Hydrogen

Definition

Hydrogen,  ₁H

Purple glow in its plasma state
General properties
Appearance	colorless gas
Standard atomic weight (Ar, standard)	[1.00784, 1.00811] conventional: 1.008
Hydrogen in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson – ↑ H ↓ Li – ← hydrogen → helium
Atomic number (Z)	1
Group	group 1
Period	period 1
Element category	reactive nonmetal
Block	s-block
Electron configuration	1s
Electrons per shell	1
Physical properties
Phase at STP	gas
Melting point	13.99 K (−259.16 °C, −434.49 °F)
Boiling point	20.271 K (−252.879 °C, −423.182 °F)
Density (at STP)	0.08988 g/L
when liquid (at m.p.)	0.07 g/cm(solid: 0.0763 g/cm)
when liquid (at b.p.)	0.07099 g/cm
Triple point	13.8033 K, 7.041 kPa
Critical point	32.938 K, 1.2858 MPa
Heat of fusion	(H2) 0.117 kJ/mol
Heat of vaporization	(H2) 0.904 kJ/mol
Molar heat capacity	(H2) 28.836 J/(mol•K)
Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 15 20
Atomic properties
Oxidation states	−1, +1 (an amphotericoxide)
Electronegativity	Pauling scale: 2.20
Ionization energies	1st: 1312.0 kJ/mol
Covalent radius	31±5 pm
Van der Waals radius	120 pm
Spectral lines
Miscellanea
Crystal structure	hexagonal
Speed of sound	1310 m/s (gas, 27 °C)
Thermal conductivity	0.1805 W/(m•K)
Magnetic ordering	diamagnetic
Magnetic susceptibility	−3.98•10 cm/mol (298 K)
CAS Number	12385-13-6 1333-74-0 (H2)
History
Discovery	Henry Cavendish (1766)
Named by	Antoine Lavoisier (1783)
Main isotopes of hydrogen
Iso­tope Abun­dance Half-life(t1/2) Decay mode Pro­duct H 99.98% stable H 0.02% stable H trace 12.32 y β He

Hydrogen is a chemical element with symbol H and atomic number 1. With a standard atomic weight of 1.008, hydrogen is the lightest element on the periodic table. Its monatomic form (H) is the most abundant chemical substance in the Universe, constituting roughly 75% of all baryonic mass. Non-remnant stars are mainly composed of hydrogen in the plasma state. The most common isotope of hydrogen, termed protium (name rarely used, symbol H), has one proton and no neutrons.

The universal emergence of atomic hydrogen first occurred during the recombination epoch. At standard temperature and pressure, hydrogen is a colorless, odorless, tasteless, non-toxic, nonmetallic, highly combustible diatomic gas with the molecular formula H₂. Since hydrogen readily forms covalent compounds with most nonmetallic elements, most of the hydrogen on Earth exists in molecular forms such as water or organic compounds. Hydrogen plays a particularly important role in acid–base reactions because most acid-base reactions involve the exchange of protons between soluble molecules. In ionic compounds, hydrogen can take the form of a negative charge (i.e., anion) when it is known as a hydride, or as a positively charged (i.e., cation) species denoted by the symbol H. The hydrogen cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds are always more complex. As the only neutral atom for which the Schrödinger equation can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics.

Hydrogen gas was first artificially produced in the early 16th century by the reaction of acids on metals. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, and that it produces water when burned, the property for which it was later named: in Greek, hydrogen means "water-former".

Industrial production is mainly from steam reforming natural gas, and less often from more energy-intensive methods such as the electrolysis of water. Most hydrogen is used near the site of its production, the two largest uses being fossil fuel processing (e.g., hydrocracking) and ammonia production, mostly for the fertilizer market. Hydrogen is a concern in metallurgy as it can embrittle many metals, complicating the design of pipelines and storage tanks.

Properties

Combustion

The Space Shuttle Main Engine burnt hydrogen with oxygen, producing a nearly invisible flame at full thrust.

Explosion of a hydrogen–air mixture.

Hydrogen gas (dihydrogen or molecular hydrogen, also called diprotium when consisting specifically of a pair of protium atoms) is highly flammable and will burn in air at a very wide range of concentrations between 4% and 75% by volume. The enthalpy of combustion is −286 kJ/mol:

2 H₂(g) + O₂(g) → 2 H₂O(l) + 572 kJ (286 kJ/mol)

Hydrogen gas forms explosive mixtures with air in concentrations from 4–74% and with chlorine at 5–95%. The explosive reactions may be triggered by spark, heat, or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C (932 °F). Pure hydrogen-oxygen flames emit ultraviolet light and with high oxygen mix are nearly invisible to the naked eye, as illustrated by the faint plume of the Space Shuttle Main Engine, compared to the highly visible plume of a Space Shuttle Solid Rocket Booster, which uses an ammonium perchlorate composite. The detection of a burning hydrogen leak may require a flame detector; such leaks can be very dangerous. Hydrogen flames in other conditions are blue, resembling blue natural gas flames.

The destruction of the Hindenburg airship was a notorious example of hydrogen combustion and the cause is still debated. The visible orange flames in that incident were the result of a rich mixture of hydrogen to oxygen combined with carbon compounds from the airship skin.

H₂ reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding hydrogen halides, hydrogen chloride and hydrogen fluoride, which are also potentially dangerous acids.

Electron energy levels

Depiction of a hydrogen atom with size of central proton shown, and the atomic diameter shown as about twice the Bohr model radius (image not to scale)

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 91 nm wavelength.

The energy levels of hydrogen can be calculated fairly accurately using the Bohr model of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the Sun. However, the atomic electron and proton are held together by electromagnetic force, while planets and celestial objects are held by gravity. Because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, and therefore only certain allowed energies.

A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation, Dirac equation or even the Feynman path integral formulation to calculate the probability density of the electron around the proton. The most complicated treatments allow for the small effects of special relativity and vacuum polarization. In the quantum mechanical treatment, the electron in a ground state hydrogen atom has no angular momentum at all—illustrating how the "planetary orbit" differs from electron motion.

Elemental molecular forms

First tracks observed in liquid hydrogen bubble chamber at the Bevatron

There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei. In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state with a molecular spin quantum number of 1 (⁄₂+⁄₂); in the parahydrogen form the spins are antiparallel and form a singlet with a molecular spin quantum number of 0 (⁄₂–⁄₂). At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form, also known as the "normal form". The equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but because the ortho form is an excited state and has a higher energy than the para form, it is unstable and cannot be purified. At very low temperatures, the equilibrium state is composed almost exclusively of the para form. The liquid and gas phase thermal properties of pure parahydrogen differ significantly from those of the normal form because of differences in rotational heat capacities, as discussed more fully in spin isomers of hydrogen. The ortho/para distinction also occurs in other hydrogen-containing molecules or functional groups, such as water and methylene, but is of little significance for their thermal properties.

The uncatalyzed interconversion between para and ortho H₂ increases with increasing temperature; thus rapidly condensed H₂ contains large quantities of the high-energy ortho form that converts to the para form very slowly. The ortho/para ratio in condensed H₂ is an important consideration in the preparation and storage of liquid hydrogen: the conversion from ortho to para is exothermic and produces enough heat to evaporate some of the hydrogen liquid, leading to loss of liquefied material. Catalysts for the ortho-para interconversion, such as ferric oxide, activated carbon, platinized asbestos, rare earth metals, uranium compounds, chromic oxide, or some nickel compounds, are used during hydrogen cooling.

Phases

• Gaseous hydrogen
• Liquid hydrogen
• Slush hydrogen
• Solid hydrogen
• Metallic hydrogen

Compounds

Covalent and organic compounds

While H₂ is not very reactive under standard conditions, it does form compounds with most elements. Hydrogen can form compounds with elements that are more electronegative, such as halogens (e.g., F, Cl, Br, I), or oxygen; in these compounds hydrogen takes on a partial positive charge. When bonded to fluorine, oxygen, or nitrogen, hydrogen can participate in a form of medium-strength noncovalent bonding with the hydrogen of other similar molecules, a phenomenon called hydrogen bonding that is critical to the stability of many biological molecules. Hydrogen also forms compounds with less electronegative elements, such as metals and metalloids, where it takes on a partial negative charge. These compounds are often known as hydrides.

Hydrogen forms a vast array of compounds with carbon called the hydrocarbons, and an even vaster array with heteroatoms that, because of their general association with living things, are called organic compounds. The study of their properties is known as organic chemistry and their study in the context of living organisms is known as biochemistry. By some definitions, "organic" compounds are only required to contain carbon. However, most of them also contain hydrogen, and because it is the carbon-hydrogen bond which gives this class of compounds most of its particular chemical characteristics, carbon-hydrogen bonds are required in some definitions of the word "organic" in chemistry. Millions of hydrocarbons are known, and they are usually formed by complicated synthetic pathways that seldom involve elementary hydrogen.

Hydrides

Compounds of hydrogen are often called hydrides, a term that is used fairly loosely. The term "hydride" suggests that the H atom has acquired a negative or anionic character, denoted H, and is used when hydrogen forms a compound with a more electropositive element. The existence of the hydride anion, suggested by Gilbert N. Lewis in 1916 for group 1 and 2 salt-like hydrides, was demonstrated by Moers in 1920 by the electrolysis of molten lithium hydride (LiH), producing a stoichiometry quantity of hydrogen at the anode. For hydrides other than group 1 and 2 metals, the term is quite misleading, considering the low electronegativity of hydrogen. An exception in group 2 hydrides is BeH
2, which is polymeric. In lithium aluminium hydride, the AlH−
4 anion carries hydridic centers firmly attached to the Al(III).

Although hydrides can be formed with almost all main-group elements, the number and combination of possible compounds varies widely; for example, more than 100 binary borane hydrides are known, but only one binary aluminium hydride. Binary indium hydride has not yet been identified, although larger complexes exist.

In inorganic chemistry, hydrides can also serve as bridging ligands that link two metal centers in a coordination complex. This function is particularly common in group 13 elements, especially in boranes (boron hydrides) and aluminium complexes, as well as in clustered carboranes.

Protons and acids

Oxidation of hydrogen removes its electron and gives H, which contains no electrons and a nucleus which is usually composed of one proton. That is why H+
 is often called a proton. This species is central to discussion of acids. Under the Brønsted–Lowry acid–base theory, acids are proton donors, while bases are proton acceptors.

A bare proton, H+
, cannot exist in solution or in ionic crystals because of its unstoppable attraction to other atoms or molecules with electrons. Except at the high temperatures associated with plasmas, such protons cannot be removed from the electron clouds of atoms and molecules, and will remain attached to them. However, the term 'proton' is sometimes used loosely and metaphorically to refer to positively charged or cationic hydrogen attached to other species in this fashion, and as such is denoted "H+
" without any implication that any single protons exist freely as a species.

To avoid the implication of the naked "solvated proton" in solution, acidic aqueous solutions are sometimes considered to contain a less unlikely fictitious species, termed the "hydronium ion" (H
3O+
). However, even in this case, such solvated hydrogen cations are more realistically conceived as being organized into clusters that form species closer to H
9O+
4. Other oxonium ionsare found when water is in acidic solution with other solvents.

Although exotic on Earth, one of the most common ions in the universe is the H+
3 ion, known as protonated molecular hydrogen or the trihydrogen cation.

Atomic hydrogen

NASA has investigated the use of atomic hydrogen as a rocket propellant. It could be stored in liquid helium to prevent it from recombining into molecular hydrogen. When the helium is vaporized, the atomic hydrogen would be released and combine back to molecular hydrogen. The result would be a intensely hot stream of hydrogen and helium gas. The liftoff weight of rockets could be reduced by 50% by this method.

Most interstellar hydrogen is in the form of atomic hydrogen because the atoms can seldom collide and combine. They are the source of the important 21 cm hydrogen line in astronomy at 1420 MHz.

Isotopes

Hydrogen discharge (spectrum) tube

Deuterium discharge (spectrum) tube

Protium, the most common isotope of hydrogen, has one proton and one electron. Unique among all stable isotopes, it has no neutrons (see diproton for a discussion of why others do not exist).

Hydrogen has three naturally occurring isotopes, denoted 1
H, 2
H and 3
H. Other, highly unstable nuclei (4
H to 7
H) have been synthesized in the laboratory but not observed in nature.

• 1
  H is the most common hydrogen isotope with an abundance of more than 99.98%. Because the nucleus of this isotope consists of only a single proton, it is given the descriptive but rarely used formal name protium.
• 2
  H, the other stable hydrogen isotope, is known as deuterium and contains one proton and one neutron in the nucleus. All deuterium in the universe is thought to have been produced at the time of the Big Bang, and has endured since that time. Deuterium is not radioactive, and does not represent a significant toxicity hazard. Water enriched in molecules that include deuterium instead of normal hydrogen is called heavy water. Deuterium and its compounds are used as a non-radioactive label in chemical experiments and in solvents for 1
  H-NMR spectroscopy. Heavy water is used as a neutron moderator and coolant for nuclear reactors. Deuterium is also a potential fuel for commercial nuclear fusion.
• 3
  H is known as tritium and contains one proton and two neutrons in its nucleus. It is radioactive, decaying into helium-3 through beta decay with a half-life of 12.32 years. It is so radioactive that it can be used in luminous paint, making it useful in such things as watches. The glass prevents the small amount of radiation from getting out. Small amounts of tritium are produced naturally by the interaction of cosmic rays with atmospheric gases; tritium has also been released during nuclear weapons tests. It is used in nuclear fusion reactions, as a tracer in isotope geochemistry, and in specialized self-powered lighting devices. Tritium has also been used in chemical and biological labeling experiments as a radiolabel.

Hydrogen is the only element that has different names for its isotopes in common use today. During the early study of radioactivity, various heavy radioactive isotopes were given their own names, but such names are no longer used, except for deuterium and tritium. The symbols D and T (instead of 2
H and 3
H) are sometimes used for deuterium and tritium, but the corresponding symbol for protium, P, is already in use for phosphorus and thus is not available for protium. In its nomenclatural guidelines, the International Union of Pure and Applied Chemistry (IUPAC) allows any of D, T, 2
H, and 3
H to be used, although 2
H and 3
H are preferred.

The exotic atom muonium (symbol Mu), composed of an antimuon and an electron, is also sometimes considered as a light radioisotope of hydrogen, due to the mass difference between the antimuon and the electron. Muonium was discovered in 1960. During the muon's 2.2 µs lifetime, muonium can enter into compounds such as muonium chloride (MuCl) or sodium muonide (NaMu), analogous to hydrogen chloride and sodium hydride respectively.

History

Discovery and use

In 1671, Robert Boyle discovered and described the reaction between iron filings and dilute acids, which results in the production of hydrogen gas. In 1766, Henry Cavendish was the first to recognize hydrogen gas as a discrete substance, by naming the gas from a metal-acid reaction "inflammable air". He speculated that "inflammable air" was in fact identical to the hypothetical substance called "phlogiston" and further finding in 1781 that the gas produces water when burned. He is usually given credit for the discovery of hydrogen as an element. In 1783, Antoine Lavoisier gave the element the name hydrogen (from the Greek ὑδρο- hydro meaning "water" and -γενής genes meaning "creator") when he and Laplacereproduced Cavendish's finding that water is produced when hydrogen is burned.

Antoine-Laurent de Lavoisier

Lavoisier produced hydrogen for his experiments on mass conservation by reacting a flux of steam with metallic iron through an incandescent iron tube heated in a fire. Anaerobic oxidation of iron by the protons of water at high temperature can be schematically represented by the set of following reactions:

   Fe +    H₂O → FeO + H₂

2 Fe + 3 H₂O → Fe₂O₃ + 3 H₂

3 Fe + 4 H₂O → Fe₃O₄ + 4 H₂

Many metals such as zirconium undergo a similar reaction with water leading to the production of hydrogen.

Hydrogen was liquefied for the first time by James Dewar in 1898 by using regenerative cooling and his invention, the vacuum flask. He produced solid hydrogen the next year. Deuterium was discovered in December 1931 by Harold Urey, and tritium was prepared in 1934 by Ernest Rutherford, Mark Oliphant, and Paul Harteck. Heavy water, which consists of deuterium in the place of regular hydrogen, was discovered by Urey's group in 1932.François Isaac de Rivaz built the first de Rivaz engine, an internal combustion engine powered by a mixture of hydrogen and oxygen in 1806. Edward Daniel Clarke invented the hydrogen gas blowpipe in 1819. The Döbereiner's lamp and limelight were invented in 1823.

The first hydrogen-filled balloon was invented by Jacques Charles in 1783. Hydrogen provided the lift for the first reliable form of air-travel following the 1852 invention of the first hydrogen-lifted airship by Henri Giffard. German count Ferdinand von Zeppelin promoted the idea of rigid airships lifted by hydrogen that later were called Zeppelins; the first of which had its maiden flight in 1900. Regularly scheduled flights started in 1910 and by the outbreak of World War I in August 1914, they had carried 35,000 passengers without a serious incident. Hydrogen-lifted airships were used as observation platforms and bombers during the war.

The first non-stop transatlantic crossing was made by the British airship R34 in 1919. Regular passenger service resumed in the 1920s and the discovery of helium reserves in the United States promised increased safety, but the U.S. government refused to sell the gas for this purpose. Therefore, H₂ was used in the Hindenburg airship, which was destroyed in a midair fire over New Jersey on 6 May 1937. The incident was broadcast live on radio and filmed. Ignition of leaking hydrogen is widely assumed to be the cause, but later investigations pointed to the ignition of the aluminized fabric coating by static electricity. But the damage to hydrogen's reputation as a lifting gas was already done and commercial hydrogen airship travel ceased. Hydrogen is still used, in preference to non-flammable but more expensive helium, as a lifting gas for weather balloons.

In the same year the first hydrogen-cooled turbogenerator went into service with gaseous hydrogen as a coolant in the rotor and the stator in 1937 at Dayton, Ohio, by the Dayton Power & Light Co.; because of the thermal conductivity of hydrogen gas, this is the most common type in its field today.

The nickel hydrogen battery was used for the first time in 1977 aboard the U.S. Navy's Navigation technology satellite-2 (NTS-2). For example, the ISS, Mars Odyssey and the Mars Global Surveyor are equipped with nickel-hydrogen batteries. In the dark part of its orbit, the Hubble Space Telescope is also powered by nickel-hydrogen batteries, which were finally replaced in May 2009, more than 19 years after launch and 13 years beyond their design life.

Role in quantum theory

Hydrogen emission spectrum lines in the visible range. These are the four visible lines of the Balmer series

Because of its simple atomic structure, consisting only of a proton and an electron, the hydrogen atom, together with the spectrum of light produced from it or absorbed by it, has been central to the development of the theory of atomic structure. Furthermore, study of the corresponding simplicity of the hydrogen molecule and the corresponding cation 
H

+
2

 brought understanding of the nature of the chemical bond, which followed shortly after the quantum mechanical treatment of the hydrogen atom had been developed in the mid-1920s.

One of the first quantum effects to be explicitly noticed (but not understood at the time) was a Maxwell observation involving hydrogen, half a century before full quantum mechanical theory arrived. Maxwell observed that the specific heat capacity of H₂ unaccountably departs from that of a diatomic gas below room temperature and begins to increasingly resemble that of a monatomic gas at cryogenic temperatures. According to quantum theory, this behavior arises from the spacing of the (quantized) rotational energy levels, which are particularly wide-spaced in H₂ because of its low mass. These widely spaced levels inhibit equal partition of heat energy into rotational motion in hydrogen at low temperatures. Diatomic gases composed of heavier atoms do not have such widely spaced levels and do not exhibit the same effect.

Antihydrogen (
H
) is the antimatter counterpart to hydrogen. It consists of an antiproton with a positron. Antihydrogen is the only type of antimatter atom to have been produced as of 2015.

Natural occurrence

NGC 604, a giant region of ionized hydrogen in the Triangulum Galaxy

Hydrogen, as atomic H, is the most abundant chemical element in the universe, making up 75% of normal matter by mass and more than 90% by number of atoms. (Most of the mass of the universe, however, is not in the form of chemical-element type matter, but rather is postulated to occur as yet-undetected forms of mass such as dark matter and dark energy.) This element is found in great abundance in stars and gas giant planets. Molecular clouds of H₂ are associated with star formation. Hydrogen plays a vital role in powering stars through the proton-proton reaction and the CNO cycle nuclear fusion.

Throughout the universe, hydrogen is mostly found in the atomic and plasma states, with properties quite different from those of molecular hydrogen. As a plasma, hydrogen's electron and proton are not bound together, resulting in very high electrical conductivity and high emissivity (producing the light from the Sun and other stars). 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 is found in the neutral atomic state in the interstellar medium. The large amount of neutral hydrogen found in the damped Lyman-alpha systems is thought to dominate the cosmological baryonic density of the Universe up to redshift z=4.

Under ordinary conditions on Earth, elemental hydrogen exists as the diatomic gas, H₂. However, hydrogen gas is very rare in the Earth's atmosphere (1 ppm by volume) because of its light weight, which enables it to escape from Earth's gravity more easily than heavier gases. However, hydrogen is the third most abundant element on the Earth's surface, mostly in the form of chemical compounds such as hydrocarbons and water. Hydrogen gas is produced by some bacteria and algae and is a natural component of flatus, as is methane, itself a hydrogen source of increasing importance.

A molecular form called protonated molecular hydrogen (H+
3) is found in the interstellar medium, where it is generated by ionization of molecular hydrogen from cosmic rays. This charged ion has also been observed in the upper atmosphere of the planet Jupiter. The ion is relatively stable in the environment of outer space due to the low temperature and density. H+
3 is one of the most abundant ions in the Universe, and it plays a notable role in the chemistry of the interstellar medium. Neutral triatomic hydrogen H₃ can exist only in an excited form and is unstable. By contrast, the positive hydrogen molecular ion (H+
2) is a rare molecule in the universe.

Production

H
2 is produced in chemistry and biology laboratories, often as a by-product of other reactions; in industry for the hydrogenation of unsaturated substrates; and in nature as a means of expelling reducing equivalents in biochemical reactions.

Electrolysis of Water

The electrolysis of water is a simple method of producing hydrogen. A low voltage current is run through the water, and gaseous oxygen forms at the anode while gaseous hydrogen forms at the cathode. Typically the cathode is made from platinum or another inert metal when producing hydrogen for storage. If, however, the gas is to be burnt on site, oxygen is desirable to assist the combustion, and so both electrodes would be made from inert metals. (Iron, for instance, would oxidize, and thus decrease the amount of oxygen given off.) The theoretical maximum efficiency (electricity used vs. energetic value of hydrogen produced) is in the range 88-94%.

2 H2O(l) → 2 H2(g) + O2(g)

When determining the electrical efficiency of PEM (proton exchange membrane) electrolysis, the higher heat value (HHV) is used. This is because the catalyst layer interacts with water as steam. As the process operates at 80 °C for PEM electrolysers the waste heat can be redirected through the system to create the steam, resulting in a higher overall electrical efficiency. The lower heat value (LHV) must be used for alkaline electrolysers as the process within these electrolysers requires water in liquid form and uses alkalinity to facilitate the breaking of the bond holding the hydrogen and oxygen atoms together. The lower heat value must also be used for fuel cells, as steam is the output rather than input.

Steam reforming

Hydrogen is often produced using natural gas, which involves the removal of hydrogen from hydrocarbons at very high temperatures, with about 95% of hydrogen production coming from steam reforming around year 2000. Commercial bulk hydrogen is usually produced by the steam reforming of natural gas. At high temperatures (1000–1400 K, 700–1100 °C or 1300–2000 °F), steam (water vapor) reacts with methane to yield carbon monoxide and H
2.

CH4 + H2O → CO + 3 H2

This reaction is favored at low pressures but is nonetheless conducted at high pressures (2.0  MPa, 20 atm or 600 inHg). This is because high-pressure H
2 is the most marketable product and pressure swing adsorption (PSA) purification systems work better at higher pressures. The product mixture is known as "synthesis gas" because it is often used directly for the production of methanol and related compounds. Hydrocarbons other than methane can be used to produce synthesis gas with varying product ratios. One of the many complications to this highly optimized technology is the formation of coke or carbon:

CH4 → C + 2 H2

Consequently, steam reforming typically employs an excess of H
2O. Additional hydrogen can be recovered from the steam by use of carbon monoxide through the water gas shift reaction, especially with an iron oxide catalyst. This reaction is also a common industrial source of carbon dioxide:

CO + H2O → CO2 + H2

Other important methods for H
2 production include partial oxidation of hydrocarbons:

2 CH4 + O2 → 2 CO + 4 H2

and the coal reaction, which can serve as a prelude to the shift reaction above:

C + H2O → CO + H2

Hydrogen is sometimes produced and consumed in the same industrial process, without being separated. In the Haber process for the production of ammonia, hydrogen is generated from natural gas. Electrolysis of brine to yield chlorine also produces hydrogen as a co-product.

Metal-acid

In the laboratory, H
2 is usually prepared by the reaction of dilute non-oxidizing acids on some reactive metals such as zinc with Kipp's apparatus.

Zn + 2 H+ → Zn2+ + H2

Aluminium can also produce H
2 upon treatment with bases:

2 Al + 6 H2O + 2 OH− → 2 Al(OH)− 4 + 3 H2

An alloy of aluminium and gallium in pellet form added to water can be used to generate hydrogen. The process also produces alumina, but the expensive gallium, which prevents the formation of an oxide skin on the pellets, can be re-used. This has important potential implications for a hydrogen economy, as hydrogen can be produced on-site and does not need to be transported.

Thermochemical

There are more than 200 thermochemical cycles which can be used for water splitting, around a dozen of these cycles such as the iron oxide cycle, cerium(IV) oxide–cerium(III) oxide cycle, zinc zinc-oxide cycle, sulfur-iodine cycle, copper-chlorine cycle and hybrid sulfur cycle are under research and in testing phase to produce hydrogen and oxygen from water and heat without using electricity. A number of laboratories (including in France, Germany, Greece, Japan, and the USA) are developing thermochemical methods to produce hydrogen from solar energy and water.

Anaerobic corrosion

Under anaerobic conditions, iron and steel alloys are slowly oxidized by the protons of water concomitantly reduced in molecular hydrogen (H
2). The anaerobic corrosion of iron leads first to the formation of ferrous hydroxide (green rust) and can be described by the following reaction:

Fe + 2 H2O → Fe(OH)2 + H2

In its turn, under anaerobic conditions, the ferrous hydroxide (Fe(OH)
2) can be oxidized by the protons of water to form magnetite and molecular hydrogen. This process is described by the Schikorr reaction:

3 Fe(OH)2 → Fe3O4 + 2 H2O + H2

ferrous hydroxide → magnetite + water + hydrogen

The well crystallized magnetite (Fe
3O
4) is thermodynamically more stable than the ferrous hydroxide (Fe(OH)
2).

This process occurs during the anaerobic corrosion of iron and steel in oxygen-free groundwater and in reducing soils below the water table.

Geological occurrence: the serpentinization reaction

In the absence of atmospheric oxygen (O
2), in deep geological conditions prevailing far away from Earth atmosphere, hydrogen (H
2) is produced during the process of serpentinization by the anaerobic oxidation by the water protons (H) of the ferrous (Fe) silicate present in the crystal lattice of the fayalite (Fe
2SiO
4, the olivine iron-endmember). The corresponding reaction leading to the formation of magnetite (Fe
3O
4), quartz (SiO
2) and hydrogen (H
2) is the following:

3Fe2SiO4 + 2 H2O → 2 Fe3O4 + 3 SiO2 + 3 H2

fayalite + water → magnetite + quartz + hydrogen

This reaction closely resembles the Schikorr reaction observed in the anaerobic oxidation of the ferrous hydroxide in contact with water.

Formation in transformers

From all the fault gases formed in power transformers, hydrogen is the most common and is generated under most fault conditions; thus, formation of hydrogen is an early indication of serious problems in the transformer's life cycle.

Applications

Consumption in processes

Large quantities of H
2 are needed in the petroleum and chemical industries. The largest application of H
2 is for the processing ("upgrading") of fossil fuels, and in the production of ammonia. The key consumers of H
2 in the petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking. H
2 has several other important uses. H
2 is used as a hydrogenating agent, particularly in increasing the level of saturation of unsaturated fats and oils (found in items such as margarine), and in the production of methanol. It is similarly the source of hydrogen in the manufacture of hydrochloric acid. H
2 is also used as a reducing agent of metallic ores.

Hydrogen is highly soluble in many rare earth and transition metals and is soluble in both nanocrystalline and amorphous metals. Hydrogen solubility in metals is influenced by local distortions or impurities in the crystal lattice. These properties may be useful when hydrogen is purified by passage through hot palladium disks, but the gas's high solubility is a metallurgical problem, contributing to the embrittlement of many metals, complicating the design of pipelines and storage tanks.

Apart from its use as a reactant, H
2 has wide applications in physics and engineering. It is used as a shielding gas in welding methods such as atomic hydrogen welding. H₂ is used as the rotor coolant in electrical generators at power stations, because it has the highest thermal conductivity of any gas. Liquid H₂ is used in cryogenic research, including superconductivitystudies. Because H
2 is lighter than air, having a little more than ⁄₁₄ of the density of air, it was once widely used as a lifting gas in balloons and airships.

In more recent applications, hydrogen is used pure or mixed with nitrogen (sometimes called forming gas) as a tracer gas for minute leak detection. Applications can be found in the automotive, chemical, power generation, aerospace, and telecommunications industries. Hydrogen is an authorized food additive (E 949) that allows food package leak testing among other anti-oxidizing properties.

Hydrogen's rarer isotopes also each have specific applications. Deuterium (hydrogen-2) is used in nuclear fission applications as a moderator to slow neutrons, and in nuclear fusionreactions. 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.

The triple point temperature of equilibrium hydrogen is a defining fixed point on the ITS-90 temperature scale at 13.8033 kelvins.

Coolant

Hydrogen is commonly used in power stations as a coolant in generators due to a number of favorable properties that are a direct result of its light diatomic molecules. These include low density, low viscosity, and the highest specific heat and thermal conductivity of all gases.

Energy carrier

Hydrogen is not an energy resource, except in the hypothetical context of commercial nuclear fusion power plants using deuterium or tritium, a technology presently far from development. The Sun's energy comes from nuclear fusion of hydrogen, but this process is difficult to achieve controllably on Earth. Elemental hydrogen from solar, biological, or electrical sources requires more energy to make than is obtained by burning it, so in these cases hydrogen functions as an energy carrier, like a battery. Hydrogen may be obtained from fossil sources (such as methane), but these sources are unsustainable.

The energy density per unit volume of both liquid hydrogen and compressed hydrogen gas at any practicable pressure is significantly less than that of traditional fuel sources, although the energy density per unit fuel mass is higher. Nevertheless, elemental hydrogen has been widely discussed in the context of energy, as a possible future carrier of energy on an economy-wide scale. For example, CO
2 sequestration followed by carbon capture and storage could be conducted at the point of H
2 production from fossil fuels. Hydrogen used in transportation would burn relatively cleanly, with some NO
_x emissions, but without carbon emissions. However, the infrastructure costs associated with full conversion to a hydrogen economy would be substantial. Fuel cells can convert hydrogen and oxygen directly to electricity more efficiently than internal combustion engines.

Semiconductor industry

Hydrogen is employed to saturate broken ("dangling") bonds of amorphous silicon and amorphous carbon that helps stabilizing material properties. It is also a potential electron donor in various oxide materials, including  ZnO, SnO₂, CdO, MgO, ZrO₂, HfO₂, La₂O₃, Y₂O₃, TiO₂, SrTiO₃, LaAlO₃, SiO₂, Al
₂O₃, ZrSiO₄, HfSiO₄, and SrZrO₃.

Biological reactions

H₂ is a product of some types of anaerobic metabolism and is produced by several microorganisms, usually via reactions catalyzed by iron- or nickel-containing enzymes called hydrogenases. These enzymes catalyze the reversible redox reaction between H₂ and its component two protons and two electrons. Creation of hydrogen gas occurs in the transfer of reducing equivalents produced during pyruvate fermentation to water. The natural cycle of hydrogen production and consumption by organisms is called the hydrogen cycle.

Water splitting, in which water is decomposed into its component protons, electrons, and oxygen, occurs in the light reactions in all photosynthetic organisms. Some such organisms, including the alga Chlamydomonas reinhardtii and cyanobacteria, have evolved a second step in the dark reactions in which protons and electrons are reduced to form H₂ gas by specialized hydrogenases in the chloroplast. Efforts have been undertaken to genetically modify cyanobacterial hydrogenases to efficiently synthesize H₂ gas even in the presence of oxygen.Efforts have also been undertaken with genetically modified alga in a bioreactor.

Safety and precautions

Hydrogen poses a number of hazards to human safety, from potential detonations and fires when mixed with air to being an asphyxiant in its pure, oxygen-free form. In addition, liquid hydrogen is a cryogen and presents dangers (such as frostbite) associated with very cold liquids. Hydrogen dissolves in many metals, and, in addition to leaking out, may have adverse effects on them, such as hydrogen embrittlement, leading to cracks and explosions. Hydrogen gas leaking into external air may spontaneously ignite. Moreover, hydrogen fire, while being extremely hot, is almost invisible, and thus can lead to accidental burns.

Even interpreting the hydrogen data (including safety data) is confounded by a number of phenomena. Many physical and chemical properties of hydrogen depend on the parahydrogen/orthohydrogen ratio (it often takes days or weeks at a given temperature to reach the equilibrium ratio, for which the data is usually given). Hydrogen detonation parameters, such as critical detonation pressure and temperature, strongly depend on the container geometry.

Retrieved from: https://en.wikipedia.org/wiki/Hydrogen

Definition

The crystal structure of sodium chloride, NaCl, a typical ionic compound. The purple spheres represent sodium cations, Na, and the green spheres represent chloride anions, Cl.

In chemistry, an ionic compound is a chemical compound composed of ions held together by electrostatic forces termed ionic bonding. The compound is neutral overall, but consists of positively charged ions called cations and negatively charged ions called anions. These can be simple ions such as the sodium (Na) and chloride (Cl) in sodium chloride, or polyatomic species such as the ammonium (NH+
4) and carbonate (CO2−
3) ions in ammonium carbonate. Individual ions within an ionic compound usually have multiple nearest neighbours, so are not considered to be part of molecules, but instead part of a continuous three-dimensional network, usually in a crystalline structure.

Ionic compounds containing hydrogen ions (H) are classified as acids, and those containing basic ions hydroxide (OH) or oxide (O) are classified as bases. Ionic compounds without these ions are also known as salts and can be formed by acid–base reactions. Ionic compounds can also be produced from their constituent ions by evaporation of their solvent, precipitation, freezing, a solid-state reaction, or the electron transfer reaction of reactivemetals with reactive non-metals, such as halogen gases.

Ionic compounds typically have high melting and boiling points, and are hard and brittle. As solids they are almost always electrically insulating, but when melted or dissolved they become highly conductive, because the ions are mobilized.

History of discovery

The word ion is the Greek ἰόν, ion, "going", the present participle of ἰέναι, ienai, "to go". This term was introduced by English physicist and chemist Michael Faraday in 1834 for the then-unknown species that goes from one electrode to the other through an aqueous medium.

X-ray spectrometer developed by Bragg

In 1913 the crystal structure of sodium chloride was determined by William Henry Bragg and William Lawrence Bragg. This revealed that there were six equidistant nearest-neighbours for each atom, demonstrating that the constituents were not arranged in molecules or finite aggregates, but instead as a network with long-range crystalline order. Many other inorganic compounds were also found to have similar structural features. These compounds were soon described as being constituted of ions rather than neutral atoms, but proof of this hypothesis was not found until the mid-1920s, when X-ray reflection experiments (which detect the density of electrons), were performed.

Principal contributors to the development of a theoretical treatment of ionic crystal structures were Max Born, Fritz Haber, Alfred Landé, Erwin Madelung, Paul Peter Ewald, and Kazimierz Fajans. Born predicted crystal energies based on the assumption of ionic constituents, which showed good correspondence to thermochemical measurements, further supporting the assumption.

Formation

Halite, the mineral form of sodium chloride, forms when salty water evaporates leaving the ions behind.

Ionic compounds can be produced from their constituent ions by evaporation, precipitation, or freezing. Reactive metals such as the alkali metals can react directly with the highly electronegative halogen gases to form an ionic product. They can also be synthesized as the product of a high temperature reaction between solids.

If the ionic compound is soluble in a solvent, it can be obtained as a solid compound by evaporating the solvent from this electrolyte solution. As the solvent is evaporated, the ions do not go into the vapour, but stay in the remaining solution, and when they become sufficiently concentrated, nucleation occurs, and they crystallize into an ionic compound. This process occurs widely in nature, and is the means of formation of the evaporiteminerals. Another method of recovering the compound from solution involves saturating a solution at high temperature and then reducing the solubility by reducing the temperature until the solution is supersaturated and the solid compound nucleates.

Insoluble ionic compounds can be precipitated by mixing two solutions, one with the cation and one with the anion in it. Because all solutions are electrically neutral, the two solutions mixed must also contain counterions of the opposite charges. To ensure that these do not contaminate the precipitated ionic compound, it is important to ensure they do not also precipitate. If the two solutions have hydrogen ions and hydroxide ions as the counterions, they will react with one another in what is called an acid–base reaction or a neutralization reaction to form water. Alternately the counterions can be chosen to ensure that even when combined into a single solution they will remain soluble as spectator ions.

If the solvent is water in either the evaporation or precipitation method of formation, in many cases the ionic crystal formed also includes water of crystallization, so the product is known as a hydrate, and can have very different chemical properties.

Molten salts will solidify on cooling to below their freezing point. This is sometimes used for the solid-state synthesis of complex ionic compounds from solid reactants, which are first melted together. In other cases, the solid reactants do not need to be melted, but instead can react through a solid-state reaction route. In this method the reactants are repeatedly finely ground into a paste, and then heated to a temperature where the ions in neighbouring reactants can diffuse together during the time the reactant mixture remains in the oven. Other synthetic routes use a solid precursor with the correct stoichiometric ratio of non-volatile ions, which is heated to drive off other species.

In some reactions between highly reactive metals (usually from Group 1 or Group 2) and highly electronegative halogen gases, or water, the atoms can be ionized by electron transfer, a process thermodynamically understood using the Born–Haber cycle.

Bonding

A schematic electron shell diagram of sodium and fluorine atoms undergoing a redox reaction to form sodium fluoride. Sodium loses its outer electron to give it a stable electron configuration, and this electron enters the fluorine atom exothermically. The oppositely charged ions – typically a great many of them – are then attracted to each other to form a solid.

Ions in ionic compounds are primarily held together by the electrostatic forces between the charge distribution of these bodies, and in particular the ionic bond resulting from the long-ranged Coulomb attraction between the net negative charge of the anions and net positive charge of the cations. There is also a small additional attractive force from van der Waals interactions which contributes only around 1–2% of the cohesive energy for small ions. When a pair of ions comes close enough for their outer electron shells (most simple ions have closed shells) to overlap, a short-ranged repulsive force occurs, due to the Pauli exclusion principle. The balance between these forces leads to a potential energy well with a minimum energy when the nuclei are separated by a specific equilibrium distance.

If the electronic structure of the two interacting bodies is affected by the presence of one another, covalent interactions (non-ionic) also contribute to the overall energy of the compound formed. Ionic compounds are rarely purely ionic, i.e. held together only by electrostatic forces. The bonds between even the most electronegative/electropositive pairs such as those in caesium fluoride exhibit a small degree of covalency. Conversely, covalent bonds between unlike atoms often exhibit some charge separation and can be considered to have a partial ionic character. The circumstances under which a compound will have ionic or covalent character can typically be understood using Fajans' rules, which use only charges and the sizes of each ion. According to these rules, compounds with the most ionic character will have large positive ions with a low charge, bonded to a small negative ion with a high charge. More generally HSAB theory can be applied, whereby the compounds with the most ionic character are those consisting of hard acids and hard bases: small, highly charged ions with a high difference in electronegativities between the anion and cation. This difference in electronegativities means that the charge separation, and resulting dipole moment, is maintained even when the ions are in contact (the excess electrons on the anions are not transferred or polarized to neutralize the cations).

Structure

The unit cell of the zinc blende structure

Ions typically pack into extremely regular crystalline structures, in an arrangement that minimizes the lattice energy (maximizing attractions and minimizing repulsions). The lattice energy is the summation of the interaction of all sites with all other sites. For unpolarizable spherical ions only the charges and distances are required to determine the electrostatic interaction energy. For any particular ideal crystal structure, all distances are geometrically related to the smallest internuclear distance. So for each possible crystal structure, the total electrostatic energy can be related to the electrostatic energy of unit charges at the nearest neighbour distance by a multiplicative constant called the Madelung constant that can be efficiently computed using an Ewald sum. When a reasonable form is assumed for the additional repulsive energy, the total lattice energy can be modelled using the Born–Landé equation, the Born–Mayer equation, or in the absence of structural information, the Kapustinskii equation.

Using an even simpler approximation of the ions as impenetrable hard spheres, the arrangement of anions in these systems are often related to close-packed arrangements of spheres, with the cations occupying tetrahedral or octahedral interstices. Depending on the stoichiometry of the ionic compound, and the coordination (principally determined by the radius ratio) of cations and anions, a variety of structures are commonly observed,and theoretically rationalized by Pauling's rules.

Common ionic compound structures with close-packed anions

Stoichiometry	Cation:anion coordination	Interstitial sites		Cubic close packing of anions		Hexagonal close packing of anions
		occupancy	critical radius ratio	name	Madelung constant	name	Madelung constant
MX	6:6	all octahedral	0.4142	sodium chloride	1.747565	nickeline	<1.73
	4:4	alternate tetrahedral	0.2247	zinc blende	1.6381	wurtzite	1.641
MX2	8:4	all tetrahedral	0.2247	fluorite	5.03878
	6:3	half octahedral (alternate layers fully occupied)	0.4142	cadmium chloride	5.61	cadmium iodide	4.71
MX3	6:2	one-third octahedral	0.4142	rhodium(III) bromide	6.67	bismuth iodide	8.26
M2X3	6:4	two-thirds octahedral	0.4142			corundum	25.0312
ABO3		two-thirds octahedral	0.4142			ilmenite	depends on charges and structure
AB2O4		one-eighth tetrahedral and one-half octahedral	rA/rO = 0.2247, rB/rO = 0.4142	spinel, inverse spinel	depends on cation site distributions	olivine	depends on cation site distributions

In some cases the anions take on a simple cubic packing, and the resulting common structures observed are:

Common ionic compound structures with simple cubic packed anions

Stoichiometry	Cation:anion coordination	Interstitial sites occupied	Example structure
			name	critical radius ratio	Madelung constant
MX	8:8	entirely filled	cesium chloride	0.7321	1.762675
MX2	8:4	half filled	calcium fluoride
M2X	4:8	half filled	lithium oxide

Some ionic liquids, particularly with mixtures of anions or cations, can be cooled rapidly enough that there is not enough time for crystal nucleation to occur, so an ionic glass is formed (with no long-range order).

Defects

Frenkel defect

Schottky defect

Within an ionic crystal, there will usually be some point defects, but to maintain electroneutrality, these defects come in pairs.Frenkel defects consist of a cation vacancy paired with a cation interstitial and can be generated anywhere in the bulk of the crystal,occurring most commonly in compounds with a low coordination number and cations that are much smaller than the anions.Schottky defects consist of one vacancy of each type, and are generated at the surfaces of a crystal, occurring most commonly in compounds with a high coordination number and when the anions and cations are of similar size. If the cations have multiple possible oxidation states, then it is possible for cation vacancies to compensate for electron deficiencies on cation sites with higher oxidation numbers, resulting in a non-stoichiometric compound. Another non-stoichiometric possibility is the formation of an F-center, a free electron occupying an anion vacancy. When the compound has three or more ionic components, even more defect types are possible. All of these point defects can be generated via thermal vibrations and have an equilibrium concentration. Because they are energetically costly, but entropically beneficial, they occur in greater concentration at higher temperatures. Once generated, these pairs of defects can diffuse mostly independently of one another, by hopping between lattice sites. This defect mobility is the source of most transport phenomena within an ionic crystal, including diffusion and solid state ionic conductivity. When vacancies collide with interstitials (Frenkel), they can recombine and annihilate one another. Similarly vacancies are removed when they reach the surface of the crystal (Schottky). Defects in the crystal structure generally expand the lattice parameters, reducing the overall density of the crystal.Defects also result in ions in distinctly different local environments, which causes them to experience a different crystal-field symmetry, especially in the case of different cations exchanging lattice sites. This results in a different splitting of d-electron orbitals, so that the optical absorption (and hence colour) can change with defect concentration.

Properties

Acidity/basicity

Ionic compounds containing hydrogen ions (H) are classified as acids, and those containing electropositive cations and basic anions ions hydroxide (OH) or oxide (O) are classified as bases. Other ionic compounds are known as salts and can be formed by acid–base reactions. If the compound is the result of a reaction between a strong acid and a weak base, the result is an acidic salt. If it is the result of a reaction between a strong base and a weak acid, the result is a basic salt. If it is the result of a reaction between a strong acid and a strong base, the result is a neutral salt. Weak acids reacted with weak bases can produce ionic compounds with both the conjugate base ion and conjugate acid ion, such as ammonium acetate.

Some ions are classed as amphoteric, being able to react with either an acid or a base. This is also true of some compounds with ionic character, typically oxides or hydroxides of less-electropositive metals (so the compound also has significant covalent character), such as zinc oxide, aluminium hydroxide, aluminium oxide and lead(II) oxide.

Melting and boiling points

Electrostatic forces between particles are strongest when the charges are high, and the distance between the nuclei of the ions is small. In such cases, the compounds generally have very high melting and boiling points and a low vapour pressure. Trends in melting points can be even better explained when the structure and ionic size ratio is taken into account. Above their melting point ionic solids melt and become molten salts (although some ionic compounds such as aluminium chloride and iron(III) chloride show molecule-like structures in the liquid phase). Inorganic compounds with simple ions typically have small ions, and thus have high melting points, so are solids at room temperature. Some substances with larger ions, however, have a melting point below or near room temperature (often defined as up to 100 °C), and are termed ionic liquids. Ions in ionic liquids often have uneven charge distributions, or bulky substituents like hydrocarbon chains, which also play a role in determining the strength of the interactions and propensity to melt.

Even when the local structure and bonding of an ionic solid is disrupted sufficiently to melt it, there are still strong long-range electrostatic forces of attraction holding the liquid together and preventing ions boiling to form a gas phase. This means that even room temperature ionic liquids have low vapour pressures, and require substantially higher temperatures to boil.Boiling points exhibit similar trends to melting points in terms of the size of ions and strength of other interactions. When vapourized, the ions are still not freed of one another. For example, in the vapour phase sodium chloride exists as diatomic "molecules".

Brittleness

Most ionic compounds are very brittle. Once they reach the limit of their strength, they cannot deform malleably, because the strict alignment of positive and negative ions must be maintained. Instead the material undergoes fracture via cleavage. As the temperature is elevated (usually close to the melting point) a ductile–brittle transition occurs, and plastic flowbecomes possible by the motion of dislocations.

Compressibility

The compressibility of an ionic compound is strongly determined by its structure, and in particular the coordination number. For example, halides with the caesium chloride structure (coordination number 8) are less compressible than those with the sodium chloride structure (coordination number 6), and less again than those with a coordination number of 4.

Solubility

When ionic compounds dissolve, the individual ions dissociate and are solvated by the solvent and dispersed throughout the resulting solution. Because the ions are released into solution when dissolved, and can conduct charge, soluble ionic compounds are the most common class of strong electrolytes, and their solutions have a high electrical conductivity.

The aqueous solubility of a variety of ionic compounds as a function of temperature. Some compounds exhibiting unusual solubility behaviour have been included.

The solubility is highest in polar solvents (such as water) or ionic liquids, but tends to be low in nonpolar solvents (such as petrol/gasoline). This is principally because the resulting ion–dipole interactions are significantly stronger than ion-induced dipole interactions, so the heat of solution is higher. When the oppositely charged ions in the solid ionic lattice are surrounded by the opposite pole of a polar molecule, the solid ions are pulled out of the lattice and into the liquid. If the solvation energy exceeds the lattice energy, the negative net enthalpy change of solution provides a thermodynamic drive to remove ions from their positions in the crystal and dissolve in the liquid. In addition, the entropy change of solution is usually positive for most solid solutes like ionic compounds, which means that their solubility increases when the temperature increases. There are some unusual ionic compounds such as cerium(III) sulfate, where this entropy change is negative, due to extra order induced in the water upon solution, and the solubility decreases with temperature.

Electrical conductivity

Although ionic compounds contain charged atoms or clusters, these materials do not typically conduct electricity to any significant extent when the substance is solid. In order to conduct, the charged particles must be mobile rather than stationary in a crystal lattice. This is achieved to some degree at high temperatures when the defect concentration increases the ionic mobility and solid state ionic conductivity is observed. When the ionic compounds are dissolved in a liquid or are melted into a liquid, they can conduct electricity because the ions become completely mobile. This conductivity gain upon dissolving or melting is sometimes used as a defining characteristic of ionic compounds.

In some unusual ionic compounds: fast ion conductors, and ionic glasses, one or more of the ionic components has a significant mobility, allowing conductivity even while the material as a whole remains solid. This is often highly temperature dependant, and may be the result of either a phase change or a high defect concentration. These materials are used in all solid-state supercapacitors, batteries, and fuel cells, and in various kinds of chemical sensors.

Colour

Anhydrous cobalt(II) chloride,
CoCl₂

Cobalt(II) chloride hexahydrate,
CoCl₂•6H₂O

The colour of an ionic compound is often different to the colour of an aqueous solution containing the constituent ions,or the hydrated form of the same compound.

The anions in compounds with bonds with the most ionic character tend to be colourless (with an absorption band in the ultraviolet part of the spectrum). In compounds with less ionic character, their colour deepens through yellow, orange, red and black (as the absorption band shifts to longer wavelengths into the visible spectrum).

The absorption band of simple cations shift toward shorter wavelength when they are involved in more covalent interactions. This occurs during hydration of metal ions, so colourless anhydrous ionic compounds with an anion absorbing in the infrared can become colourful in solution.

Uses

Ionic compounds have long had a wide variety of uses and applications. Many minerals are ionic. Humans have processed common salt (sodium chloride) for over 8000 years, using it first as a food seasoning and preservative, and now also in manufacturing, agriculture, water conditioning, for de-icing roads, and many other uses. Many ionic compounds are so widely used in society that they go by common names unrelated to their chemical identity. Examples of this include borax, calomel, milk of magnesia, muriatic acid, oil of vitriol, saltpeter, and slaked lime.

Soluble ionic compounds like salt can easily be dissolved to provide electrolyte solutions. This is a simple way to control the concentration and ionic strength. The concentration of solutes affects many colligative properties, including increasing the osmotic pressure, and causing freezing-point depression and boiling-point elevation. Because the solutes are charged ions they also increase the electrical conductivity of the solution. The increased ionic strength reduces the thickness of the electrical double layer around colloidal particles, and therefore the stability of emulsions and suspensions.

The chemical identity of the ions added is also important in many uses. For example, fluoride containing compounds are dissolved to supply fluoride ions for water fluoridation.

Solid ionic compounds have long been used as paint pigments, and are resistant to organic solvents, but are sensitive to acidity or basicity. Since 1801 pyrotechnicians have described and widely used metal-containing ionic compounds as sources of colour in fireworks. Under intense heat, the electrons in the metal ions or small molecules can be excited. These electrons later return to lower energy states, and release light with a colour spectrum characteristic of the species present.

In chemistry, ionic compounds are often used as precursors for high-temperature solid-state synthesis.

Many metals are geologically most abundant as ionic compounds within ores. To obtain the elemental materials, these ores are processed by smelting or electrolysis, in which redox reactions occur (often with a reducing agent such as carbon) such that the metal ions gain electrons to become neutral atoms.

Nomenclature

According to the nomenclature recommended by IUPAC, ionic compounds are named according to their composition, not their structure. In the most simple case of a binary ionic compound with no possible ambiguity about the charges and thus the stoichiometry, the common name is written using two words. The name of the cation (the unmodified element name for monatomic cations) comes first, followed by the name of the anion. For example, MgCl₂ is named magnesium chloride, and Na₂SO₄ is named sodium sulfate (SO2−
4, sulfate, is an example of a polyatomic ion). To obtain the empirical formula from these names, the stoichiometry can be deduced from the charges on the ions, and the requirement of overall charge neutrality.

If there are multiple different cations and/or anions, multiplicative prefixes (di-, tri-, tetra-, ...) are often required to indicate the relative compositions, and cations then anions are listed in alphabetical order. For example, KMgCl₃ is named magnesium potassium trichloride to distinguish it from K₂MgCl₄, magnesium dipotassium tetrachloride (note that in both the empirical formula and the written name, the cations appear in alphabetical order, but the order varies between them because the symbol for potassium is K). When one of the ions already has a multiplicative prefix within its name, the alternate multiplicative prefixes (bis-, tris-, tetrakis-, ...) are used. For example, Ba(BrF₄)₂ is named barium bis(tetrafluoridobromate).

Compounds containing one or more elements which can exist in a variety of charge/oxidation states will have a stoichiometry that depends on which oxidation states are present, to ensure overall neutrality. This can be indicated in the name by specifying either the oxidation state of the elements present, or the charge on the ions. Because of the risk of ambiguity in allocating oxidation states, IUPAC prefers direct indication of the ionic charge numbers. These are written as an arabic integer followed by the sign (... , 2−, 1−, 1+, 2+, ...) in parentheses directly after the name of the cation (without a space separating them). For example, FeSO₄ is named iron(2+) sulfate (with the 2+ charge on the Fe ions balancing the 2− charge on the sulfate ion), whereas Fe₂(SO₄)₃ is named iron(3+) sulfate (because the two iron ions in each formula unit each have a charge of 3+, to balance the 2− on each of the three sulfate ions).Stock nomenclature, still in common use, writes the oxidation number in Roman numerals (... , −II, −I, 0, I, II, ...). So the examples given above would be named iron(II) sulfate and iron(III) sulfate respectively. For simple ions the ionic charge and the oxidation number are identical, but for polyatomic ions they often differ. For example, the uranyl(2+) ion, UO2+
2, has uranium in an oxidation state of +6, so would be called a dioxouranium(VI) ion in Stock nomenclature. An even older naming system for metal cations, also still widely used, appended the suffixes -ous and -ic to the Latin root of the name, to give special names for the low and high oxidation states. For example, this scheme uses "ferrous" and "ferric", for iron(II) and iron(III) respectively, so the examples given above were classically named ferrous sulfate and ferric sulfate.

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