Salt (chemistry)

File:NaCl bonds.svg

The crystal structure of sodium chloride, NaCl, a typical salt. The purple spheres represent sodium cations, Na⁺, and the green spheres represent chloride anions, Cl⁻. The yellow stipples show the electrostatic forces.

In chemistry, a salt or ionic compound is a chemical compound consisting of an assembly of positively charged ions (cations) and negatively charged ions (anions),<ref>Template:GoldBookRef</ref> which results in a compound with no net electric charge (electrically neutral). The constituent ions are held together by electrostatic forces termed ionic bonds.

The component ions in a salt can be either inorganic, such as chloride (Cl⁻), or organic, such as acetate (Template:Chem). Each ion can be either monatomic, such as sodium (Na⁺) and chloride (Cl⁻) in sodium chloride, or polyatomic, such as ammonium (Template:Chem) and carbonate (Template:Chem) ions in ammonium carbonate. Salts containing basic ions hydroxide (OH⁻) or oxide (O²⁻) are classified as bases, such as sodium hydroxide and potassium oxide.

Individual ions within a salt usually have multiple near neighbours, so they are not considered to be part of molecules, but instead part of a continuous three-dimensional network. Salts usually form crystalline structures when solid.

Salts composed of small ions 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 become mobile. Some salts have large cations, large anions, or both. In terms of their properties, such species often are more similar to organic compounds.

History of discovery

File:X-ray spectrometer, 1912. (9660569929).jpg

X-ray spectrometer developed by W. H. Bragg

In 1913 the structure of sodium chloride was determined by William Henry Bragg and his son William Lawrence Bragg.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref><ref name=sherman/> 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.<ref name=sherman/> Many other inorganic compounds were also found to have similar structural features.<ref name=sherman/> 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.<ref name=sherman>Template:Cite journal</ref><ref>Template:Cite journal</ref>

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.Template:Sfn Born predicted crystal energies based on the assumption of ionic constituents, which showed good correspondence to thermochemical measurements, further supporting the assumption.<ref name=sherman/>

Formation

File:Halite-57430.jpg

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

File:Lead(II) sulfate.jpg

Solid lead(II) sulfate (PbSO₄)

Many metals such as the alkali metals react directly with the electronegative halogens gases to form salts.Template:SfnTemplate:Sfn

Salts form upon evaporation of their solutions.Template:Sfn Once the solution is supersaturated and the solid compound nucleates.Template:Sfn This process occurs widely in nature and is the means of formation of the evaporite minerals.<ref>Template:Cite book</ref>

Insoluble salts 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 salt, it is important to ensure they do not also precipitate.Template:Sfn 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.Template:Sfn Alternately the counterions can be chosen to ensure that even when combined into a single solution they will remain soluble as spectator ions.Template:Sfn

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 compared to the anhydrous material.Template:Sfn

Molten salts will solidify on cooling to below their freezing point.Template:Sfn This is sometimes used for the solid-state synthesis of complex salts from solid reactants, which are first melted together.Template:Sfn 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 neighboring reactants can diffuse together during the time the reactant mixture remains in the oven.Template:Sfn Other synthetic routes use a solid precursor with the correct stoichiometric ratio of non-volatile ions, which is heated to drive off other species.Template:Sfn

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,Template:Sfn a process thermodynamically understood using the Born–Haber cycle.Template:Sfn

Salts are formed by salt-forming reactions

• A base and an acid, e.g., NH₃ + HCl → NH₄Cl
• A metal and an acid, e.g., Mg + H₂SO₄ → MgSO₄ + H₂
• A metal and a non-metal, e.g., Ca + Cl₂ → CaCl₂
• A base and an acid anhydride, e.g., 2 NaOH + Cl₂O → 2 NaClO + H₂O
• An acid and a base anhydride, e.g., 2 HNO₃ + Na₂O → 2 NaNO₃ + H₂O
• In the salt metathesis reaction where two different salts are mixed in water, their ions recombine, and the new salt is insoluble and precipitates. For example:

  Pb(NO₃)₂ + Na₂SO₄ → PbSO₄↓ + 2 NaNO₃

Bonding

File:NaF.gif

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.

{{#invoke:Labelled list hatnote|labelledList|Main article|Main articles|Main page|Main pages}} Ions in salts 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.Template:Sfn 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.Template:Sfn 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,Template:Sfn due to the Pauli exclusion principle.Template:Sfn The balance between these forces leads to a potential energy well with minimum energy when the nuclei are separated by a specific equilibrium distance.Template:Sfn

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.Template:Sfn Salts 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.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref> Conversely, covalent bonds between unlike atoms often exhibit some charge separation and can be considered to have a partial ionic character.Template:Sfn 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.<ref>Template:Cite book</ref> 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.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref> 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).Template:Sfn

Although chemists classify idealized bond types as being ionic or covalent, the existence of additional types such as hydrogen bonds and metallic bonds, for example, has led some philosophers of science to suggest that alternative approaches to understanding bonding are required. This could be by applying quantum mechanics to calculate binding energies.<ref>Template:Cite journal</ref><ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>

Structure

File:Mercury-telluride-unit-cell-3D-ionic.png

The unit cell of the zinc blende structure

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 neighboring distance by a multiplicative constant called the Madelung constantTemplate:Sfn that can be efficiently computed using an Ewald sum.Template:Sfn When a reasonable form is assumed for the additional repulsive energy, the total lattice energy can be modelled using the Born–Landé equation,Template:Sfn the Born–Mayer equation, or in the absence of structural information, the Kapustinskii equation.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>

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.Template:SfnTemplate:Sfn Depending on the stoichiometry of the salt, and the coordination (principally determined by the radius ratio) of cations and anions, a variety of structures are commonly observed,<ref name=Moore>Template:Cite book</ref> and theoretically rationalized by Pauling's rules.Template:Sfn

Common ionic compound structures with close-packed anions<ref name=Moore/>

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	0.4142Template:Sfn	sodium chloride	1.747565Template:Sfn	nickeline	<1.73Template:Efn<ref>Template:Cite journal</ref>
	4:4	alternate tetrahedral	0.2247Template:Sfn	zinc blende	1.6381Template:Sfn	wurtzite	1.641<ref name="sherman" />
MX2	8:4	all tetrahedral	0.2247	fluorite	5.03878<ref name=Dienes>Template:Cite book</ref>
	6:3	half octahedral (alternate layers fully occupied)	0.4142	cadmium chloride	5.61<ref>Template:Cite journal</ref>	cadmium iodide	4.71<ref name=Dienes/>
MX3	6:2	one-third octahedral	0.4142	rhodium(III) bromideTemplate:Efn<ref>{{#invoke:citation/CS1|citation	CitationClass=web }}</ref><ref name=Ellis/> || 6.67<ref name=Hoppe1966/>Template:Efn || bismuth iodide || 8.26<ref name=Hoppe1966>Template:Cite journal</ref>Template:Efn
M2X3	6:4	two-thirds octahedral	0.4142			corundum	25.0312<ref name=Dienes/>
ABO3		two-thirds octahedral	0.4142			ilmenite	Depends on charges and structureTemplate:Efn
AB2O4		one-eighth tetrahedral and one-half octahedral	rA/rO = 0.2247, rB/rO = 0.4142Template:Efn	spinel, inverse spinel	Depends on cation site distributions<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref><ref>Template:Cite magazine</ref>	olivine	Depends on cation site distributions<ref>Template:Cite journal</ref>

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<ref name=Ellis>Template:Cite book</ref>

Stoichiometry	Cation:anion coordination	Interstitial sites occupied	Example structure
			Name	Critical radius ratio	Madelung constant
MX	8:8	entirely filled	cesium chloride	0.7321Template:Sfn	1.762675Template:Sfn
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).<ref name=":0">Template:Cite journal</ref>

Defects

Template:See also Within any crystal, there will usually be some defects. To maintain electroneutrality of the crystals, defects that involve loss of a cation will be associated with loss of an anion, i.e. these defects come in pairs.<ref name=":3">Template:Cite journal</ref> Frenkel defects consist of a cation vacancy paired with a cation interstitial and can be generated anywhere in the bulk of the crystal,<ref name=":3" /> occurring most commonly in compounds with a low coordination number and cations that are much smaller than the anions.<ref name=Prakash/> Schottky defects consist of one vacancy of each type, and are generated at the surfaces of a crystal,<ref name=":3" /> occurring most commonly in compounds with a high coordination number and when the anions and cations are of similar size.<ref name=Prakash>Template:Cite book</ref> 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.<ref name=":3" /> Another non-stoichiometric possibility is the formation of an F-center, a free electron occupying an anion vacancy.Template:Sfn When the compound has three or more ionic components, even more defect types are possible.<ref name=":3" /> 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.<ref name=":3" /> 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.<ref name=":3" /> 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.<ref name=":3" /> This results in a different splitting of d-electron orbitals, so that the optical absorption (and hence colour) can change with defect concentration.<ref name=":3" />

Properties

File:ILfromOS.svg

[BMIM]+[PF6]−, an ionic liquid

Acidity/basicity

Ionic compounds containing hydrogen ions (H⁺) are classified as acids, and those containing electropositive cations<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> and basic anions ions hydroxide (OH⁻) or oxide (O²⁻) are classified as bases. Other ionic compounds are known as salts and can be formed by acid–base reactions.<ref>Template:Cite book</ref> Salts that produce hydroxide ions when dissolved in water are called alkali salts, and salts that produce hydrogen ions when dissolved in water are called acid salts, e.g. sodium hydrogen selenite (Template:Chem2).<ref>Template:Greenwood&Earnshaw2nd</ref> If the compound is the result of a reaction between a strong acid and a weak base, the result is an acid salt. If it is the result of a reaction between a strong base and a weak acid, the result is a base salt. If it is the result of a reaction between a strong acid and a strong base, the result is a neutral salt.Template:Cn 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.<ref>Template:Cite journal</ref> 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.<ref>Template:Cite book</ref>

Solubility

File:SolubilityVsTemperature.png

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

Template:See also When simple salts dissolve, they dissociate into individual ions, which are solvated and dispersed throughout the resulting solution. Salts do not exist in solution.Template:Sfn In contrast, molecular compounds, which includes most organic compounds, remain intact in solution.

The solubility of salts is highest in polar solvents (such as water) or ionic liquids, but tends to be low in nonpolar solvents (such as petrol/gasoline).Template:Sfn This contrast 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 salts, which means that their solubility increases when the temperature increases.Template:Sfn There are some unusual salts 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.Template:Sfn

The lattice energy, the cohesive forces between these ions within a solid, determines the solubility. The solubility is dependent on how well each ion interacts with the solvent, so certain patterns become apparent. For example, salts of sodium, potassium and ammonium are usually soluble in water. Notable exceptions include ammonium hexachloroplatinate and potassium cobaltinitrite. Most nitrates and many sulfates are water-soluble. Exceptions include barium sulfate, calcium sulfate (sparingly soluble), and lead(II) sulfate, where the 2+/2− pairing leads to high lattice energies. For similar reasons, most metal carbonates are not soluble in water. Some soluble carbonate salts are: sodium carbonate, potassium carbonate and ammonium carbonate.

Strength

Strong salts or strong electrolyte salts are chemical salts composed of strong electrolytes. These salts dissociate completely or almost completely in water. They are generally odorless and nonvolatile.

Strong salts start with Na__, K__, NH₄__, or they end with __NO₃, __ClO₄, or __CH₃COO. Most group 1 and 2 metals form strong salts. Strong salts are especially useful when creating conductive compounds as their constituent ions allow for greater conductivity.Template:Citation needed

Weak salts or weak electrolyte salts are composed of weak electrolytes. These salts do not dissociate well in water. They are generally more volatile than strong salts. They may be similar in odor to the acid or base they are derived from. For example, sodium acetate, CH₃COONa, smells similar to acetic acid CH₃COOH.

Electrical conductivity

File:SegStackEdgeOnHMTFCQ.jpg

Edge-on view of portion of crystal structure of hexamethyleneTTF/TCNQ charge transfer salt.<ref>Template:Cite journal</ref>

Salts are characteristically insulators. Although they 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 salts are dissolved in a liquid or are melted into a liquid, they can conduct electricity because the ions become completely mobile. For this reason, molten salts and solutions containing dissolved salts (e.g., sodium chloride in water) can be used as electrolytes.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref> This conductivity gain upon dissolving or melting is sometimes used as a defining characteristic of salts.Template:Sfn

In some unusual salts: fast-ion conductors, and ionic glasses,<ref name=":0" /> one or more of the ionic components has a significant mobility, allowing conductivity even while the material as a whole remains solid.<ref name=":4">Template:Cite book</ref> This is often highly temperature dependent, and may be the result of either a phase change or a high defect concentration.<ref name=":4" /> These materials are used in all solid-state supercapacitors, batteries, and fuel cells, and in various kinds of chemical sensors.<ref>Template:Cite journal</ref><ref>Template:Cite journal</ref>

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.Template:Sfn Trends in melting points can be even better explained when the structure and ionic size ratio is taken into account.<ref>Template:Cite journal</ref> Above their melting point, salts melt and become molten salts (although some salts such as aluminium chloride and iron(III) chloride show molecule-like structures in the liquid phase).<ref>Template:Cite book</ref> 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.Template:Sfn 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.Template:Sfn

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.<ref name=":1">Template:Cite journal</ref> This means that even room temperature ionic liquids have low vapour pressures, and require substantially higher temperatures to boil.<ref name=":1" /> Boiling points exhibit similar trends to melting points in terms of the size of ions and strength of other interactions.<ref name=":1" /> When vapourized, the ions are still not freed of one another. For example, in the vapour phase sodium chloride exists as diatomic "molecules".<ref>Template:Cite book</ref>

Brittleness

Most salts 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.<ref name=":2">Template:Cite magazine</ref> As the temperature is elevated (usually close to the melting point) a ductile–brittle transition occurs, and plastic flow becomes possible by the motion of dislocations.<ref name=":2" /><ref>Template:Cite magazine</ref>

Compressibility

The compressibility of a salt 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.<ref>Template:Cite journal</ref>

Colour

The colour of a salt is often different from the colour of an aqueous solution containing the constituent ions,Template:Sfn or the hydrated form of the same compound.Template:Sfn

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

The absorption band of simple cations shifts toward a shorter wavelength when they are involved in more covalent interactions.Template:Sfn This occurs during hydration of metal ions, so colorless anhydrous salts with an anion absorbing in the infrared can become colorful in solution.Template:Sfn

Salts exist in many different colors, which arise either from their constituent anions, cations or solvates. For example:

• sodium chromate Template:Chem2 is made yellow by the chromate ion Template:Chem2.
• potassium dichromate Template:Chem2 is made red-orange by the dichromate ion Template:Chem2.
• cobalt(II) nitrate hexahydrate Template:Chem2 is made red by the chromophore of hydrated cobalt(II) Template:Chem2.
• copper(II) sulfate pentahydrate Template:Chem2 is made blue by the hydrated copper(II) cation.
• potassium permanganate Template:Chem2 is made violet by the permanganate anion Template:Chem2.
• nickel(II) chloride hexahydrate Template:Chem2 is made green by the hydrated nickel(II) chloride Template:Chem2.
• sodium chloride NaCl and magnesium sulfate heptahydrate Template:Chem2 are colorless or white because the constituent cations and anions do not absorb light in the part of the spectrum that is visible to humans.

Some minerals are salts, some of which are soluble in water.{{ safesubst:#invoke:Unsubst||date=__DATE__ |$B= Template:Fix }}Template:Clarify Similarly, inorganic pigments tend not to be salts, because insolubility is required for fastness. Some organic dyes are salts, but they are virtually insoluble in water.

Taste

Salts can elicit all five basic tastes, e.g., salty (sodium chloride), sweet (lead diacetate, which will cause lead poisoning if ingested), sour (potassium bitartrate), bitter (magnesium sulfate), and umami or savory (monosodium glutamate).

Odor

Salts of strong acids and strong bases ("strong salts") are non-volatile and often odorless, whereas salts of either weak acids or weak bases ("weak salts") may smell like the conjugate acid (e.g., acetates like acetic acid (vinegar) and cyanides like hydrogen cyanide (almonds)) or the conjugate base (e.g., ammonium salts like ammonia) of the component ions. That slow, partial decomposition is usually accelerated by the presence of water, since hydrolysis is the other half of the reversible reaction equation of formation of weak salts.

Uses

Salts have long had a wide variety of uses and applications. Many minerals are ionic.Template:Sfn 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.<ref>Template:Cite book</ref> Many salts 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.<ref>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>

Soluble salts 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.Template:Sfn Because the solutes are charged ions they also increase the electrical conductivity of the solution.Template:Sfn The increased ionic strength reduces the thickness of the electrical double layer around colloidal particles, and therefore the stability of emulsions and suspensions.Template:Sfn

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.<ref name=Reeves>{{#invoke:citation/CS1|citation |CitationClass=web }}</ref>

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

In chemical synthesis, salts are often used as precursors for high-temperature solid-state synthesis.<ref>Template:Cite book</ref>

Many metals are geologically most abundant as salts within ores.Template:Sfn 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.Template:Sfn<ref>Template:Cite book</ref>

Nomenclature

Template:See also According to the nomenclature recommended by IUPAC, salts are named according to their composition, not their structure.Template:Sfn In the most simple case of a binary salt with no possible ambiguity about the charges and thus the stoichiometry, the common name is written using two words.Template:Sfn The name of the cation (the unmodified element name for monatomic cations) comes first, followed by the name of the anion.Template:Sfn<ref name=Kotz>Template:Cite book</ref> For example, MgCl₂ is named magnesium chloride, and Na₂SO₄ is named sodium sulfate (Template:Chem, 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.Template:Sfn

If there are multiple different cations and/or anions, multiplicative prefixes (di-, tri-, tetra-, ...) are often required to indicate the relative compositions,Template:Sfn and cations then anions are listed in alphabetical order.Template:Sfn For example, KMgCl₃ is named magnesium potassium trichloride to distinguish it from K₂MgCl₄, magnesium dipotassium tetrachloride<ref>Template:Cite journal</ref> (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).Template:Sfn When one of the ions already has a multiplicative prefix within its name, the alternate multiplicative prefixes (bis-, tris-, tetrakis-, ...) are used.Template:Sfn For example, Ba(BrF₄)₂ is named barium bis(tetrafluoridobromate).Template:Sfn

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.Template:Sfn Because of the risk of ambiguity in allocating oxidation states, IUPAC prefers direct indication of the ionic charge numbers.Template:Sfn 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).Template:Sfn For example, FeSO₄ is named iron(2+) sulfate (with the 2+ charge on the Fe²⁺ 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).Template:Sfn 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.Template:Sfn 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, Template:Chem, has uranium in an oxidation state of +6, so would be called a dioxouranium(VI) ion in Stock nomenclature.<ref>Template:Cite journal</ref> 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.Template:Sfn For example, this scheme uses "ferrous" and "ferric", for iron(II) and iron(III) respectively,Template:Sfn so the examples given above were classically named ferrous sulfate and ferric sulfate.Template:Citation needed

Common salt-forming cations include:

• Ammonium Template:Chem
• Calcium Template:Chem
• Iron Template:Chem and Template:Chem
• Magnesium Template:Chem
• Potassium Template:Chem
• Pyridinium Template:Chem
• Quaternary ammonium Template:Chem, R being an alkyl group or an aryl group
• Sodium Template:Chem
• Copper Template:Chem

Common salt-forming anions (parent acids in parentheses where available) include:

• Acetate Template:Chem (acetic acid)
• Carbonate Template:Chem (carbonic acid)
• Chloride Template:Chem (hydrochloric acid)
• Citrate Template:Chem (citric acid)
• Cyanide Template:Chem (hydrocyanic acid)
• Fluoride Template:Chem (hydrofluoric acid)
• Nitrate Template:Chem (nitric acid)
• Nitrite Template:Chem (nitrous acid)
• Oxide Template:Chem (water)
• Phosphate Template:Chem (phosphoric acid)
• Sulfate Template:Chem (sulfuric acid)

Salts with varying number of hydrogen atoms replaced by cations as compared to their parent acid can be referred to as monobasic, dibasic, or tribasic, identifying that one, two, or three hydrogen atoms have been replaced; polybasic salts refer to those with more than one hydrogen atom replaced. Examples include:

• Sodium phosphate monobasic (NaH₂PO₄)
• Sodium phosphate dibasic (Na₂HPO₄)
• Sodium phosphate tribasic (Na₃PO₄)

Non-salt

Zwitterion

Zwitterions contain an anionic and a cationic centre in the same molecule, but are not considered salts. Examples of zwitterions are amino acids, many metabolites, peptides, and proteins.<ref>Template:Cite book</ref>

See also

• Bonding in solids
• Ioliomics
• Salt metathesis reaction
• Bresle method (the method used to test for salt presence during coating applications)
• Carboxylate
• Halide
• Ionic bonds
• Natron
• Salinity

References

• Mark Kurlansky (2002). Salt: A World History. Walker Publishing Company. Template:ISBN.

Bibliography

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