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Silane

Silane

Properties
General
Name	Silane
Chemical formula	Si H₄
Appearance	Colourless gas
Physical
Formula weight	32.1 u
Melting point	88 K (−185 °C)
Boiling point	161 K (−112 °C)
Density	0.7 g/cm³ (liquid)
Solubility	insoluble
Thermochemistry
Δ_fH⁰_gas	? kJ/mol
Δ_fH⁰_liquid	? kJ/mol
Δ_fH⁰_solid	-1615 kJ/mol
S⁰_{gas, 1 bar}	? J/(mol·K)
S⁰_{liquid, 1 bar}	? J/(mol·K)
S⁰_solid	283 J/(mol·K)
Safety
Ingestion	Relatively low toxicity, but avoid exposure where possible.
Inhalation	Relatively low toxicity: may cause coughing, hyperventilation.
Skin	Irritant, may cause redness and swelling.
Eyes	Similar to skin, may cause irritation. Avoid contact.
More info	[http://ull.chemistry.uakron.edu/erd/chemicals/8/7027.html Hazardous Chemical Database]
SI units were used where possible. Unless otherwise stated, standard conditions were used. Disclaimer and references

Silane is a chemical compound with chemical formula Si H₄. It is the silicon analogue of methane. At room temperature, silane is presumed to be a pyrophoric gas — it spontaneously undergoes combustion in air without the need for external ignition. However, there is a school of thought which says that silane is stable and that it is the natural formation of larger silanes during production which causes its pyrophoricity. Above 420°C, silane decomposes into silicon and hydrogen; it can therefore be used in the chemical vapor deposition of silicon. More generally, a silane is any silicon analogue of an alkane hydrocarbon. Silanes consist of a chain of silicon atoms covalently bound to hydrogen atoms. The general formula of a silane is Si_nH_2n+2. Silanes tend to be less stable than their carbon analogues because the Si-Si bond has a strength slightly lower than the C-C bond. Oxygen decomposes silanes easily, because the silicon-oxygen bond is quite stable. There exists a regular nomenclature for silanes. Each silane's name is the word silane preceded by a numerical prefix (di, tri, tetra, etc.) for the number of silicon atoms in the molecule. Thus Si₂H₆ is disilane, Si₃H₈ is trisilane, and so forth. There is no prefix for one; SiH₄ is simply silane. Silanes can also be named like any other inorganic compound; in this naming system, silane is named silicon tetrahydride. However, with longer silanes, this becomes cumbersome. A cyclosilane is a silane in a ring, just as a cycloalkane is an alkane in a ring. Branched silanes are possible. The radical SiH₃^- is called silyl, Si₂H₅^- is disilanyl etc. If we have trisilane with a silyl group attached to the middle silicon, we have silyltrisilane. It parallels alkanes. Silanes can also take the same functional groups as alkanes, OH to make a silanol. There is (at least in principle) a silicon analogue for all carbon alkanes.

Production

Industrially, silane is produced from metallurgical grade silicon in a two-step process. In the first step, powdered silicon is reacted with hydrochloric acid at about 300 °C to produce trichlorosilane, HSiCl₃, along with hydrogen gas, according to the chemical equation: :Si + 3 HCl → HSiCl₃ + H₂ The trichlorosilane is then boiled on a resinous bed containing a catalyst which promotes its disproportionation to silane and silicon tetrachloride according to the chemical equation: :4 HSiCl₃ → SiH₄ + 3 SiCl₄ The most commonly used catalysts for this process are metal halides, particularly aluminium chloride.

Applications

Several industrial and medical applications exist for silanes. For instance, silanes are used as coupling agents to adhere glass fibers to a polymer matrix, stabilizing the composite material. They can also be used to couple a bio-inert layer on a titanium implant. Other applications include water repellants, masonry protection, control of graffiti, applying polycrystalline silicon layers on silicon wafers when manufacturing semiconductors, and sealants.

External links

- [http://box27.bluehost.com/~edsanvil/wiki/index.php?title=Silane Computational Chemistry Wiki] Category:Silicon compounds Category:Hydrides ja:シラン (化合物)

Chemical formula

A chemical formula (also called molecular formula) is a concise way of expressing information about the atoms that constitute a particular chemical compound. It identifies each type of chemical element by its element symbol and identifies the number of atoms of such element to be found in each discrete molecule of that compound. The number of atoms (if greater than one) is indicated as a subscript. For non-molecular substances the subscripts indicate the ratio of elements in the empirical formula. Chemical formula used for a series of compounds that differ from each other by a constant unit is called general formula. Such a series is called the homologous series, while its members are called homologs.

Elements

In organic chemistry most compounds consist of the following five chemical elements:
- C carbon
- H hydrogen
- N nitrogen
- O oxygen
- S sulfur For other element symbols see list of elements by symbol.

Molecular and structural formulas

For example methane, a simple molecule consisting of one carbon atom bonded to four hydrogen atoms has the chemical formula: : CH₄ and glucose with six carbon atoms, twelve hydrogen atoms and six oxygen atoms has the chemical formula: : C₆H₁₂O₆. A chemical formula may also supply information about the types and spatial arrangement of bonds in the chemical, though it does not necessarily specify the exact isomer. For example ethane consists of two carbon atoms single-bonded to each other, each having three hydrogen atoms bonded to it. Its chemical formula can be rendered as CH₃CH₃. If there were a double bond between the carbon atoms (and thus each carbon only had two hydrogens), the chemical formula may be written: CH₂CH₂, and the fact that there is a double bond between the carbons is assumed. However, a more explicit and correct method is to write H₂C:CH₂ or H₂C=CH₂. The two dots or lines indicate that a double bond connects the atoms on either side of them. A triple bond may be expressed with three dots or lines, and if there may be ambiguity, a single dot or line may be used to indicate a single bond. Molecules with multiple functional groups that are the same may be expressed in the following way: (CH₃)₃CH. However, this implies a different structure from other molecules that can be formed using the same atoms (isomers). The formula (CH₃)₃CH implies a chain of three carbon atoms, with the middle carbon atom bonded to another carbon: Carbon chain and the remaining bonds on the carbons all leading to hydrogen atoms. However, the same number of atoms (10 hydrogens and 4 carbons, or C₄H₁₀) may be used to make a straight chain: CH₃CH₂CH₂CH₃. The alkene 2-butene has two isomers which the chemical formula CH₃CH=CHCH₃ does not identify. The relative position of the two methyl groups must be indicated by additional notation denoting whether the methyl groups are on the same side of the double bond (cis or Z) or on the opposite sides from each other.(trans or E)

Polymers

For polymers, parentheses are placed around the repeating unit. For example, a hydrocarbon molecule that is described as: CH₃(CH₂)₅₀CH₃, is a molecule with 50 repeating units. If the number of repeating units is unknown or variable, the letter n may be used to indicate this: CH₃(CH₂)_nCH₃.

Ions

For ions, the charge on a particular atom may be denoted with a right-hand superscript. For example Na⁺, or Cu²⁺. The total charge on a charged molecule or a polyatomic ion may also be shown in this way. For example: hydronium, H₃O⁺ or sulfate, SO₄^2-.

Isotopes

Although isotopes are more relevant to nuclear chemistry or stable isotope chemistry than to conventional chemistry, different isotopes may be indicated with a left-hand superscript in a chemical formula. For example, the phosphate ion containing radioactive phosphorus-32 is ³²PO₄^3-. Also a study involving stable isotope ratios might include ¹⁸O:¹⁶O. A left-hand subscript is sometimes used to indicate redundantly, for convenience, the atomic number.

Empirical formula

In chemistry, the empirical formula of a chemical is a simple expression of the relative number of each type of atom or ratio of the elements in it. Empirical formulas are the standard for ionic compounds, such as CaCl₂, and for macromolecules, such as SiO₂. An empirical formula makes no reference to isomerism, structure, or absolute number of atoms. The term empirical refers to the process of elemental analysis, a technique of analytical chemistry used to determine the relative percent composition of a pure chemical substance by element. For example, hexane could have a chemical formula of CH₃CH₂CH₂CH₂CH₂CH₃, implying that it has a straight chain structure, 6 carbon atoms, and 14 hydrogen atoms. However the empirical formula for the same molecule would be C₃H₇.

Hydrogen

|- | Critical temperature || 32.19 K |- | Critical pressure || 1.315 MPa |- | Critical density || 30.12 g/L (Bohr radius) Hydrogen (Latin: hydrogenium, from Greek: hydro: water, genes: forming) is a chemical element in the periodic table that has the symbol H and atomic number 1. At standard temperature and pressure it is a colorless, odorless, nonmetallic, univalent, highly flammable diatomic gas. Hydrogen is the lightest and most abundant element in the universe. It is present in water, all organic compounds (rare exceptions exist, like buckminsterfullerene) and in all living organisms. Hydrogen is able to react chemically with most other elements. Stars in their main sequence are overwhelmingly composed of hydrogen in its plasma state. The element is used in ammonia production, as a lifting gas, as an alternative fuel, and more recently as a power source of fuel cells. Despite its ubiquity in the universe, hydrogen is surprisingly hard to produce in large quantities on the Earth. In the laboratory, the element is prepared by the reaction of acids on metals such as zinc. The electrolysis of water is a simple method of producing hydrogen, but is economically inefficient for mass production. Large-scale production is usually achieved by steam reforming natural gas. Scientists are now researching new methods for hydrogen production; if they succeed in developing a cost-efficient method of large-scale production, hydrogen may become a viable alternative to greenhouse-gas-producing fossil fuels. One of the methods under investigation involves use of green algae; another promising method involves the conversion of biomass derivatives such as glucose or sorbitol at low temperatures using a catalyst. Yet another method is the "steaming" of Carbon, whereby hydrocarbons are broken down with heat to release hydrogen.

Basic features

Hydrogen is the lightest chemical element; its most common isotope comprises just one negatively charged electron, distributed around a positively charged proton (the nucleus of the atom). The electron is bound to the proton by the Coulomb force, the electrical force that one stationary, electrically charged nanoparticle exerts on another. The hydrogen atom has special significance in quantum mechanics as a simple physical system for which there is an exact solution to the Schrödinger equation; from that equation, the experimentally observed frequencies and intensities of the hydrogen's spectral lines can be calculated. Spectral lines are dark or bright lines in an otherwise uniform and continuous spectrum, resulting from an excess or deficiency of photons in a narrow frequency range, compared with the nearby frequencies. At standard temperature and pressure, hydrogen forms a diatomic gas, H₂, with a boiling point of only 20.27 K and a melting point of 14.02 K. Under extreme pressures, such as those at the center of gas giants, the molecules lose their identity and the hydrogen becomes a liquid metal. Under the extremely low pressure in space—virtually a vacuum—the element tends to exist as individual atoms, simply because there is no way for them to combine. However, clouds of H₂ and singular hydrogen atoms are said to form in H I and H II regions and are associated with star formation, however the existence of singular hydrogen atoms is disputed.. Hydrogen plays a vital role in powering stars through the proton–proton and carbon–nitrogen cycle. These are nuclear fusion processes, which release huge amounts of energy in stars and other hot celestial bodies as hydrogen atoms combine into helium atoms. H₂ is highly soluble in water, alcohol, and ether. It has a high capacity for adsorption, in which it is attached to and held to the surface of some substances. It is an odorless, tasteless, colorless, and highly flammable gas that burns at concentrations as low as 4% H₂ in air. It reacts violently with chlorine and fluorine, forming hydrohalic acids that can damage the lungs and other tissues. When mixed with oxygen, hydrogen explodes on ignition. A unique property of hydrogen is that its flame is completely invisible in air. This makes it difficult to tell if a leak is burning, and carries the added risk that it is easy to walk into a hydrogen fire inadvertently. See also: hydrogen atom.

Applications

Large quantities of hydrogen are needed in the chemical and petrolium industries, notably in the Haber process for the production of ammonia, which by mass ranks as the world's fifth most highly produced industrial compound. Hydrogen is used in the hydrogenation of fats and oils (into items such as margarine), and in the production of methanol. Hydrogen is used in hydrodealkylation, hydrodesulfurization, and hydrocracking. The element has several other important uses.
- The element is used in the manufacture of hydrochloric acid, in welding processes, and in the reduction of metallic ores.
- It is an ingredient in rocket fuels.
- It is used as the rotor coolant in electrical generators at power stations, because it has the highest thermal conductivity of any gas.
- Liquid hydrogen is used in cryogenic research, including superconductivity studies.
- Since hydrogen is 14.5 times lighter than air, it was once widely used as a lifting agent in balloons and airships. However, this use was curtailed when the Hindenburg disaster convinced the public that the gas was too dangerous for this purpose.
- Deuterium, an isotope of hydrogen (hydrogen-2), is used in nuclear fission applications as a moderator to slow neutrons, and in nuclear fusion reactions. Deuterium compounds have applications in chemistry and biology in studies of reaction isotope effects.
- Tritium (hydrogen-3), produced in nuclear reactors, is used in the production of hydrogen bombs, as an isotopic label in the biosciences, and as a radiation source in luminous paints. There are no "hydrogen wells" or "hydrogen mines" on Earth, so hydrogen cannot be considered a primary energy source like fossil fuels or uranium. Hydrogen can however be burned in internal combustion engines, an approach advocated by BMW's experimental hydrogen car. However, it is currently difficult and dangerous to store and handle in sufficient quantity for motor fuel use. Hydrogen fuel cells are being investigated as mobile power sources with lower emissions than hydrogen-burning internal combustion engines. The low emissions of hydrogen in internal combustion engines and fuel cells are currently offset by the pollution created by hydrogen production. This may change if the substantial amounts of electricity required for water electrolysis can be generated primarily from low pollution sources such as nuclear energy or wind. Research is being conducted on hydrogen as a replacement for fossil fuels. It could become the link between a range of energy sources, carriers and storage. Hydrogen can be converted to and from electricity (solving the electricity storage and transport issues), from bio-fuels, and from and into natural gas and diesel fuel. All of this can theoretically be achieved with zero emissions of CO₂ and toxic pollutants.

History

Hydrogen was first produced by Theophratus Bombastus von Hohenheim (1493–1541)—also known as Paracelsus—by mixing metals with acids. He was unaware that the explosive gas produced by this chemical reaction was hydrogen. In 1671, Robert Boyle described the reaction between two iron fillings and dilute acids, which results in the production of gaseous hydrogen. In 1766, Henry Cavendish was the first to recognize hydrogen as a discrete substance, by identifying the gas from this reaction as "inflammable" and finding that the gas produces water when burned in air. Cavendish stumbled on hydrogen when experimenting with acids and mercury. Although he wrongly assumed that hydrogen was a compound of mercury—and not of the acid—he was still able to accurately describe several key properties of hydrogen. Antoine Lavoisier gave the element its name and proved that water is composed of hydrogen and oxygen. One of the first uses of the element was for balloons. The hydrogen was obtained by mixing sulfuric acid and iron. Harold C. Urey discovered Deuterium, an isotope of hydrogen, by repeated distilling the same sample of water. For this discovery, Urey received the Nobel prize for in 1934. In the same year, the third isotope, tritium, was discovered. Because of its relatively simple structure, hydrogen has often been used in models of how an atom works.

Electron energy levels

The ground state energy level of the electron in a Hydrogen atom is 13.6 eV, which is equivalent to an ultraviolet photon of roughly 92 nm. With the Bohr Model the energy levels of Hydrogen can be calculated fairly accurately. This is done by modeling the electron as revolving around the proton, much like the earth revolving around the sun. Except the sun holds earth in orbit with the force of gravity, but the proton holds the electron in orbit with the force of electromagnetism. Another difference between the Earth-Sun system and the Electron-Proton system is that, in this model, due to quantum mechanics the electron is allowed to only be at very specific distances from the proton. Modeling the hydrogen atom in this fashion yields the correct energy levels and spectrum.

Occurrence

quantum mechanics.]] Hydrogen is the most abundant element in the universe, making up 75% of normal matter by mass and over 90% by number of atoms. This element is found in great abundance in stars and gas giant planets. It is very rare in the Earth's atmosphere (1 ppm by volume), because being the lightest gas causes it to escape Earth's gravity, though when compounds are considered, it is the tenth most abundant element on Earth. The most common source for this element on Earth is water, which is composed two parts hydrogen to one part oxygen (H₂O). Other sources include most forms of organic matter (currently all known life forms) including coal, natural gas, and other fossil fuels. Methane (CH₄) is an increasingly important source of hydrogen. Throughout the Universe, hydrogen is mostly found in the plasma state whose properties are quite different to molecular hydrogen. As a plasma, hydrogen's electron and proton are not bound together, resulting in very high electrical conductivity, even when the gas is only partially ionised. The charged particles are highly influenced by magnetic and electric fields, for example, in the Solar Wind they interact with the Earth's magnetosphere giving rise to Birkeland currents and the aurora. Hydrogen can be prepared in several different ways: steam on heated carbon, hydrocarbon decomposition with heat, reaction of a strong base in an aqueous solution with aluminium, water electrolysis, or displacement from acids with certain metals. Commercial bulk hydrogen is usually produced by the steam reforming of natural gas. At high temperatures (700–1100 °C), steam reacts with methane to yield carbon monoxide and hydrogen. :CH₄ + H₂O → CO + 3 H₂ Additional hydrogen can be recovered from the carbon monoxide through the water-gas shift reaction: :CO + H₂O → CO₂ + H₂

Compounds

The lightest of all gases, hydrogen combines with most other elements to form compounds. Hydrogen has an electronegativity of 2.2, so it forms compounds where it is the more nonmetallic and where it is the more metallic element. The former are called hydrides, where hydrogen either exists as H^- ions or just as a solute within the other element (as in palladium hydride). The latter tend to be covalent, since the H⁺ ion would be a bare nucleus and so has a strong tendency to pull electrons to itself. These both form acids. Thus even in an acidic solution one sees ions like hydronium (H₃O⁺) as the protons latch on to something. Although exotic on earth, one of the most common ions in the universe is the H₃⁺ ion. Hydrogen combines with oxygen to form water, H₂O, and releases a lot of energy in doing so, burning explosively in air. Deuterium oxide, or D₂O, is commonly referred to as heavy water. Hydrogen also forms a vast array of compounds with carbon. Because of their association with living things, these compounds are called organic compounds, and the study of the properties of these compounds is called organic chemistry. organic chemistry

Forms

Under normal conditions, hydrogen gas is a mix of two different kinds of molecules which differ from one another by the relative spin of the nuclei. These two forms are known as ortho- and para-hydrogen (this is different from isotopes, see below). In ortho-hydrogen the nuclear spins are parallel (form a triplet), while in para they are antiparallel (form a singlet). At standard conditions hydrogen is composed of about 25% of the para form and 75% of the ortho form (the so-called "normal" form). The equilibrium ratio of these two forms depends on temperature, but since the ortho form has higher energy (is an excited state), it cannot be stable in its pure form. In low temperatures (around boiling point), the equilibrium state is comprised almost entirely of the para form. The conversion process between the forms is slow, and if hydrogen is cooled down and condensed rapidly, it contains large quantities of the ortho form. It is important in preparation and storage of liquid hydrogen, since the ortho-para conversion produces more heat than the heat of its evaporation, and a lot of hydrogen can be lost by evaporation in this way during several days after liquefying. Therefore, some catalysts of the ortho-para conversion process are used during hydrogen cooling. The two forms have also slightly different physical properties. For example, the melting and boiling points of parahydrogen are about 0.1 K lower than of the "normal" form.

Isotopes

Hydrogen is the only element that has different names for its isotopes. (During the early study of radioactivity, various heavy radioactive isotopes were given names, but such names are no longer used, although one element, radon, has a name that originally applied to only one of its isotopes.) The symbols D and T (instead of ²H and ³H) are sometimes used for deuterium and tritium, although this is not officially sanctioned. (The symbol P is already in use for phosphorus and is not available for protium.) ;¹H The most common isotope of hydrogen, this stable isotope has a nucleus consisting of a single proton; hence the descriptive, although rarely used, name protium. The spin of a protium atom is 1/2+. ;²H The other stable isotope is deuterium, with an extra neutron in the nucleus. Deuterium comprises 0.0184%–0.0082% of all hydrogen (IUPAC); ratios of deuterium to protium are reported relative to the VSMOW standard reference water. The spin of a deuterium atom is 1+. ;³H The third naturally occurring hydrogen isotope is the radioactive tritium. The tritium nucleus contains two neutrons in addition to the proton. It decays through beta decay and has a half-life of 12.32 years. Tritium occurs naturally due to cosmic rays interacting with atmospheric gases. Like ordinary hydrogen, tritium reacts with the oxygen in the atmosphere to form T₂O. This radioactive "water" molecule constantly enters the Earth's seas and lakes in the form of slightly radioactive rain, but its half-life is short enough to prevent a buildup of hazardous radioactivity. The spin of a tritium atom is 1/2+. ;⁴H Hydrogen-4 was synthesized by bombarding tritium with fast-moving deuterium nuclei. It decays through neutron emission and has a half-life of 9.93696x10^-23 seconds. The spin of a hydrogen-4 atom is 2-. ;⁵H In 2001 scientists detected hydrogen-5 by bombarding a hydrogen target with heavy ions. It decays through neutron emission and has a half-life of 8.01930x10^-23 seconds. ;⁶H Hydrogen-6 decays through triple neutron emission and has a half-life of 3.26500^-22 seconds. ;⁷H In 2003 hydrogen-7 was created ([http://physicsweb.org/articles/news/7/3/3 article]) at the RIKEN laboratory in Japan by colliding a high-energy beam of helium-8 atoms with a cryogenic hydrogen target and detecting tritons—the nuclei of tritium atoms—and neutrons from the breakup of hydrogen-7, the same method used to produce and detect hydrogen-5.

References

# # # # # #
- [http://www.riken.go.jp/engn/r-world/research/lab/wako/ribeam/ RIKEN Beam Science Laboratory, Japan - Heavy hydrogen research]
- [http://chartofthenuclides.com/default.html Nuclides and Isotopes] Fourteenth Edition: Chart of the Nuclides, General Electric Company, 1989 ;Book references:
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External links

- [http://www.hydropole.ch/Hydropole/Intro/Phasediag.gif Hydrogen phase diagram.]
- [http://www.compchemwiki.org/index.php?title=Hydrogen Computational Chemistry Wiki] Category:Nonmetals Category:Fuels Category:Chemical elements ko:수소 ms:Hidrogen ja:水素 simple:Hydrogen th:ไฮโดรเจน

Atomic weight

The atomic mass of a chemical element (also known as the relative atomic mass or average atomic mass or atomic weight) is the average atomic mass of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance. Periodic tables usually list these with reference to the local environment of Earth's crust and atmosphere. For artificial elements the nucleon count of the most stable isotope is listed in parentheses as the atomic mass. The atomic mass of an isotope 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. 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, 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. A similar definition applies to molecules; it is then called molecular mass. One can compute the molecular mass of a compound by adding the atomic masses of its constituent atoms multiplied by the ratios of elements given in the chemical formula. A similar formula mass can be calculated for those compounds which do not form molecules. Direct comparison and measurement of the masses of atoms and molecules is achieved with mass spectrometry. One mole of a substance always contains almost exactly the atomic or molecular mass of that substance, expressed in grams. For example, the atomic mass of iron is 55.847, and therefore one mole of iron has a mass of 55.847 grams.

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 term standard atomic weight refers to the mean relative atomic mass of an element.

External links

- [http://physics.nist.gov/cgi-bin/Compositions/stand_alone.pl?ele=&ascii=html&isotype=some Atomic masses of all isotopes] Category:Chemical properties Category:Mass ko:원자 질량 ja:原子量 th:มวลอะตอม

Atomic mass unit

AMU redirects here, but may also refer to the Arab Maghreb Union The unified atomic mass unit (u), or dalton (Da), is a small unit of mass used to express atomic masses and molecular masses. It is defined to be 1/12 of the mass of one atom of carbon-12. :1 u = 1/N_A gram = 1/(1000 N_A) kg (where N_A is Avogadro's number) :1 u ≈ 1.66053886 x 10^-27 kg See 1 E-27 kg for a list of objects which have a mass of about 1 u. The symbol amu for atomic mass unit can sometimes still be found, particularly in older works. Atomic masses are often written without any unit and then the atomic mass unit is implied. In biochemistry and molecular biology literature (particularly in reference to proteins), the term "dalton" is used, with the symbol "Da". Because proteins are large molecules, they are typically referred to in kilodaltons, or "kDa", with one kilodalton being equal to 1000 daltons. The unified atomic mass unit is not an SI unit of mass, although it is (only by that name, and only with the symbol u) accepted for use with SI. See SI website link below. The unit is convenient because one hydrogen atom has a mass of approximately 1 u, and more generally an atom or molecule that contains n protons and neutrons will have a mass approximately equal to n u. (The reason is that a carbon-12 atom contains 6 protons, 6 neutrons and 6 electrons, with the protons and neutrons having about the same mass and the electron mass being negligible in comparison.) This is only a rough approximation however, since it does not account for the mass contained in the binding energy of an atom's nucleus; this binding energy mass is not a fixed fraction of an atom's total mass. Another reason the unit is used is that it is experimentally much easier and more precise to compare masses of atoms and molecules (determine relative masses) than to measure their absolute masses. Masses are compared with a mass spectrometer (see below). Avogadro's number (N_A) and the mole are defined so that one mole of a substance with atomic or molecular mass 1 u will have a mass of precisely 1 gram. For example, the molecular mass of water is 18.01508 u, and this means that one mole of water has a mass of 18.01508 grams, or conversely that 1 gram of water contains N_A/18.01508 ≈ 3.3428 × 10²² molecules.

Measuring relative atomic masses

The relative atomic mass is measured with a mass spectrometer. After placing a sample of the element to be measured in the mass spectrometer it is bombarded with electrons which turns the atoms into positive ions. An electric field is then used to accelerate these positive ions, after which the ions are deflected using a magnetic field. As a result the various isotopes are separated out due to the ions of lighter isotopes being deflected more than those heavier. This produces a mass spectrum. This spectrum provides two things: # Relative isotopic masses in the sample # Abundances of the isotopes

History

The chemist John Dalton was the first to suggest the mass of one atom of hydrogen as the atomic mass unit. Francis Aston, inventor of the mass spectrometer, later used 1/16 of the mass of one atom of oxygen-16 as his unit. Before 1961, the physical atomic mass unit was defined as 1/16 of the mass of one atom of oxygen-16, while the chemical atomic mass unit was defined as 1/16 of the average mass of an oxygen atom (taking the natural abundance of the different oxygen isotopes into account). Both units are slightly smaller than the unified atomic mass unit, which was adopted by the International Union of Pure and Applied Physics in 1960 and by the International Union of Pure and Applied Chemistry in 1961.

External links

- [http://www1.bipm.org/en/si/si_brochure/chapter4/table7.html SI website on acceptable non-SI units]
- [http://physics.nist.gov/cgi-bin/cuu/Value?ukg Accepted value of 1 u as of 2002] Category:Nuclear chemistry Category:Units of mass ja:原子質量単位 ko:아보가드로 수 th:หน่วยมวลอะตอม

Melting point

The melting point of a solid is the temperature at which it changes state from solid to liquid. When considered as the temperature of the reverse change from liquid to solid, it is referred to as the freezing point. For most substances, melting and freezing points are equal. For example, the melting point and freezing point of the element mercury is 234.32 kelvins (−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 32 to 40 °C (89.6 to 104 °F); this phenomenon is known as hysteresis. Certain materials, such as glass, may harden without crystalizing; this is called an amorphous solid. Unlike the boiling point, the melting point is relatively insensitive to pressure. The material with the highest known melting point at atmospheric pressure is graphite, with a melting point of 3,948 kelvins (3,674.8 °C or 6,646.5 °F). Water's Melting/Freezing point is 0 C, or 32 F. Melting point is often used to ascertain purity of and characterise organic compounds. The melting point of a pure substance is always higher than the melting point of an impure sample of that particular substance. When two chemical substances are mixed, the melting point of the resultant mixture will be lower than the melting point of either constituent. The mixing ratio that results in the lowest possible melting point is known as the eutectic point.

Kelvin

The kelvin (symbol: K) is the SI unit of temperature, and is one of the seven SI base units. It is defined as the fraction 1/273.16 of the thermodynamic temperature of the triple point of water. A temperature given in kelvins, without further qualification, is measured with respect to absolute zero, where molecular motion stops. It is also common to give a temperature relative to the reference temperature of 273.15 K, approximately the melting point of water under ordinary conditions; this convention is the Celsius temperature scale. The kelvin is named after the British physicist and engineer William Thomson, 1st Baron Kelvin; his barony was in turn named after the River Kelvin, which runs through the grounds of the University of Glasgow.

SI multiples

Typographical conventions

The word kelvin as an SI unit is correctly written with a lowercase k (unless at the beginning of a sentence), and is never preceded by the words degree or degrees, or the symbol °, unlike degrees Fahrenheit, or degrees Celsius. This is because the latter are adjectives, whereas kelvin is a noun. It takes the normal plural form by adding an s in English: kelvins. When the kelvin was introduced in 1954 (10th General Conference on Weights and Measures (CGPM), Resolution 3, CR 79), it was the "degree Kelvin", and written °K; the "degree" was dropped in 1967 (13th CGPM, Resolution 3, CR 104). Note that the symbol for the kelvin unit is always a capital K and never italicised. There is a space between the number and the K, as with all other SI units. Unicode includes the "kelvin sign" at U+212A (in your browser it looks like K). However, the "kelvin sign" is canonically decomposed into U+004B, thereby seen as a (preexisting) encoding mistake, and it is better to use U+004B (K) directly.

Conversion factors

Kelvins and Celsius

The Celsius temperature scale is now defined in terms of the kelvin, with 0 °C corresponding to 273.15 kelvins.
- kelvins to degrees Celsius
- :

\mathrm = \mathrm - 273.15

Temperature and energy

In a thermodynamic system, the energy of the particles of a perfect gas is proportional to the absolute temperature, where the constant of proportionality is the Boltzmann constant. As a result, it is possible to determine the average kinetic energy

\overline

of the gas particles at the temperature T or to calculate the temperature of the gas from the average kinetic energy of the particles: :

\overline = \frac \cdot k_B \cdot \mathrm

External link

- [http://www1.bipm.org/en/si/si_brochure/chapter2/2-1/2-1-1/kelvin.html BIPM brochure on the kelvin] Category:SI base units Category:Units of temperature ko:켈빈 ja:ケルビン simple:Kelvin th:เคลวิน

Boiling point

:Alternate use: Boiling Point, a film by Takeshi Kitano; Boiling Points, a television series The boiling point of a substance is the temperature at which it can change its state from a liquid to a gas throughout the bulk of the liquid. A 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 definition of boiling point is that it is the temperature at which the vapor pressure of the liquid equals the pressure of the environment. Something that should be remembered is that boiling is evidenced by the appearance of bubbles containing vapor from the liquid. Production of this vapor requires energy and thus does not occur without some source of energy. This source can be a hot surface or even the liquid itself. Hot liquid will boil as it rises through the bulk liquid when the pressure of the environment drops to the vapor pressure of the liquid at its temperature. This production of vapor will quickly stop because the temperature of the liquid will be reduced by the vaporization thus reducing the vapor pressure. 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. 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. Likewise, the boiling point decreases with decreasing ambient pressure until the triple point is reached. The boiling point cannot be reduced below the triple point. The process of changing from a liquid to a gas requires an amount of heat called the latent heat of vaporization. As heat is added to a liquid at its boiling point, all of this heat goes toward the phase change from liquid to gas, thus the temperature of the substance remains constant even though heat has been added. The word latent, which comes from Latin and means hidden, is used to describe this "disappearing" heat that is added, but doesn't result in an increase in temperature. Since heat is added with no corresponding change in temperature, the heat capacity of the liquid is essentially infinite at the boiling point.
Intermolecular interactions
In terms of intermolecular interactions, the boiling point represents the point at which the liquid molecules possess enough heat 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. 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. (Cf. DeVoe, Howard, Thermodynamics and Chemistry. Prentice-Hall, 2001)
See also

- Leidenfrost effect
- flash point
- boiling delay
- critical temperature
- triple point
- boiling-point elevation Category:Chemical properties Category:Thermodynamics Category:Fluid dynamics ko:끓는점 ja:沸点 th:จุดเดือด

Kilogram per cubic metre

Kilogram per cubic metre is the SI measure of density and is represented as kg/m³, where kg stands for kilogram and m³ stands for cubic metre. The density of water is about 1000 kg/m³ (and is exactly this at 277 K), since a cubic metre of water weighs about a tonne. kg/m³ is sometimes equivalently written as kg m^-3. To convert from g/cm³ to kg/m³, divide by 1000; multiply by 1000 for the opposite conversion. A gram per litre is identical in value to kg/m³ .

External link

- [http://www.ex.ac.uk/trol/scol/index.htm Conversion Calculator for Units of DENSITY] Category:Units of density Category:SI derived units ko:킬로그램 매 세제곱미터 ja:キログラム毎立方メートル

Solubility

A substance is soluble in a fluid if it dissolves in that fluid. The dissolved substance is called the solute and the dissolving fluid (usually present in excess) is called the solvent, which together form a solution. The process of dissolving is called solvation, or hydration if the solvent is water. A solution at equilibrium cannot hold any more solute and is said to be saturated. Solutions may, under special conditions, hold more solute than the solvent can normally dissolve. This is called supersaturation. The maximum equilibrium amount of solute which can normally dissolve per amount of solvent (or solution) is the solubility of that solute in that solvent. It is often expressed as a maximum concentration of a saturated solution. The solubility of one substance dissolving in another is determined by the intermolecular forces between the solvent and solute, temperature, the entropy change that accompanies the solvation, the presence and amount of other substances, and sometimes pressure or partial pressure of a solute gas. For salts, solubility in aqueous solutions is often dependent on a solubility constant. The solubility constant is a special case of an equilibrium constant for the reaction of dissolving the salt in question, with the concentration of undissolved compound not in the expression because it is not in the aqueous phase. The solubility constant is also "applicable" (i. e. useful) to precipitation, the reverse of the dissolving reaction. As with other equilibrium constants, temperature can affect the numerical value of solubility constant. Solvents are normally characterized as polar or nonpolar. Polar solvents will dissolve ionic compounds and covalent compounds which ionize, while nonpolar solvents will dissolve nonpolar covalent compounds. For example, ordinary table salt, an ionic compound, will dissolve in water, but not in ethanol. Common solvents used in organic chemistry include acetone, ethanol, water, and benzene. Water and nonpolar solvents are immiscible; they do not form homogeneous mixtures but separate into two distinct phases or form milky emulsions. While solutions are typically thought of as solids being mixed into liquids, any two states of matter can be mixed and be called a solution. Carbonated water is a solution of a gas in a liquid, hydrogen (a gas) can dissolve in palladium (a solid), and stainless steel is a solution of a solid in a solid (called an alloy).

Solubility of bonding type in water

Bonding type	Solubility in water	Example
ionic	most soluble	See below
metallic	insoluble	Fe
metallic	unless they react with water	K
polar covalent	soluble if it H bonds	glucose
	soluble by reaction	HCl
	insoluble otherwise	ether
non-polar covalent	most insoluble	benzene
non-polar covalent	some slightly soluble	O₂
covalent lattice	insoluble	diamond

Solubility of ionic compounds

Soluble	Insoluble
Group 1 and NH₄⁺ compounds	carbonates (except Group 1 and NH₄⁺ compounds)
nitrates	sulfites (except Group 1 and NH₄⁺ compounds)
acetates (ethanoates)	phosphates (except Group 1 and NH₄⁺ compounds)
chlorides, bromides and iodides (except Ag⁺, Pb²⁺, Cu⁺ and Hg₂²⁺)	hydroxides and oxides (except Group 1, NH₄⁺, Ba²⁺, Sr²⁺ and Ca²⁺)
sulfates (except Ag⁺, Pb²⁺, Ba²⁺, Sr²⁺ and Ca²⁺)	sulfides (except Group 1, Group 2 and NH₄⁺ compounds)

Software tools for prediction of solution

One of the most recent and prominent solution (solubility) prediction technologies is applied in [http://www.q-pharm.com/home/contents/drug_d/soft/ Quantum 3.1] that is a suite of Molecular Modeling software for Linux and Windows. The software calculates the solvation energy and solubility for a molecule or a library of molecules in a number of solvents (e.g. water and DMSO). The Quantum 3.1 [http://www.q-pharm.com/home developer] is also a service provider.

Joule

The joule (symbol: J) is the SI unit of energy, or work. It is named in honour of the physicist James Prescott Joule (1818–1889).

Definition

The joule is a derived unit defined as the work done, or energy required, to exert a force of one newton for a distance of one metre, so the same quantity may be referred to as a newton metre or newton-metre (also with meter spelling), with the symbol N·m or N m. It can also be written as kg·m²·s⁻². However, the newton metre is usually used as a measure of torque, not energy. One joule is also:
- The work required to move an electric charge of one coulomb through an electrical potential difference of one volt; or one coulomb volt, with the symbol C·V.
- The work done to produce power of one watt continuously for one second; or one watt second (compare kilowatt-hour), with the symbol W·s

Conversions

1 joule is exactly 10⁷ erg. 1 joule is approximately equal to:
- 6.241506363 eV (electron-volts)
- 0.239 cal (calorie) (small calories)
- 2.390 Calorie or kilocalorie (food)
- 9.48 BTU (British thermal unit)
- 0.738 ft·lbf (foot pound force)
- 23.7 ft·pdl (foot poundals)
- 2.7778 kilowatt-hour
- 2.7778 watt-hour
- 9.8692 litre-atmosphere
- the energy required to lift a small apple (102 g) one metre against Earth's gravity Units defined in terms of the joule include:
- 1 thermochemical calorie = 4.184 J (exact)
- 1 International Table calorie = 4.1868 J (exact)
- 1 watt-hour = 3600 J (exact)

Mole unit

The mole (symbol: mol) is the SI term identifying the number of particles in a given amount of matter. It is a dimensionless quantity (meaning a number without units) numerically equal to Avogadro's number.

Definition

The formal definition of the mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 12 grams of carbon 12, where the carbon 12 atoms are unbound, at rest and in their ground state. The number of atoms in 0.012 kilogram of carbon 12 is known as Avogadro's number. It is approximately 6.0221415 (2002 CODATA value). A mole is a dimensionless name for an integer, much like dozen or googol. Although the exact value of the mole is not known at present, it is equal to Avogadro's number, which is known to 1 part in 10 million. Because of the relationship of the atomic mass unit to Avogadro's number, a practical way of stating this for atoms or molecules is: That amount of the substance containing exactly the same number of grams as the number of the atomic weight of the substance. Since iron, for example, has an atomic weight of 55.845, there are 55.845 grams in a mole of iron.

Elementary entities

When the mole is used to specify the amount of a substance, the kind of elementary entities (particles) in the substance must be identified. The particles can be atoms, molecules, ions, formula units, electrons, or other particles. For example, one mole of water is equivalent to about 18 grams of water and contains one mole of H₂O molecules, but three moles of atoms (two moles H and one mole O). When the substance of interest is a gas, the particles are usually molecules. However, the noble gases (He, Ar, Ne, Kr, Xe, Rn) are all monoatomic, that is each particle of gas is a single atom. All gases have the same molar volume of 22.4 litres per mole at STP (see Avogadro's Law). A mole of atoms or molecules is also called a "gram atom" or "gram molecule".

History

The name mole is attributed to Wilhelm Ostwald who introduced the concept in the year 1902. He used it to express the gram molecular weight of a substance. So, for example, 1 mole of hydrochloric acid (HCl) has a mass of 36.5 grams (atomic weights Cl: 35.5 u, H: 1.0 u). Prior to 1959 both the IUPAP and IUPAC used oxygen to define the mole, the chemists defining the mole as the number of atoms of oxygen which had mass 16 g, the physicists using a similar definition but with the oxygen-16 isotope only. The two organizations agreed in 1959/1960 to define the mole as such: The mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon 12; its symbol is "mol." This was adopted by the CIPM (International Committee for Weights and Measures) in 1967, and in 1971 it was adopted by the 14th CGPM (General Conference on Weights and Measures) In 1980 the CIPM clarified the above definition, defining that the carbon-12 atoms are unbound and in their ground state.

Utility of moles

The mole is useful in chemistry because it allows different substances to be measured in a comparable way. Using the same number of moles of two substances, both amounts have the same number of molecules or atoms. The mole makes it easier to interpret chemical equations in practical terms. Thus the equation: :2H₂ + O₂ = 2H₂O can be understood as "two moles of hydrogen plus one mole of oxygen yields two moles of water." Moles are useful in chemical calculations, because they enable the calculation of yields and other values when dealing with particles of different mass. Number of particles is a more useful unit in chemistry than mass or weight, because reactions take place between atoms (for example, two hydrogen atoms and one oxygen atom make one molecule of water) that have very different weights (one oxygen atom weighs almost 16 times as much as a hydrogen atom). However, the raw numbers of atoms in a reaction are not convenient, because they are very large; for example, just one mL of water contains over 3 × 10²² (or 30,000,000,000,000,000,000,000) molecules.

Example calculation

In this example, moles are used to calculate the mass of CO₂ given off when 1 g of ethane is burnt. The equation for this chemical reaction is: :7 O₂ + 2 C₂H₆ → 4 CO₂ + 6 H₂O Here, 7 moles of oxygen react with 2 moles of ethane to give 4 moles of carbon dioxide and 6 moles of water. Notice that the number of moles does not need to balance on either side of the equation. This is because a mole does not count mass or the number of atoms involved, simply the number of individual particles. In our calculation it is first necessary to work out the number of moles of ethane that has been burnt. The mass in grams of one mole of a substance is by definition its atomic or molecular mass. The atomic mass of hydrogen is 1, and the atomic mass of carbon is 12, so the molecular mass of C₂H₆ is (2 × 12) + (6 × 1) = 30. One mole of ethane is 30 g. The amount burnt was 1 g, or 1/30th of a mole. The molecular mass of CO₂ (the atomic mass of carbon is 12 and that of oxygen is 16) is 2 × 16 + 12 = 44, so one mole of carbon dioxide is 44 g. From the formula we know that :1 mole of ethane gives off 2 moles of carbon dioxide (because 2 give off 4). We also know the masses of a mole of both ethane and carbon dioxide, so :30 g of ethane gives off 2 × 44 g of carbon dioxide. It is necessary to multiply the mass of carbon dioxide by 2 because two moles are produced. However, we also know that just 1/30th of a mole of ethane was burnt. Again: :1/30th of a mole of ethane gives off 2 × 1/30th of a mole of carbon dioxide, so finally: :30 × 1/30 g ethane gives off 44 × 2/30 g of carbon dioxide = 2.93 g.

References

# [http://www.bipm.org/en/si/base_units/ Official SI Unit definitions] Category:SI base units Category:Units of amount of substance ko:몰 ja:モル

Standard enthalpy change of formation

The standard enthalpy of formation or "standard heat of formation" of a compound is the change of enthalpy that accompanies the formation of 1 mole of a substance in its standard state from its constituent elements in their standard states (the most stable form of the element at 1 atmosphere of pressure and the specified temperature, usually 25 degrees Celsius). Its symbol is ΔH_f^O. The standard enthalpy change of formation is measured in units of energy per amount of substance. Most are defined in kilojoules per mole or kJ/mol but can also be measured in calories per mole, joules per mole or kilocalories per gram (any combination of these units conforming to the energy per mass or amount guideline) All elements in their standard states oxygen gas, graphite, etc.) have a standard enthalpy of formation of zero, as there is no change involved. The standard enthalpy change of formation is used in thermochemistry to find the standard enthalpy change of reaction. This is done by subtacting the standard enthalpy of formation of the reactants from the standard enthalpy change of the products, as shown in the equation below. ΔH_reaction^O = ΔH_f^O (Products) - ΔH_f^O (Reactants) The standard enthalpy of formation is equivalent to the sum of many separate processes included in the Born-Haber cycle of synthesis reactions. For example, to calculate the standard enthalpy of formation of sodium chloride, we use the following reaction: :Na_(s) + (1/2)Cl_2(g) → NaCl_(s) This process is made of many separate sub-processes, each with their own enthalpies. Therefore, we must take into account: #The standard enthalpy of atomization of solid sodium #The first ionization energy of gaseous sodium #The standard enthalpy of atomization of chlorine gas #The electron affinity of chlorine atoms #The lattice enthalpy of sodium chloride The sum of all these values will give the standard enthalpy of formation of sodium chloride. Additionally, applying Hess's Law shows that the sum of the individual reactions corresponding to the enthalpy change of formation for each substance in the reaction is equal to the enthalpy change of the overall reaction, regardless of the number of steps or intermediate reactions involved. In the example above the standard enthalpy change of formation for sodium chloride is equal to the sum of the standard enthalpy change of formation for each of the steps involved in the process. This is especially useful for very long reactions with many intermediate steps and compounds. This leads us to the point that at times chemists may define a standard enthalpy of formation for a reaction that can be (and often is) hypothetical. For instance we cannot combine carbon and hydrogen in the laboratory to make methane, yet we define the standard enthalpy of formation for methane as -74.8 kJ/mol.

Mole unit

Definition

Elementary entities

History

Utility of moles

Example calculation

References

# [http://www.bipm.org/en/si/base_units/ Official SI Unit definitions] Category:SI base units Category:Units of amount of substance ko:몰 ja:モル

Standard enthalpy change of formation

Standard molar entropy

In chemistry, the standard molar entropy is the entropy content of one mole of substance, under conditions of standard temperature and pressure. By comparing the entropies S and enthalpies H of products and reactants in a chemical reaction, we can determine whether the reaction will go forward or backwards. If the Gibbs free energy (G = H - TS) is negative, the forward reaction will take place. The standard molar entropy is usually given the symbol S⁰, and the units J/(mol·K) (joules per mole-kelvin). Category:Chemical properties Category:Entropy

Standard molar entropy

SI

The International System of Units (abbreviated SI from the French language name Système International d'Unités) is the modern form of the metric system. It is the world's most widely used system of units, both in everyday commerce and in science. The older metric system included several groupings of units. The SI was developed in 1960 from one of these, the metre-kilogram-second (MKS) system, rather than the centimetre-gram-second (CGS) system, which, in turn, had many variants. The SI introduced several newly named units. The SI is not static; it is a living set of standards where units are created and definitions are modified with international agreement as measurement technology progresses. With few exceptions (such as draught beer sales in the United Kingdom), the system is legally being used in every country in the world, and many countries do not maintain official definitions of other units. In the United States, industrial use of SI is increasing, but popular use is still limited. In the United Kingdom, conversion to metric units is official policy but not yet complete. Those countries that still recognize non-SI units (e.g. the US and UK) have redefined most of their traditional, non-SI units in terms of SI units.

History

:See main articles: metre, kilogram, second, ampere, Kelvin, and candela. The metric system was officially adopted in France after the French Revolution. During the history of the metric system a number of variations have evolved and their use spread around the world replacing many traditional measurement systems. By the end of World War II a number of different systems of measurement were still in use throughout the world. Some of these systems were metric system variations whilst others were based on the Imperial and American systems. It was recognised that additional steps were needed to promote a worldwide measurement system. As a result the 9th General Conference on Weights and Measures (CGPM), in 1948, asked the International Committee for Weights and Measures (CIPM) to conduct an international study of the measurement needs of the scientific, technical, and educational communities. Based on the findings of this study, the 10th CGPM in 1954 decided that an international system should be derived from six base units to provide for the measurement of temperature and optical radiation in addition to mechanical and electromagnetic quantities. The six base units recommended were the