Salt (chemistry)

Source: Wikipedia, the free encyclopedia.
(Redirected from
Ionic compound
)
"Ionic compound" redirects here. Not to be confused with Salt or Sodium chloride.
anions
, Cl−. The yellow stipples show the electrostatic forces.

In

anions,[1] which results in a neutral compound with no net electric charge. The constituent ions are held together by electrostatic forces termed ionic bonds
.

The component ions in a salt can be either

simple ion), such as fluoride (F−), and sodium (Na+) and chloride (Cl−) in sodium chloride, or polyatomic, such as sulfate (SO2−
4), and ammonium (NH+
4) and carbonate (CO2−
3) ions in ammonium carbonate. Salt containing basic ions hydroxide (OH−) or oxide (O2−) are classified as bases, for example sodium hydroxide
.

Individual ions within a salt usually have multiple near neighbours, so 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

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

X-ray spectrometer developed by W. H. Bragg

In 1913 the structure of sodium chloride was determined by

William Lawrence Bragg.[2][3][4] 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.[4] Many other inorganic compounds were also found to have similar structural features.[4] 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.[4][5]

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.[6] Born predicted crystal energies based on the assumption of ionic constituents, which showed good correspondence to thermochemical measurements, further supporting the assumption.[4]

Formation

White crystals form a mineral sample of halite, shown against a black background.
Halite, the mineral form of sodium chloride, forms when salty water evaporates leaving the ions behind.
Solid lead(II) sulfate (PbSO4)

Many metals such as the alkali metals react directly with the electronegative halogens gases to salts.[7][8]

Salts form upon evaporation of their solutions.[9] Once the solution is supersaturated and the solid compound nucleates.[9] This process occurs widely in nature and is the means of formation of the evaporite minerals.[10]

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

spectator ions.[11]

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.[13]

Molten salts will solidify on cooling to below their

freezing point.[14] This is sometimes used for the solid-state synthesis of complex ionic compounds from solid reactants, which are first melted together.[15] 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.[8] Other synthetic routes use a solid precursor with the correct stoichiometric ratio of non-volatile ions, which is heated to drive off other species.[8]

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,[16] a process thermodynamically understood using the Born–Haber cycle.[17]

Salts are formed by salt-forming reactions

Bonding

exothermically
. The oppositely charged ions – typically a great many of them – are then attracted to each other to form a solid.
Main article: Ionic bonding

Ions in ionic compounds are primarily held together by the

closed shells) to overlap, a short-ranged repulsive force occurs,[20] due to the Pauli exclusion principle.[21] The balance between these forces leads to a potential energy well with minimum energy when the nuclei are separated by a specific equilibrium distance.[20]

If the

electropositive pairs such as those in caesium fluoride exhibit a small degree of covalency.[23][24] Conversely, covalent bonds between unlike atoms often exhibit some charge separation and can be considered to have a partial ionic character.[22] 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.[25] 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.[26][27] 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).[28]

Although chemists classify idealized bond types as being ionic or covalent, the existence of additional types such as

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.[29][30]

Structure

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

Ewald sum.[31] When a reasonable form is assumed for the additional repulsive energy, the total lattice energy can be modelled using the Born–Landé equation,[32] the Born–Mayer equation, or in the absence of structural information, the Kapustinskii equation.[33]

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.[34][35] 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,[36] and theoretically rationalized by Pauling's rules.[37]

Common ionic compound structures with close-packed anions[36]
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[34] sodium chloride 1.747565[38] nickeline <1.73[a][39]
4:4 alternate tetrahedral 0.2247[40]
zinc blende
 1.6381[38] wurtzite 1.641[4]
MX2 8:4 all tetrahedral 0.2247 fluorite 5.03878[41]
6:3 half octahedral (alternate layers fully occupied) 0.4142 cadmium chloride 5.61[42] cadmium iodide 4.71[41]
MX3 6:2 one-third octahedral 0.4142 rhodium(III) bromide[b][43][44] 6.67[45][c]
bismuth iodide
 8.26[45][d]
M2X3 6:4 two-thirds octahedral 0.4142 corundum 25.0312[41]
ABO3 two-thirds octahedral 0.4142 ilmenite Depends on charges
and structure [e]
AB2O4 one-eighth tetrahedral and one-half octahedral rA/rO = 0.2247,
rB/rO = 0.4142[f]		 spinel, inverse spinel Depends on cation
site distributions[48][49][50]		 olivine Depends on cation
site distributions[51]

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[44]
Stoichiometry Cation:anion
coordination 		Interstitial sites occupied Example structure
Name Critical radius
ratio 		Madelung constant
MX 8:8 entirely filled
cesium chloride
 0.7321[52] 1.762675[38]
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).[53]

Defects

Diagram of charged ions with a positive ion out of place in the structure
Frenkel defect
Diagram of charged ions with a positive and negative missing from the structure
Schottky defect
See also: crystallographic defect

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.

d-electron orbitals, so that the optical absorption (and hence colour) can change with defect concentration.[54]

Properties

[BMIM]+[PF6]−, an ionic liquid

Acidity/basicity

Ionic compounds containing

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.[59] 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.[60]

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.[61] Trends in melting points can be even better explained when the structure and ionic size ratio is taken into account.[62] 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).[63] 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.[64] 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.[65]

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.[66] This means that even room temperature ionic liquids have low vapour pressures, and require substantially higher temperatures to boil.[66] Boiling points exhibit similar trends to melting points in terms of the size of ions and strength of other interactions.[66] When vapourized, the ions are still not freed of one another. For example, in the vapour phase sodium chloride exists as diatomic "molecules".[67]

Brittleness

Most ionic compounds are very

plastic flow becomes possible by the motion of dislocations.[68][69]

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.[70]

Solubility

The aqueous solubility of a variety of ionic compounds as a function of temperature. Some compounds exhibiting unusual solubility behavior have been included.
See also: Solubility § Solubility of ionic compounds in water

When simple salts

dissolve, they dissociate into individual ions, which are solvated and dispersed throughout the resulting solution. Salts do not exist in solution. [71]
In contrast, molecular compounds, which includes most organic compounds, remain intact in solution.

The

petrol/gasoline).[72] 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 ionic compounds, which means that their solubility increases when the temperature increases.[73] 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.[73]

The

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
.

Electrical conductivity

TCNQ charge transfer salt.[74]

Salts are characteristically

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. For this reason, liquified (molten) salts and solutions containing dissolved salts (e.g., sodium chloride in water) can be used as electrolytes.[75] This conductivity gain upon dissolving or melting is sometimes used as a defining characteristic of ionic compounds.[76]

In some unusual ionic compounds:

chemical sensors.[78][79]

Colour

blue powder on a watch glass
Anhydrous cobalt(II) chloride,
CoCl2
a pile of red granules on white paper
Cobalt(II) chloride hexahydrate,
CoCl2·6H2O
See also: Color of chemicals

The

colour of an aqueous solution containing the constituent ions,[80] or the hydrated form of the same compound.[13]

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).[81] 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). [81]

The absorption band of simple cations shifts toward a shorter wavelength when they are involved in more covalent interactions.[81] This occurs during hydration of metal ions, so colorless anhydrous ionic compounds with an anion absorbing in the infrared can become colorful in solution.[81]

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

  • chromate ion
    CrO2−4.
  • dichromate ion
    Cr2O2−7.
  • cobalt(II) nitrate hexahydrate Co(NO3)2·6H2O is made red by the chromophore of hydrated cobalt(II) [Co(H2O)6]2+.
  • copper(II) sulfate pentahydrate CuSO4·5H2O is made blue by the hydrated copper(II) cation.
  • potassium permanganate KMnO4 is made violet by the permanganate anion MnO−4.
  • nickel(II) chloride hexahydrate NiCl2·6H2O is made green by the hydrated nickel(II) chloride [NiCl2(H2O)4].
  • anions
    do not absorb light in the part of the spectrum that is visible to humans.

Some

soluble in water.[dubious – discuss][clarification needed] 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 and odor

Salts can elicit all five

bitter (magnesium sulfate), and umami or savory (monosodium glutamate
).

Salts of strong acids and strong bases ("

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

Ionic compounds have long had a wide variety of uses and applications. Many

slaked lime.[84]

Soluble ionic compounds like salt can easily be dissolved to provide

electrical double layer around colloidal particles, and therefore the stability of emulsions and suspensions.[87]

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.[88]

Solid ionic compounds have long been used as paint pigments, and are resistant to organic solvents, but are sensitive to acidity or basicity.[89] Since 1801 pyrotechnicians have described and widely used metal-containing ionic compounds as sources of colour in fireworks.[90] Under intense heat, the electrons in the metal ions or small molecules can be excited.[91] These electrons later return to lower energy states, and release light with a colour spectrum characteristic of the species present.[92][93]

In chemistry, ionic compounds are often used as precursors for high-temperature solid-state synthesis.[94]

Many metals are geologically most abundant as ionic compounds within

redox reactions occur (often with a reducing agent such as carbon) such that the metal ions gain electrons to become neutral atoms.[96][97]

Nomenclature

See also: IUPAC nomenclature of inorganic chemistry

According to the

IUPAC, ionic compounds are named according to their composition, not their structure.[98] 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.[99] The name of the cation (the unmodified element name for monatomic cations) comes first, followed by the name of the anion.[100][101] 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.[102]

If there are multiple different cations and/or anions, multiplicative prefixes (di-, tri-, tetra-, ...) are often required to indicate the relative compositions,

symbol for potassium is K).[106] When one of the ions already has a multiplicative prefix within its name, the alternate multiplicative prefixes (bis-, tris-, tetrakis-, ...) are used.[107] For example, Ba(BrF4)2 is named barium bis(tetrafluoridobromate).[108]

Compounds containing one or more elements which can exist in a variety of charge/

ferric sulfate.[citation needed
]

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:

Types of salt

Acidity and basicity

Salts can be classified in a variety of ways. Salts that produce

ions when dissolved in water are called acid salts. Neutral salts are those salts that are neither acidic nor alkaline. 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.[112]

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.[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.

See also

Notes

  1. ^ This structure type has a variable lattice parameter c/a ratio, and the exact Madelung constant depends on this.
  2. ^ This structure has been referred to in references as yttrium(III) chloride and chromium(III) chloride, but both are now known as the RhBr3 structure type.
  3. ^ The reference lists this structure as MoCl3, which is now known as the RhBr3 structure.
  4. ^ The reference lists this structure as FeCl3, which is now known as the BiI3 structure type.
  5. ^ This structure type can accommodate any charges on A and B that add up to six. When both are three the charge structure is equivalent to that of corrundum.[46] The structure also has a variable lattice parameter c/a ratio, and the exact Madelung constant depends on this.
  6. ^ However, in some cases such as MgAl2O4 the larger cation occupies the smaller tetrahedral site.[47]

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Bibliography

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