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In chemistry, a salt or ionic compound is a chemical compound consisting of an assembly of positively charged ions (Cation) and negatively charged ions (Anion), which results in a compound with no net electric charge (electrically neutral). The constituent ions are held together by electrostatic forces termed ionic bonding.
The component ions in a salt can be either inorganic, such as chloride (Cl−), or organic, such as acetate (). Each ion can be either monatomic ion, such as sodium (Na+) and chloride (Cl−) in sodium chloride, or polyatomic ion, such as ammonium () and carbonate () 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 point and , and are Hardness and Brittleness. As solids they are almost always electrically insulating, but when melting 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.
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 thermochemistry measurements, further supporting the assumption.
Salts can form upon evaporation of solvent from their solutions once the solution is supersaturation and the solid compound nucleates. This process occurs widely in nature and is the means of formation of the evaporite minerals.
Insoluble salts can be precipitated by mixing two solutions, one containing the cation and one containing the anion. Because all solutions are electrically neutral, the two solutions mixed must also contain of the opposite charges. To ensure that these do not contaminate the precipitated salt, 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 compared to the anhydrous material.
Molten salts will solidify on cooling to below their freezing point. This is sometimes used for the solid-state synthesis of complex salts 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 neighboring reactants can diffuse together during the time the reactant mixture remains in the oven. Other synthetic routes use a solid precursor with the correct Stoichiometry ratio of non-volatile ions, which is heated to drive off other species.
In some reactions between highly reactive metals (usually from Alkali metal 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.
Salts can be formed through a variety of reaction types, such as those between:
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. 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 covalent bond. 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.
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.
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 salt, and the coordination (principally determined by the radius ratio) of cations and anions, a variety of structures are commonly observed,
| + Common ionic compound structures with close-packed anions ! rowspan=2 | Stoichiometry ! rowspan=2 | Cation:anion coordination ! colspan="2" | Interstitial sites ! colspan="2" | Cubic close packing of anions ! colspan="2" | Hexagonal close packing of anions |
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| 1.641 | |||||
| 4.71 | |||||
| 8.26 | |||||
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| Depends on charges and structure | |||||
| 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 (1995). 9780841227255, American Chemical Society. ISBN 9780841227255
!rowspan=2 | Stoichiometry !rowspan=2 | Cation:anion coordination !rowspan=2 | Interstitial sites occupied ! colspan=3 | Example structure |
| 1.762675 | ||||
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).
Some ions are classed as Amphoterism, 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.
The solubility of salts is highest in (such as water) or , but tends to be low in (such as petrol/gasoline). 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. 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.
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 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 are not soluble in water. Some soluble carbonate salts are: sodium carbonate, potassium carbonate and ammonium carbonate.
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.
Weak salts or weak electrolyte salts are composed of weak . 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.
In some unusual salts: fast-ion conductors, and , 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 dependent, and may be the result of either a phase change or a high defect concentration. These materials are used in all solid-state , batteries, and , and in various kinds of .
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".
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). 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).
The absorption band of simple cations shifts toward a shorter wavelength when they are involved in more covalent interactions. This occurs during solvation of metal ions, so colorless anhydrous salts with an anion absorbing in the infrared can become colorful in solution.
Salts exist in many different , which arise either from their constituent anions, cations or Solvation. For example:
Some minerals are salts, some of which are soluble in water. Similarly, inorganic tend not to be salts, because insolubility is required for fastness. Some organic are salts, but they are virtually insoluble in water.
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. 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 particles, and therefore the stability of 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 salts have long been used as paint pigments, and are resistant to organic solvents, but are sensitive to acidity or basicity.
Since 1801 have described and widely used metal-containing salts 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 chemical synthesis, salts are often used as precursors for high-temperature solid-state synthesis.
Many metals are geologically most abundant as salts within . To obtain the Chemical element materials, these ores are processed by smelting or electrolysis, in which occur (often with a reducing agent such as carbon) such that the metal ions gain electrons to become neutral atoms.
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/ 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 numerals 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 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). 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, , 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.
Common salt-forming cations include:
Common salt-forming anions (parent acids in parentheses where available) include:
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:
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