) g·mol−1
Electron configuration perhaps [Rn] 5f14 6d10 7s2 7p2
(guess based on lead)
Electrons per shell 2, 8, 18, 32, 32, 18, 4
Phase presumably a solid
CAS registry number 54085-16-4
Selected isotopes
Main article: Isotopes of ununquadium iso NA half-life DM DE (MeV) DP
288Uuq syn 2.8 s

References
Ununquadium (IPA: /ˌjuːˌnʌnˈkwɒdiəm/), or eka-lead, is the temporary name of a radioactive chemical element in the periodic table that has the temporary symbol Uuq and has the atomic number 114.

Contents [hide]
1 History
2 Synthesis
2.1 Synthesis of Ununquadium-298
3 (more) Stable ununquadium
4 See also
5 References
6 External links

[edit] History
The discovery of ununquadium in December 1998 was reported in January 1999 by scientists at Dubna (Joint Institute for Nuclear Research) in Russia.[1] The same team produced another isotope of Uuq three months later[2] and confirmed the synthesis in 2004 and 2006.

In 2004 in the Joint Institute for Nuclear Research the synthesis of this element was confirmed by another method (the chemical identifying on final products of decay of element).

Ununquadium is a temporary IUPAC systematic element name. Some have termed it eka-lead, as its properties are conjectured to be similar to those of lead. It is expected to be a soft, dense metal that tarnishes in air, with a melting point around 200 degrees Celsius.

[edit] Synthesis
Ununquadium can be synthesized by bombarding plutonium-244 targets with calcium-48 heavy ion beams.

[edit] Synthesis of Ununquadium-298
Manufacturing ununquadium-298 would be very difficult, because nuclei summing to 114 protons and 184 neutrons are not available in weighable quantities.

However it may be possible to generate ununquadium-298, if nuclear transfer reactions can be achieved.[citation needed] One of these reactions may be

[edit] (more) Stable ununquadium
According to the island of stability theory, some nuclides around the area of 114 protons and 184 neutrons (i.e. isotope Uuq-298) can be expected to be relatively stable in comparison to the surrounding nuclides. Ununquadium does not occur naturally, so it is entirely synthesized in laboratories. All isotopes of ununquadium synthesized so far are neutron-poor. This means that they contain significantly fewer neutrons than 184, which is one of the magic number of neutrons that is believed to make the isotope more stable. Neutron-poor also indicates that the isotopes decay either by spontaneous fission producing a variety of radionuclides, positron emission or electron capture to yield element ununtrium. So far, all three that have been made so far have achieved spontanious fission in the first .0012 milliseconds, and therefore have never been able to be studied.

[edit] See also
Island of stability: Ununquadium – Unbinilium – Unbihexium
Lead
Periodic table (extended)
Isotopes of ununquadium

[edit] References
^ Oganessian, Yu. Ts.; et al. (October 1999). "Synthesis of Superheavy Nuclei in the 48Ca + 244Pu Reaction". Physical Review Letters 83: 3154. DOI:10.1103/PhysRevLett.83.3154.
^ Oganessian; et al. (July 1999). "Synthesis of nuclei of the superheavy element 114 in reactions induced by 48Ca". Nature 400: 242. DOI:10.1038/22281.

[edit] External links
Wikimedia Commons has media related to:
UnunquadiumWebElements.com - Uuq
Apsidium - Ununquadium
First postcard from the island of nuclear stability
Second postcard from the island of stability
Retrieved from "http://en.wikipedia.org/wiki/Ununquadium"
Categories: All articles with unsourced statements | Articles with unsourced statements since February 2007 | Chemical elements | Poor metals | Nuclear physics

Ununpentium

2007-08-14T08:54:00.000-07:00

Your continued donations keep Wikipedia running!
Ununpentium
From Wikipedia, the free encyclopedia
Jump to: navigation, search
115 ununquadium ← ununpentium → ununhexium
Bi
↑
Uup
↓
(Uhp)
Periodic Table - Extended Periodic Table

General
Name, Symbol, Number ununpentium, Uup, 115
Group, Period, Block 15, 7, p
Standard atomic weight (299) g·mol−1
Electron configuration perhaps [Rn] 5f14 6d10 7s2 7p3
(guess based on bismuth)
Electrons per shell 2, 8, 18, 32, 32, 18, 5
CAS registry number 54085-64-2
Selected isotopes
Main article: Isotopes of ununpentium iso NA half-life DM DE (MeV) DP
288Uup syn 87 ms

References
Ununpentium (IPA: /ˌjuːnʌnˈpɛntiəm/) is the temporary name of a synthetic superheavy element in the periodic table that has the temporary symbol Uup and has the atomic number 115. Multiple isotopes have been made by a fusion of calcium and americium (Uup-288 with the most neutrons). It can be referred to as eka-bismuth.

Element 115 also falls in the center of the theoretical island of stability. Although no stable isotopes have yet been found, conventional models predict that if stable isotopes of element 115 can be produced, they will most likely need the "magic number" of 184 neutrons, which would be Uup-299. The currently fabricated isotopes only had at most 173 neutrons (Uup-288).

Contents [hide]
1 History
2 Chemical properties
3 In popular culture
4 See also
5 References
6 External links

[edit] History
On February 2, 2004, synthesis of ununpentium and ununtrium were reported in Physical Review C by a team composed of Russian scientists at Dubna University's Joint Institute for Nuclear Research and American scientists at the Lawrence Livermore National Laboratory.[1][2]

The team reported that they bombarded americium (element 95) with calcium (element 20) to produce four atoms of ununpentium (element 115). These atoms, they report, alpha decayed to ununtrium (element 113) in approximately 100 milliseconds. The ununtrium produced then existed for 1.2 seconds before decaying into natural elements.

The synthesizing of the element was also reported by scientists of Japan.

In May 2006 in the Joint Institute for Nuclear Research the synthesis of this element was confirmed by another method (the chemical identifying on final products of decay of element).

Ununpentium is a temporary IUPAC systematic element name. Element 115 is also sometimes called eka-bismuth.

[edit] Chemical properties
For now element 115 has only been manufactured in the amount of a few atoms, so the chemistry of element 115 has yet to be researched, but chemistry and physics can tell us a lot about what to expect. Although element 115 is in the same group as bismuth, its chemistry will probably be strongly altered by relativistic effects.[3] One important predicted difference from bismuth is the presence of a stable oxidation state of +1, and a Uup+ ion with a chemistry similar to Tl+. There has been some experimental data for other superheavy elements, such as element 112, which seems to confirm relativistic effects for superheavy elements.

[edit] In popular culture
Ununpentium has been theorized to be inside the island of stability. This probably explains why it was mentioned regularly in popular culture, especially in UFO conspiracy theories.

See ununpentium's entries at fictional applications of real materials.
The most popular account of element 115, from Bob Lazar, is considered pseudoscience [1]. Although it is reasonable to suppose that element 115 will have unique properties, there is no openly available empirical evidence to back up Lazar's claims.

[edit] See also
Island of stability

[edit] References
^ Oganessian, Yu. Ts.; et al. (2004). "Experiments on the synthesis of element 115 in the reaction 243Am(48Ca,xn)291−x115". Physical Review C 69: 021601. DOI:10.1103/PhysRevC.69.021601.
^ Oganessian, Yu. Ts.; et al. (2005). "Synthesis of elements 115 and 113 in the reaction 243Am + 48Ca". Physical Review C 72: 034611. DOI:10.1103/PhysRevC.72.034611.
^ Keller, O. L., Jr.; C. W. Nestor, Jr. (1974). "Predicted properties of the superheavy elements. III. Element 115, Eka-bismuth". Journal of Physical Chemistry 78: 1945. DOI:10.1021/j100612a015.

[edit] External links
Wikimedia Commons has media related to:
UnunpentiumUut and Uup Add Their Atomic Mass to Periodic Table
Apsidium - Ununpentium
Superheavy elements
History & Etymology
[2]
Retrieved from "http://en.wikipedia.org/wiki/Ununpentium"
Categories: Chemical elements | Poor metals

Tungsten

2007-08-14T08:52:00.000-07:00

Tungsten
From Wikipedia, the free encyclopedia
Jump to: navigation, search
For other uses, see Tungsten (disambiguation).
74 tantalum ← tungsten → rhenium
Mo
↑
W
↓
Sg
Periodic Table - Extended Periodic Table

General
Name, Symbol, Number tungsten, W, 74
Chemical series transition metals
Group, Period, Block 6, 6, d
Appearance grayish white, lustrous

Standard atomic weight 183.84(1) g·mol−1
Electron configuration [Xe] 4f14 5d4 6s2
Electrons per shell 2, 8, 18, 32, 12, 2
Physical properties
Phase solid
Density (near r.t.) 19.25 g·cm−3
Liquid density at m.p. 17.6 g·cm−3
Melting point 3695 K
(3422 °C, 6192 °F)
Boiling point 5828 K
(5555 °C, 10031 °F)
Heat of fusion 52.31 kJ·mol−1
Heat of vaporization 806.7 kJ·mol−1
Heat capacity (25 °C) 24.27 J·mol−1·K−1
Vapor pressure P(Pa) 1 10 100 1 k 10 k 100 k
at T(K) 3477 3773 4137 4579 5127 5823

Atomic properties
Crystal structure cubic body centered
Oxidation states 6, 5, 4, 3, 2, 1, 0, −1
(mildly acidic oxide)
Electronegativity 2.36 (scale Pauling)
Ionization energies 1st: 770 kJ/mol
2nd: 1700 kJ/mol
Atomic radius 135 pm
Atomic radius (calc.) 193 pm
Covalent radius 146 pm
Miscellaneous
Magnetic ordering no data
Electrical resistivity (20 °C) 52.8 n Ω·m
Thermal conductivity (300 K) 173 W·m−1·K−1
Thermal expansion (25 °C) 4.5 µm·m−1·K−1
Speed of sound (thin rod) (r.t.) (annealed)
4620 m·s−1
Young's modulus 411 GPa
Shear modulus 161 GPa
Bulk modulus 310 GPa
Poisson ratio 0.28
Mohs hardness 7.5
Vickers hardness 3430 MPa
Brinell hardness 2570 MPa
CAS registry number 7440-33-7
Selected isotopes
Main article: Isotopes of tungsten iso NA half-life DM DE (MeV) DP
180W 0.12% 1.8×1018 y α 2.516 176Hf
181W syn 121.2 d ε 0.188 181Ta
182W 26.50% W is stable with 108 neutrons
183W 14.31% W is stable with 109 neutrons
184W 30.64% W is stable with 110 neutrons
185W syn 75.1 d β- 0.433 185Re
186W 28.43% W is stable with 112 neutrons

References
Tungsten (IPA: /ˈtʊŋstən/), also called wolfram (IPA: /ˈwʊlfrəm, -am/), is a chemical element that has the symbol W (New Latin: wolframium) and atomic number 74. A very hard, heavy, steel-gray to white transition metal, tungsten is found in several ores including wolframite and scheelite and is remarkable for its robust physical properties, especially the fact that it has the highest melting point of all the non-alloyed metals and the second highest of all the elements after carbon. The pure form is used mainly in electrical applications but its many compounds and alloys are widely used in many applications, most notably in light bulb filaments, in X-ray tubes (as both the filament and target), and in superalloys. Tungsten is the only metal from the third transition series that is known to occur in biomolecules.

Contents [hide]
1 Notable characteristics
2 Applications
3 History
4 Biological role
5 Production trends
6 Compounds
6.1 Aqueous polyoxoanions
7 Isotopes
8 References
9 See also
10 External links

[edit] Notable characteristics
Pure tungsten is steel-gray to tin-white and is a hard metal. Tungsten can be cut with a hacksaw when it is very pure (it is brittle and hard to work when impure) and is otherwise worked by forging, drawing, extruding, or sintering. This element has the highest melting point (3422 °C) (6192 °F), lowest vapor pressure and the highest tensile strength at temperatures above 1650 °C (3000 °F) of all metals. Tungsten has the lowest coefficient of thermal expansion of any pure metal. Its corrosion resistance is excellent and it can be attacked only slightly by most mineral acids. Tungsten metal forms a protective oxide when exposed to air but can be oxidized at high temperature. Steel alloyed with small quantities of tungsten greatly increases its toughness.

[edit] Applications
This section may require cleanup to meet Wikipedia's quality standards.
Please discuss this issue on the talk page, and/or replace this tag with a more specific message. Editing help is available.
This section has been tagged since February 2007.
Tungsten is a metal with a wide range of uses, the largest of which is as tungsten carbide (W2C, WC) in cemented carbides. Cemented carbides (also called hardmetals) are wear-resistant materials used by the metalworking, mining, petroleum and construction industries. Tungsten is widely used in light bulb and vacuum tube filaments, as well as electrodes, because it can be drawn into very thin wire with a high melting point. Other uses:

Its high melting point makes tungsten suitable for aerospace and high temperature uses which include electrical, heating, and welding applications, notably in the GTAW process (also called TIG welding).
Hardness and density properties make this metal ideal for making heavy metal alloys that are used in armament, heat sinks, and high density applications, such as weights, counterweights, ballast keels for yachts and tail ballast for commercial aircraft.
The high density makes it an ideal ingredient for darts, normally 80% and sometimes up to 97%. This allows darts containing tungsten to have a smaller diameter than those of other metals at the same weight, permitting tighter groupings.
High speed steel contains tungsten and some tungsten steels contain as much as 18% tungsten.
Superalloys containing tungsten are used in turbine blades and wear resistant parts and coatings. Examples are Hastelloy and Stellite.
Tungsten powder is used as a filler material in thermoplastic composites which are used as a nontoxic substitute for lead, in bullets, shot, and radiation shields.
Tungsten chemical compounds are used in catalysts, inorganic pigments, and tungsten disulfide high-temperature lubricants which are stable to 500 °C (930 °F).
Since this element's thermal expansion is similar to borosilicate glass, it is used for making glass-to-metal seals.
It is used in kinetic energy penetrators, usually alloyed with nickel and iron or cobalt, to form heavy alloys, used as an alternative to depleted uranium.
Tungsten is used as an interconnect material in integrated circuits. Contact holes are etched in silicon dioxide dielectric material, filled with tungsten and polished to form connections to transistors. Typical contact holes can be as small as 65 nm.
Tungsten carbide is one of the hardest carbides and is used in machine tools such as make milling and turning tools, and used together with cobalt and carbon is often the best choice for such applications.
Used extensively for shielding in the radiopharmaceutical industry. It is often employed when transporting individual FDG doses (called 'pigs') - the high energy of fluorine-18 makes lead much less effective.
Tungsten is used in the emitters of focused ion beam and electron microscopes.
Tungsten is also beginning to be used in jewelry. Its hardness makes it ideal for rings that will never scratch, are hypoallergenic and will not need polishing. This property is especially useful in designs with a brushed finish.
Also used in fishing lures like the Mormyshka.
Miscellaneous: Oxides are used in ceramic glazes and calcium/magnesium tungstates are used widely in fluorescent lighting. Crystal tungstates are used as scintillation detectors in nuclear physics and nuclear medicine. The metal is also used in X-ray targets and heating elements for electrical furnaces. Salts that contain tungsten are used in the chemical and tanning industries. Tungsten 'bronzes' (so-called due to the colour of the tungsten oxides) along with other compounds are used in paints. Some types of strings for musical instruments are wound with tungsten wire.

Closeup of a tungsten filament inside a halogen lamp.
[edit] History
Tungsten (Swedish tung sten meaning "heavy stone"), even though the current name for the element in Swedish is wolfram (sometimes spelled in Swedish as volfram), from the denomination volf rahm by Wallerius in 1747, translated from the description by Agricola in 1546 as Lupi spuma, meaning "wolf's froth" after the way tin is eaten up like a wolf after sheep in the process of its extraction[1].

It was first hypothesized to exist by Peter Woulfe in 1779 who examined wolframite and concluded that it must contain a new substance. In 1781 Carl Wilhelm Scheele ascertained that a new acid could be made from tungstenite. Scheele and Torbern Bergman suggested that it could be possible to obtain a new metal by reducing tungstic acid. In 1783 José and Fausto Elhuyar found an acid in wolframite that was identical to tungstic acid. In Spain later that year the brothers succeeded in isolating tungsten through reduction of this acid with charcoal. They are credited with the discovery of the element.

In World War II, tungsten played an enormous role in background political dealings. Portugal, as the main European source of the element, was put under pressure from both sides, because of its sources of wolframite ore. The resistance to high temperatures, as well as the extreme strength of its alloys, made the metal into a very important raw material for the weaponry industry.

[edit] Biological role
Tungsten is an essential nutrient for some organisms.

Enzymes called oxidoreductases use tungsten in a way that is similar to molybdenum by using it in a tungsten-pterin complex.

On August 20, 2002, officials representing the U.S.-based Centers for Disease Control and Prevention announced that urine tests on leukemia patient families and control group families in the Fallon, Nevada area had shown elevated levels of the metal tungsten in the bodies of both groups.[2] Sixteen recent cases of cancer in children were discovered in the Fallon area which has now been identified as a cancer cluster, (it should be noted, however, that the majority of the cancer victims are not long time residents of Fallon). Dr. Carol H. Rubin, a branch chief at the CDC, said data demonstrating a link between tungsten and leukemia is not available at present.[3]

[edit] Production trends
The quality of this article or section may be compromised by "weasel words".
You can help Wikipedia by removing weasel words.

Tungsten output in 2005Tungsten is found in the minerals wolframite (iron-manganese tungstate, FeWO4/MnWO4), scheelite (calcium tungstate, CaWO4), ferberite and hübnerite. There are important deposits of these minerals in China (with about 80% world share), Russia, Austria and Portugal, reports the British Geological Survey. The metal is commercially produced by reducing tungsten oxide with hydrogen or carbon.

World tungsten reserves have been estimated at 7 million t W. Unfortunately, most of these reserves are not economically workable so far. At our current annual consumption rate, these reserves will only last for about 140 years. According to further estimates, it has been suggested that 30% of the reserves are Wolframite and 70% are Scheelite ores. Another factor that controls the tungsten supply is scrap recycling of tungsten and it has been proven to be a very valuable raw material in comparison to ore.

[edit] Compounds
The most common formal oxidation state of tungsten is +6, but it exhibits all oxidation states from -1 to +6. [1] Tungsten typically combines with oxygen to form the yellow tungstic oxide, WO3, which dissolves in aqueous alkaline solutions to form tungstate ions, WO42−.

[edit] Aqueous polyoxoanions
Aqueous tungstate solutions are noted for the formation of polyoxoanions under neutral and acidic conditions. As tungstate is progressively treated with acid, it first yields the soluble, metastable "paratungstate A" anion, W7O246−, which over hours or days converts to the less soluble "paratungstate B" anion, H2W12O4210−. Further acidification produces the very soluble metatungstate anion, H2W12O406−, after equilibrium is reached. The metatungstate ion exists as a symmetric cluster of twelve tungsten-oxygen octahedra known as the "Keggin" anion. Many other polyoxoanions exist as metastable species. The inclusion of a different atom such as phosphorus in place of the two central hydrogens in metatungstate produces a wide variety of the so-called heteropolyanions.

See also tungsten compounds.

[edit] Isotopes
Main article: isotopes of tungsten
Naturally occurring tungsten consists of five isotopes whose half-lives are so long that they can be considered stable. All can decay into isotopes of element 72 (hafnium) by alpha emission; 180W has been observed to have a half life of 1.8 +- 0.2 Ea. The other naturally occurring isotopes have not been observed to decay, constraining their half-lives to be: 182W, T1/2 > 8.3 Ea; 184W, T1/2 > 29 Ea; 185W, T1/2 > 13 Ea; 186W, T1/2 > 27 Ea. [2] On average, two alpha decays of 180W occur in one gram of natural tungsten per year.

27 artificial radioisotopes of tungsten have been characterized, the most stable of which are 181W with a half-life of 121.2 days, 185W with a half-life of 75.1 days, 188W with a half-life of 69.4 days and 178W with a half-life of 21.6 days. All of the remaining radioactive isotopes have half-lives of less than 24 hours, and most of these have half-lives that are less than 8 minutes. Tungsten also has 4 meta states, the most stable being 179mW (t½ 6.4 minutes).

[edit] References
Los Alamos National Laboratory - Tungsten
DC/AC Circuits and Electronics: Principles & Applications by Robert K. Herrick, Published by Delmar Learning 2003 for Purdue University
^ Emsley, John (2000). The Elements, 3rd edition.
^ National Nuclear Data Center table of nuclides, http://www.nndc.bnl.gov/chart/

[edit] See also
Field emission gun
Oliver Sacks: "Uncle Tungsten: Memories of a Chemical Boyhood"

[edit] External links
Wikimedia Commons has media related to:
TungstenLook up tungsten in
Wiktionary, the free dictionary.Tungsten Disulfide Applications WS2
WebElements.com – Tungsten
Properties, Photos, History, MSDS
ScienceLab.com – Tungsten
Picture in the collection from Heinrich Pniok
Elementymology & Elements Multidict by Peter van der Krogt – Tungsten
Detection of the Natural Alpha Decay of Tungsten
Retrieved from "http://en.wikipedia.org/wiki/Tungsten"
Categories: Cleanup from February 2007 | Chemical elements | Transition metals | Refractory materials | Tungsten

Thorium

2007-08-14T08:50:00.000-07:00

Thorium
From Wikipedia, the free encyclopedia
Jump to: navigation, search
90 actinium ← thorium → protactinium
Ce
↑
Th
↓
(Uqn)
Periodic Table - Extended Periodic Table

General
Name, Symbol, Number thorium, Th, 90
Chemical series Actinides
Group, Period, Block n/a, 7, f
Appearance silvery white
Standard atomic weight 232.03806(2) g·mol−1
Electron configuration [Rn] 6d2 7s2
Electrons per shell 2, 8, 18, 32, 18, 10, 2
Physical properties
Phase solid
Density (near r.t.) 11.7 g·cm−3
Melting point 2115 K
(1842 °C, 3348 °F)
Boiling point 5061 K
(4788 °C, 8650 °F)
Heat of fusion 13.81 kJ·mol−1
Heat of vaporization 514 kJ·mol−1
Heat capacity (25 °C) 26.230 J·mol−1·K−1
Vapor pressure P(Pa) 1 10 100 1 k 10 k 100 k
at T(K) 2633 2907 3248 3683 4259 5055

Atomic properties
Crystal structure cubic face centered
Oxidation states 4
(weakly basic oxide)
Electronegativity 1.3 (scale Pauling)
Ionization energies
(more) 1st: 587 kJ·mol−1
2nd: 1110 kJ·mol−1
3rd: 1930 kJ·mol−1
Atomic radius 180 pm
Miscellaneous
Magnetic ordering no data
Electrical resistivity (0 °C) 147 nΩ·m
Thermal conductivity (300 K) 54.0 W·m−1·K−1
Thermal expansion (25 °C) 11.0 µm·m−1·K−1
Speed of sound (thin rod) (20 °C) 2490 m/s
Young's modulus 79 GPa
Shear modulus 31 GPa
Bulk modulus 54 GPa
Poisson ratio 0.27
Mohs hardness 3.0
Vickers hardness 350 MPa
Brinell hardness 400 MPa
CAS registry number 7440-29-1
Selected isotopes
Main article: Isotopes of thorium iso NA half-life DM DE (MeV) DP
228Th syn 1.9116 years α 5.520 224Ra
229Th syn 7340 years α 5.168 225Ra
230Th syn 75380 years α 4.770 226Ra
231Th trace 25.5 hours β 0.39 231Pa
232Th 100% 1.405×1010 years α 4.083 228Ra
234Th trace 24.1 days β 0.27 234Pa

References
Thorium (IPA: /ˈθɔːriəm/) is a chemical element in the periodic table that has the symbol Th and atomic number 90. As a naturally occurring, slightly radioactive metal, it has been considered as an alternative nuclear fuel to uranium.

Contents [hide]
1 Notable characteristics
2 Applications
3 History
4 Occurrence
4.1 Distribution
5 Thorium as a nuclear fuel
6 Isotopes
7 Precautions
8 Thorium Extraction
9 See also
10 References
11 External links

[edit] Notable characteristics
When pure, thorium is a silvery white metal that retains its luster for several months. However, when it is exposed to oxygen, thorium slowly tarnishes in air, becoming grey and eventually black. Thorium dioxide (ThO2), also called thoria, has the highest melting point of any oxide (3300°C)[1]. When heated in air, thorium metal turnings ignite and burn brilliantly with a white light.

Thorium has the largest liquid range of any element: 2946°C (2946 K) between the melting point and boiling point.

See Actinides in the environment for details of the environmental aspects of thorium.

[edit] Applications
Applications of thorium:

As an alloying element in magnesium, used in aircraft engines, imparting high strength and creep resistance at elevated temperatures.
Thorium is used to coat tungsten wire used in electronic equipment, improving the electron emission of heated cathodes.
Thorium has been used in gas tungsten arc welding electrodes and heat-resistant ceramics.
Uranium-thorium age dating has been used to date hominid fossils.
As a fertile material for producing nuclear fuel. In particular, the proposed energy amplifier reactor design would employ thorium. Since thorium is more abundant than uranium, some nuclear reactor designs incorporate thorium in their fuel cycle.
Thorium is a very effective radiation shield, although it has not been used for this purpose as much as lead or depleted uranium.
Thorium may be used in nuclear reactors instead of uranium as fuel. This produces less transuranic waste.
Applications of thorium dioxide (ThO2):

Mantles in portable gas lights. These mantles glow with a dazzling light (unrelated to radioactivity) when heated in a gas flame.
Used to control the grain size of tungsten used for electric lamps.
Used for high-temperature laboratory crucibles.
Added to glass, it helps create glasses of a high refractive index and with low dispersion. Consequently, they find application in high-quality lenses for cameras and scientific instruments.
Has been used as a catalyst:
In the conversion of ammonia to nitric acid.
In petroleum cracking.
In producing sulfuric acid.
Thorium dioxide is the active ingredient of Thorotrast, which was used as part of X-ray diagnostics. This use has been abandoned due to the carcinogenic nature of Thorotrast.

[edit] History
M. T. Esmark found a black mineral on Løvøy Island, Norway and gave a sample to Professor Jens Esmark, a noted mineralogist who was not able to identify it so he sent a sample to the Swedish chemist Jöns Jakob Berzelius for examination in 1828.[2] Berzelius analysed it and named it after Thor, the Norse god of thunder. The metal had virtually no uses until the invention of the gas mantle in 1885.

The crystal bar process (or Iodide process) was discovered by Anton Eduard van Arkel and Jan Hendrik de Boer in 1925 to produce high-purity metallic thorium. [3]

The name ionium was given early in the study of radioactive elements to the 230Th isotope produced in the decay chain of 238U before it was realized that ionium and thorium were chemically identical. The symbol Io was used for this supposed element.

[edit] Occurrence

Monazite, a rare-earth-and-thorium-phosphate mineral, is the primary source of the world's thoriumThorium is found in small amounts in most rocks and soils, where it is about three times more abundant than uranium, and is about as common as lead. Soil commonly contains an average of around 12 parts per million (ppm) of thorium. Thorium occurs in several minerals, the most common being the rare earth-thorium-phosphate mineral, monazite, which contains up to about 12% thorium oxide. There are substantial deposits in several countries. 232Th decays very slowly (its half-life is about three times the age of the earth) but other thorium isotopes occur in the thorium and uranium decay chains. Most of these are short-lived and hence much more radioactive than 232Th, though on a mass basis they are negligible. India is believed to have 25% of the world's Thorium reserves. [4]

See also thorium minerals.

[edit] Distribution
Present knowledge of the distribution of Thorium resources is poor because of the relatively low-key exploration efforts arising out of insignificant demand.[5] Under the prevailing estimate, Australia and India have particularly large reserves of thorium.

The prevailing estimate of the economically available thorium reserves comes from the US Geological Survey, Mineral Commodity Summaries (1997-2006):[6][7]
Country Th Reserves (tonnes) Th Reserve Base (tonnes)
Australia 300,000 340,000
India 290,000 300,000
Norway 170,000 180,000
United States 160,000 300,000
Canada 100,000 100,000
South Africa 35,000 39,000
Brazil 16,000 18,000
Malaysia 4,500 4,500
Other Countries 95,000 100,000
World Total 1,200,000 1,400,000

Another estimate of Reasonably Assured Reserves (RAR) and Estimated Additional Reserves (EAR) of thorium comes from OECD/NEA, Nuclear Energy, "Trends in Nuclear Fuel Cycle", Paris, France (2001)[8]
Country RAR Th (tonnes) EAR Th (tonnes)
Brazil 606,000 700,000
Turkey 380,000 500,000
India 319,000 -
United States 137,000 295,000
Norway 132,000 132,000
Greenland 54,000 32,000
Canada 45,000 128,000
Australia 19,000 -
South Africa 18,000 -
Egypt 15,000 309,000
Other Countries 505,000 -
World Total 2,230,000 2,130,000

The two sources vary wildly for countries such as Brazil, Turkey, and Australia.

[edit] Thorium as a nuclear fuel

Thorium metal foil (approximately 0.5 mm thick) sealed in a glass ampoule under an argon atmosphere to prevent oxidationThorium, as well as uranium and plutonium, can be used as fuel in a nuclear reactor. Although not fissile itself, 232Th will absorb slow neutrons to produce uranium-233 (233U), which is fissile. Hence, like 238U, it is fertile. In one significant respect 233U is better than the other two fissile isotopes used for nuclear fuel, 235U and plutonium-239 (239Pu), because of its higher neutron yield per neutron absorbed. Given a start with some other fissile material (235U or 239Pu), a breeding cycle similar to, but more efficient than that currently possible with the 238U-to-239Pu cycle (in slow-neutron reactors), can be set up. The 232Th absorbs a neutron to become 233Th which normally decays to protactinium-233 (233Pa) and then 233U. The irradiated fuel can then be unloaded from the reactor, the 233U separated from the thorium (a relatively simple process since it involves chemical instead of isotopic separation), and fed back into another reactor as part of a closed nuclear fuel cycle.

Problems include the high cost of fuel fabrication due partly to the high radioactivity of 233U which is a result of its contamination with traces of the short-lived 232U; the similar problems in recycling thorium due to highly radioactive 228Th; some weapons proliferation risk of 233U; and the technical problems (not yet satisfactorily solved) in reprocessing. Much development work is still required before the thorium fuel cycle can be commercialised, and the effort required seems unlikely while (or where) abundant uranium is available.

Nevertheless, the thorium fuel cycle, with its potential for breeding fuel without fast neutron reactors, holds considerable potential long-term benefits. Thorium is significantly more abundant than uranium, and is a key factor in sustainable nuclear energy.

India, having about 25% of the world's reserves [4], has planned its nuclear power program to eventually use thorium exclusively, phasing out uranium as a feed stock. This ambitious plan uses both fast and thermal breeder reactors. The Advanced Heavy Water Reactor and KAMINI reactor are efforts in this direction.

In 2007, Norway was debating whether or not to focus on Thorium plants.

The primary fuel of the HT3R Project in Odessa, TX, will be Ceramic-coated thorium beads.

[edit] Isotopes
Main article: isotopes of thorium
Naturally occurring thorium is composed of one isotope: 232Th. Twenty-seven radioisotopes have been characterized, with the most abundant and/or stable being 232Th with a half-life of 14.05 billion years, 230Th with a half-life of 75,380 years, 229Th with a half-life of 7340 years, and 228Th with a half-life of 1.92 years. All of the remaining radioactive isotopes have half-lives that are less than thirty days and the majority of these have half-lives that are less than ten minutes. One isotope, 229Th, has a nuclear isomer (or metastable state) with a remarkably low excitation energy of 3.5 eV. [9]

The known isotopes of thorium range in atomic weight from 210 u (210Th)[10] to 236 u (236Th).

[edit] Precautions
Powdered thorium metal is often pyrophoric and should be handled carefully.

Natural thorium decays very slowly compared to many other radioactive materials, and the alpha radiation emitted cannot penetrate human skin. Owning and handling small amounts of thorium, such as a gas mantle, is considered safe if care is taken not to ingest the thorium -- lungs and other internal organs can be penetrated by alpha radiation. Exposure to aerosolized thorium can lead to increased risk of cancers of the lung, pancreas and blood. Exposure to thorium internally leads to increased risk of liver diseases. This element has no known biological role. See also Thorotrast.

[edit] Thorium Extraction
Thorium has been extracted chiefly from monazite through a multi-stage process. In the first stage, the monazite sand is dissolved in an inorganic acid such as sulfuric acid (H2SO4). In the second, the Thorium is extracted into an organic phase containing an amine. Next it is separated or "stripped" using an anion such as nitrate, chloride, hydroxide, or carbonate, returning the thorium to an aqueous phase. Finally, the thorium is precipitated and collected.[11]

[edit] See also
David Hahn, who produced small quantities of fissionable material in his backyard.
Periodic table
Nuclear reactor
Decay chain
Sylvania Electric Products explosion
Thorium's entries at fictional applications of real materials.

[edit] References
^ Emsley, John (2001). Nature's Building Blocks, (Hardcover, First Edition), Oxford University Press, page 441. ISBN 0198503407.
^ Thorium. BBC.co. Retrieved on 2007-01-18.
^ van Arkel, A.E.; de Boer, J.H. (1925). "Preparation of pure titanium, zirconium, hafnium, and thorium metal". Zeitschrift für Anorganische und Allgemeine Chemie 148: 345-350. Retrieved on 2006-05-06.
^ a b US approves Indian nuclear deal. BBC News (2006-12-09).
^ K.M.V. Jayaram. An Overview of World Thorium Resources, Incentives for Further Exploration and Forecast for Thorium Requirements in the Near Future.
^ U.S. Geological Survey, Mineral Commodity Summaries - Thorium.
^ Information and Issue Briefs - Thorium. World Nuclear Association. Retrieved on 2006-11-01.
^ IAEA: Thorium fuel cycle -- Potential benefits and challenges, pp 45(table 8), 97(ref 78).
^ Phys. Rev. C 73 044326 (April 2006)
^ Phys. Rev. C 52, 113–116 (1995)
^ Crouse, David; Brown, Keith (December 1959). "The Amex Process for Extracting Thorium Ores with Alkyl Amines".Industrial & Engineering Chemistry 51 (12): 1461. Retrieved on 2007-03-09
Los Alamos National Laboratory — Thorium
WebElements.com — Thorium
The Uranium Information Centre provided some of the original material in this article.
European Nuclear Society — Natural Decay Chains

[edit] External links
Wikimedia Commons has media related to:
ThoriumLook up thorium in
Wiktionary, the free dictionary.Thorium information page
New Age Nuclear: article on thorium reactors | Cosmos Magazine
ATSDR ToxFAQs — Thorium
USGS data — Thorium
The Endless Refrigerator/Freezer Deodorizer, a commercial product which claimed to destroy odours 'forever.' Made with thorium-232.
Is thorium the answer to our energy crisis?
Retrieved from "http://en.wikipedia.org/wiki/Thorium"
Categories: Chemical elements | Actinides | Nuclear materials | Carcinogens | Thorium