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Salt (chemistry)

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"Ionic compound" redirects here. Not to be confused withSaltorSodium chloride.
Thecrystalstructure ofsodium chloride,NaCl, a typical salt. The purple spheres representsodiumcations,Na+,and the green spheres representchlorideanions,Cl−.The yellow stipples show the electrostatic forces.

Inchemistry,asaltorionic compoundis achemical compoundconsisting of an assembly of positively chargedions(cations) and negatively charged ions (anions),[1]which results in a compound with no netelectric charge(electrically neutral). The constituent ions are held together byelectrostatic forcestermedionic bonds.

The component ions in a salt can be eitherinorganic,such aschloride(Cl−), ororganic,such asacetate(CH
3COO−
). Each ion can be eithermonatomic(termedsimple ion), such asfluoride(F−), andsodium(Na+) andchloride(Cl−) insodium chloride,orpolyatomic,such assulfate(SO2−
4), andammonium(NH+
4) andcarbonate(CO2−
3) ions inammonium carbonate.Salt containing basic ionshydroxide(OH−) oroxide(O2−) are classified asbases,for examplesodium hydroxide.

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 formcrystalline structureswhen solid.

Salts composed of small ions typically have highmeltingandboiling points,and arehardandbrittle.As solids they are almost alwayselectrically insulating,but whenmeltedordissolvedthey become highlyconductive,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[edit]

X-ray spectrometer developed by W. H. Bragg

In 1913 the structure ofsodium chloridewas determined byWilliam Henry BraggandWilliam Lawrence Bragg.[2][3][4]This revealed that there were six equidistantnearest-neighboursfor each atom, demonstrating that the constituents were not arranged in molecules or finite aggregates, but instead as a network with long-rangecrystallineorder.[4]Many otherinorganic compoundswere also found to have similar structural features.[4]These compounds were soon described as being constituted of ions rather than neutralatoms,but proof of this hypothesis was not found until the mid-1920s, whenX-ray reflectionexperiments (which detect the density of electrons), were performed.[4][5]

Principal contributors to the development of atheoreticaltreatment of ionic crystal structures wereMax Born,Fritz Haber,Alfred Landé,Erwin Madelung,Paul Peter Ewald,andKazimierz Fajans.[6]Born predicted crystal energies based on the assumption of ionic constituents, which showed good correspondence tothermochemicalmeasurements, further supporting the assumption.[4]

Formation[edit]

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

Many metals such as thealkali metalsreact directly with theelectronegativehalogensgases to salts.[7][8]

Salts form upon evaporation of theirsolutions.[9]Once the solution issupersaturatedand the solid compound nucleates.[9]This process occurs widely in nature and is the means of formation of theevaporiteminerals.[10]

Insoluble salts can be precipitated by mi xing 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 containcounterionsof the opposite charges. To ensure that these do not contaminate the precipitated salt, it is important to ensure they do not also precipitate.[11]If the two solutions have hydrogen ions and hydroxide ions as the counterions, they will react with one another in what is called anacid–base reactionor aneutralization reactionto form water.[12]Alternately the counterions can be chosen to ensure that even when combined into a single solution they will remain soluble asspectator ions.[11]

If the solvent is water in either the evaporation or precipitation method of formation, in many cases theionic crystalformed also includeswater of crystallization,so the product is known as ahydrate,and can have very different chemical properties compared to theanhydrousmaterial.[13]

Molten salts will solidify on cooling to below theirfreezing point.[14]This is sometimes used for thesolid-state synthesisof complex salts 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 asolid-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 correctstoichiometricratio of non-volatile ions, which is heated to drive off other species.[8]

In some reactions between highly reactive metals (usually fromGroup 1orGroup 2) and highly electronegative halogen gases, or water, the atoms can be ionized byelectron transfer,[16]a process thermodynamically understood using theBorn–Haber cycle.[17]

Salts are formed bysalt-forming reactions

Bonding[edit]

A schematicelectron shelldiagram ofsodiumandfluorineatoms undergoing a redox reaction to formsodium fluoride.Sodium loses its outerelectronto give it a stableelectron configuration,and this electron enters the fluorine atomexothermically.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 salts are primarily held together by theelectrostatic forcesbetween the charge distribution of these bodies, and in particular, the ionic bond resulting from the long-rangedCoulombattraction between the net negative charge of the anions and net positive charge of the cations.[18]There is also a small additional attractive force fromvan der Waals interactionswhich contributes only around 1–2% of the cohesive energy for small ions.[19]When a pair of ions comes close enough for theirouterelectron shells(most simple ions haveclosed shells) to overlap, a short-ranged repulsive force occurs,[20]due to thePauli 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 theelectronic structureof 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.[22]Salts are rarely purely ionic, i.e. held together only by electrostatic forces. The bonds between even the mostelectronegative/electropositivepairs such as those incaesium fluorideexhibit a small degree ofcovalency.[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 usingFajans' 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 generallyHSAB theorycan 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 ashydrogen bondsandmetallic bonds,for example, has led some philosophers of science to suggest that alternative approaches to understanding bonding are required. This could be by applyingquantum mechanicsto calculate binding energies.[29][30]

Structure[edit]

The unit cell of thezinc blendestructure

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 theMadelung constant[20]that can be efficiently computed using anEwald sum.[31]When a reasonable form is assumed for the additional repulsive energy, the total lattice energy can be modelled using theBorn–Landé equation,[32]theBorn–Mayer equation,or in the absence of structural information, theKapustinskii equation.[33]

Using an even simpler approximation of the ions as impenetrable hard spheres, the arrangement of anions in these systems are often related toclose-packedarrangements of spheres, with the cations occupying tetrahedral or octahedralinterstices.[34][35]Depending on thestoichiometryof the salt, and thecoordination(principally determined by theradius ratio) of cations and anions, a variety of structures are commonly observed,[36]and theoretically rationalized byPauling'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

Someionic liquids,particularly with mixtures of anions or cations, can be cooled rapidly enough that there is not enough time for crystalnucleationto occur, so anionic glassis formed (with no long-range order).[53]

Defects[edit]

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.[54]Frenkel defectsconsist of a cation vacancy paired with a cation interstitial and can be generated anywhere in the bulk of the crystal,[54]occurring most commonly in compounds with a low coordination number and cations that are much smaller than the anions.[55]Schottky defectsconsist of one vacancy of each type, and are generated at the surfaces of a crystal,[54]occurring most commonly in compounds with a high coordination number and when the anions and cations are of similar size.[55]If the cations have multiple possibleoxidation states,then it is possible for cation vacancies to compensate for electron deficiencies on cation sites with higher oxidation numbers, resulting in anon-stoichiometric compound.[54]Another non-stoichiometric possibility is the formation of anF-center,a free electron occupying an anion vacancy.[56]When the compound has three or more ionic components, even more defect types are possible.[54]All of these point defects can be generated via thermal vibrations and have anequilibriumconcentration. Because they are energetically costly butentropicallybeneficial, 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 andsolid state ionic conductivity.[54]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 thelattice parameters,reducing the overall density of the crystal.[54]Defects also result in ions in distinctly different local environments, which causes them to experience a differentcrystal-field symmetry,especially in the case of different cations exchanging lattice sites.[54]This results in a differentsplittingofd-electron orbitals,so that the optical absorption (and hence colour) can change with defect concentration.[54]

Properties[edit]

[BMIM]+[PF6]−,anionic liquid

Acidity/basicity[edit]

Ionic compounds containinghydrogen ions(H+) are classified asacids,and those containingelectropositivecations[57]and basic anions ionshydroxide(OH−) oroxide(O2−) are classified asbases.Other ionic compounds are known as salts and can be formed byacid–base reactions.[58]Salts that producehydroxideionswhen dissolved inwaterare calledalkali salts,and salts that producehydrogenionswhen dissolved inwaterare calledacid salts.If the compound is the result of a reaction between astrong acidand aweak base,the result is anacid salt.If it is the result of a reaction between astrong baseand aweak acid,the result is abase 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 theconjugate baseion and conjugate acid ion, such asammonium acetate.

Some ions are classed asamphoteric,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 aszinc oxide,aluminium hydroxide,aluminium oxideandlead(II) oxide.[60]

Melting and boiling points[edit]

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 highmeltingandboiling pointsand a lowvapour 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, salts melt and becomemolten salts(although some salts such asaluminium chlorideandiron(III) chlorideshow 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 termedionic liquids.[64]Ions in ionic liquids often have uneven charge distributions, or bulkysubstituentslike 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[edit]

Most salts are verybrittle.Once they reach the limit of their strength, they cannot deformmalleably,because the strict alignment of positive and negative ions must be maintained. Instead the material undergoesfractureviacleavage.[68]As the temperature is elevated (usually close to the melting point) aductile–brittle transitionoccurs, andplastic flowbecomes possible by the motion ofdislocations.[68][69]

Compressibility[edit]

Thecompressibilityof an salt is strongly determined by its structure, and in particular thecoordination 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[edit]

The aqueous solubility of a variety of salts 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 saltsdissolve,theydissociateinto individual ions, which aresolvatedand 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.

Thesolubilityof salts is highest inpolar solvents(such aswater) orionic liquids,but tends to be low innonpolar solvents(such aspetrol/gasoline).[72]This contrast is principally because the resultingion–dipole interactionsare significantly stronger than ion-induced dipole interactions, so theheat of solutionis 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 thesolvationenergy exceeds thelattice energy,the negative netenthalpy change of solutionprovides a thermodynamic drive to remove ions from their positions in the crystal and dissolve in the liquid. In addition, theentropy change of solutionis usually positive for most solid solutes like salts, which means that their solubility increases when the temperature increases.[73]There are some unusual salts such ascerium(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]

Thelattice 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 ofsodium,potassiumand ammonium are usually soluble in water. Notable exceptions includeammonium hexachloroplatinateandpotassium cobaltinitrite.Mostnitratesand manysulfatesare water-soluble. Exceptions includebarium sulfate,calcium sulfate(sparingly soluble), andlead(II) sulfate,where the 2+/2− pairing leads to high lattice energies. For similar reasons, most metalcarbonatesare not soluble in water. Some soluble carbonate salts are:sodium carbonate,potassium carbonateandammonium carbonate.

Electrical conductivity[edit]

Edge-on view of portion of crystal structure of hexamethyleneTTF/TCNQcharge transfer salt.[74]

Salts are characteristicallyinsulators.Although they contain charged atoms or clusters, these materials do not typicallyconduct electricityto any significant extent when the substance is solid. In order to conduct, the charged particles must bemobilerather than stationary in acrystal lattice.This is achieved to some degree at high temperatures when the defect concentration increases the ionic mobility andsolid state ionic conductivityis observed. When the salts aredissolved in a liquidor are melted into aliquid,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 aselectrolytes.[75]This conductivity gain upon dissolving or melting is sometimes used as a defining characteristic of salts.[76]

In some unusual salts:fast ion conductors,andionic glasses,[53]one or more of the ionic components has a significant mobility, allowing conductivity even while the material as a whole remains solid.[77]This is often highly temperature dependent, and may be the result of either a phase change or a high defect concentration.[77]These materials are used in all solid-statesupercapacitors,batteries,andfuel cells,and in various kinds ofchemical sensors.[78][79]

Colour[edit]

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

Thecolour of a saltis often different from thecolour of an aqueous solutioncontaining the constituent ions,[80]or thehydratedform of the same compound.[13]

The anions in compounds with bonds with the most ionic character tend to be colorless (with anabsorption bandin 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 duringhydrationof metal ions, so colorlessanhydroussalts with an anion absorbing in the infrared can become colorful in solution.[81]

Salts exist in many differentcolors,which arise either from their constituent anions, cations orsolvates.For example:

  • sodium chromateNa2CrO4is made yellow by thechromate ionCrO2−4.
  • potassium dichromateK2Cr2O7is made red-orange by thedichromate ionCr2O2−7.
  • cobalt(II) nitratehexahydrateCo(NO3)2·6H2Ois made red by the chromophore ofhydratedcobalt(II)[Co(H2O)6]2+.
  • copper(II) sulfatepentahydrateCuSO4·5H2Ois made blue by the hydrated copper(II) cation.
  • potassium permanganateKMnO4is made violet by thepermanganateanionMnO−4.
  • nickel(II) chloridehexahydrateNiCl2·6H2Ois made green by the hydrated nickel(II) chloride[NiCl2(H2O)4].
  • sodium chlorideNaCl andmagnesium sulfateheptahydrateMgSO4·7H2Oare colorless or white because the constituentcationsandanionsdo not absorb light in the part of the spectrum that is visible to humans.

Somemineralsare salts, some of which aresolublein water.[dubious–discuss][clarification needed]Similarly, inorganicpigmentstend not to be salts, because insolubility is required for fastness. Some organicdyesare salts, but they are virtually insoluble in water.

Taste and odor[edit]

Salts can elicit all fivebasic tastes,e.g.,salty(sodium chloride),sweet(lead diacetate,which will causelead poisoningif ingested),sour(potassium bitartrate),bitter(magnesium sulfate), andumamiorsavory(monosodium glutamate).

Salts of strong acids and strong bases ( "strong salts") are non-volatileand often odorless, whereas salts of either weak acids or weak bases ( "weak salts") may smell like theconjugate acid(e.g., acetates like acetic acid (vinegar) and cyanides likehydrogen cyanide(almonds)) or the conjugate base (e.g., ammonium salts likeammonia) of the component ions. That slow, partial decomposition is usually accelerated by the presence of water, sincehydrolysisis the other half of thereversible reactionequation of formation of weak salts.

Uses[edit]

Salts have long had a wide variety of uses and applications. Manymineralsare ionic.[82]Humans have processedcommon 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.[83]Many salts are so widely used in society that they go by common names unrelated to their chemical identity. Examples of this includeborax,calomel,milk of magnesia,muriatic acid,oil of vitriol,saltpeter,andslaked lime.[84]

Soluble salts can easily be dissolved to provideelectrolytesolutions. This is a simple way to control the concentration andionic strength.The concentration of solutes affects manycolligative properties,including increasing theosmotic pressure,and causingfreezing-point depressionandboiling-point elevation.[85]Because the solutes are charged ions they also increase the electrical conductivity of the solution.[86]The increased ionic strength reduces the thickness of theelectrical double layeraroundcolloidalparticles, and therefore the stability ofemulsionsandsuspensions.[87]

The chemical identity of the ions added is also important in many uses. For example,fluoridecontaining compounds are dissolved to supply fluoride ions forwater fluoridation.[88]

Solid salts have long been used as paint pigments, and are resistant to organic solvents, but are sensitive to acidity or basicity.[89]Since 1801pyrotechnicianshave described and widely used metal-containing salts 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]

Inchemical synthesis,salts are often used asprecursorsfor high-temperature solid-state synthesis.[94]

Many metals are geologically most abundant as salts withinores.[95]To obtain theelementalmaterials, these ores are processed bysmeltingorelectrolysis,in whichredox reactionsoccur (often with a reducing agent such as carbon) such that the metal ions gain electrons to become neutral atoms.[96][97]

Nomenclature[edit]

See also:IUPAC nomenclature of inorganic chemistry

According to thenomenclaturerecommended byIUPAC,salts are named according to their composition, not their structure.[98]In the most simple case of a binary salt with no possible ambiguity about the charges and thus thestoichiometry,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, MgCl2is namedmagnesium chloride,and Na2SO4is namedsodium sulfate(SO2−
4,sulfate,is an example of apolyatomic ion). To obtain theempirical formulafrom 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,[103]and cations then anions are listed in Alpha betical order.[104]For example, KMgCl3is namedmagnesium potassium trichlorideto distinguish it from K2MgCl4,magnesium dipotassium tetrachloride[105](note that in both the empirical formula and the written name, the cations appear in Alpha betical order, but the order varies between them because thesymbolforpotassiumis 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)2is namedbarium bis(tetrafluoridobromate).[108]

Compounds containing one or more elements which can exist in a variety of charge/oxidation stateswill 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.[108]Because of the risk of ambiguity in allocating oxidation states, IUPAC prefers direct indication of the ionic charge numbers.[108]These are written as anarabicinteger followed by the sign (..., 2−, 1−, 1+, 2+,...) in parentheses directly after the name of the cation (without a space separating them).[108]For example, FeSO4is namediron(2+) sulfate(with the 2+ charge on theFe2+ions balancing the 2− charge on the sulfate ion), whereas Fe2(SO4)3is namediron(3+) sulfate(because the two iron ions in eachformula uniteach have a charge of 3+, to balance the 2− on each of the three sulfate ions).[108]Stock nomenclature,still in common use, writes theoxidation numberinRoman numerals(..., −II, −I, 0, I, II,...). So the examples given above would be namediron(II) sulfateandiron(III) sulfaterespectively.[109]For simple ions the ionic charge and the oxidation number are identical, but for polyatomic ions they often differ. For example, theuranyl(2+)ion,UO2+
2,has uranium in an oxidation state of +6, so would be called a dioxouranium(VI) ion in Stock nomenclature.[110]An even older naming system for metal cations, also still widely used, appended the suffixes-ousand-icto theLatinroot of the name, to give special names for the low and high oxidation states.[111]For example, this scheme uses "ferrous" and "ferric", for iron(II) and iron(III) respectively,[111]so the examples given above were classically namedferrous sulfateandferric 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 asmonobasic,dibasic,ortribasic,identifying that one, two, or three hydrogen atoms have been replaced;polybasicsalts refer to those with more than one hydrogen atom replaced. Examples include:

Strength[edit]

Strong salts or strongelectrolytesalts are chemical salts composed ofstrong electrolytes.These salts dissociate completely or almost completely inwater.They are generally odorless andnonvolatile.

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 weakelectrolytes.These salts do not dissociate well in water. They are generally morevolatilethan strong salts. They may be similar inodorto theacidorbasethey are derived from. For example,sodium acetate,CH3COONa, smells similar toacetic acidCH3COOH.

Zwitterion[edit]

Zwitterionscontain an anionic and a cationic centre in the samemolecule,but are not considered salts. Examples of zwitterions areamino acids,manymetabolites,peptides,andproteins.[112]

See also[edit]

Notes[edit]

  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 asyttrium(III) chlorideandchromium(III) chloride,but both are now known as the RhBr3structure type.
  3. ^The reference lists this structure asMoCl3,which is now known as the RhBr3structure.
  4. ^The reference lists this structure asFeCl3,which is now known as the BiI3structure 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 asMgAl2O4the larger cation occupies the smaller tetrahedral site.[47]

References[edit]

  1. ^IUPAC,Compendium of Chemical Terminology,2nd ed. (the "Gold Book" ) (1997). Online corrected version: (2006–) "salt".doi:10.1351/goldbook.S05447
  2. ^Bragg, W. H.; Bragg, W. L. (1 July 1913). "The Reflection of X-rays by Crystals".Proceedings of the Royal Society A: Mathematical, Physical and Engineering Sciences.88(605): 428–438.Bibcode:1913RSPSA..88..428B.doi:10.1098/rspa.1913.0040.S2CID13112732.
  3. ^Bragg, W. H. (22 September 1913)."The Reflection of X-rays by Crystals. (II.)".Proceedings of the Royal Society A: Mathematical, Physical and Engineering Sciences.89(610): 246–248.Bibcode:1913RSPSA..89..246B.doi:10.1098/rspa.1913.0082.
  4. ^abcdefSherman, Jack (August 1932). "Crystal Energies of Ionic Compounds and Thermochemical Applications".Chemical Reviews.11(1): 93–170.doi:10.1021/cr60038a002.
  5. ^James, R. W.; Brindley, G. W. (1 November 1928)."A Quantitative Study of the Reflexion of X-Rays by Sylvine".Proceedings of the Royal Society A: Mathematical, Physical and Engineering Sciences.121(787): 155–171.Bibcode:1928RSPSA.121..155J.doi:10.1098/rspa.1928.0188.
  6. ^Pauling 1960,p. 505.
  7. ^Zumdahl 1989,p. 312.
  8. ^abcWold & Dwight 1993,p. 71.
  9. ^abWold & Dwight 1993,p. 82.
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Bibliography[edit]

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