Salt (chemistry): Difference between revisions

[[Image:NaCl bonds.svg|thumb|The [[crystal]] structure of [[sodium chloride]], NaCl, a typical ionic compoundsalt. The purple spheres represent [[sodium]] [[cation]]s, Na⁺, and the green spheres represent [[chloride]] [[anion]]s, Cl⁻. The yellow stipples show the electrostatic forces.]]

The component ions in a salt can be either [[inorganic compound|inorganic]], such as [[chloride]] (Cl⁻), or [[organic chemistry|organic]], such as [[acetate]] ({{chem|CH|3|COO|−}}). Each ion can be either [[monatomic ion|monatomic]] (termed [[simple ion]]), such as [[fluoride]] (F⁻), and [[sodium]] (Na⁺) and [[chloride]] (Cl⁻) in [[sodium chloride]], or [[polyatomic ion|polyatomic]], such as [[sulfate]] ({{chem|SO|4|2−}}), and [[ammonium]] ({{chem|NH|4|+}}) and [[carbonate]] ({{chem|CO|3|2−}}) ions in [[ammonium carbonate]]. SaltSalts containing basic ions [[hydroxide]] (OH⁻) or [[oxide]] (O²⁻) are classified as [[Base (chemistry)|bases]], for example [[sodium hydroxide]].

Insoluble ionic compoundssalts 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 [[counterion]]s of the opposite charges. To ensure that these do not contaminate the precipitated ionic compoundsalt, it is important to ensure they do not also precipitate.{{sfn|Zumdahl|1989|p=133–140}} 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#Arrhenius theory|acid–base reaction]] or a [[neutralization reaction]] to form water.{{sfn|Zumdahl|1989|p=144–145}} Alternately the counterions can be chosen to ensure that even when combined into a single solution they will remain soluble as [[spectator ions]].{{sfn|Zumdahl|1989|p=133–140}}

Molten salts will solidify on cooling to below their [[freezing point]].{{sfn|Wold|Dwight|1993|page=79}} This is sometimes used for the [[Solid-state chemistry|solid-state synthesis]] of complex ionic compoundssalts from solid reactants, which are first melted together.{{sfn|Wold|Dwight|1993|pages=79–81}} 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.{{sfn|Wold|Dwight|1993|page=71}} Other synthetic routes use a solid precursor with the correct [[Stoichiometry|stoichiometric]] ratio of non-volatile ions, which is heated to drive off other species.{{sfn|Wold|Dwight|1993|page=71}}

* A [[base (chemistry)|base]] and an [[acid]], e.g., [[ammonia|NH₃]] + [[hydrochloric acid|HCl]] → [[ammonium chloride|NH₄Cl]]

* A [[metal]] and an [[acid]], e.g., [[magnesium|Mg]] + [[sulfuric acid|H₂SO₄]] → [[magnesium sulfate|MgSO₄]] + [[hydrogen|H₂]]

* A metal and a non-metal, e.g., [[calcium|Ca]] + [[chlorine|Cl₂]] → [[calcium chloride|CaCl₂]]

* A [[base (chemistry)|base]] and an [[acid anhydride]], e.g., 2 [[Sodium Hydroxide|NaOH]] + [[Dichlorine monoxide|Cl₂O]] → 2 [[Sodium hypochlorite|NaClO]] + [[Water|H₂O]]

* An [[acid]] and a [[base anhydride]], e.g., 2 [[nitric acid|HNO₃]] + [[Sodium oxide|Na₂O]] → 2 [[Sodium nitrate|NaNO₃]] + [[Water|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:

Ions in ionic compoundssalts are primarily held together by the [[electrostatic force]]s between the charge distribution of these bodies, and in particular, the ionic bond resulting from the long-ranged [[Coulomb's law|Coulomb]] attraction between the net negative charge of the anions and net positive charge of the cations.{{sfn|Pauling|1960|p=6}} 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.{{sfn|Kittel|2005|p=61}} When a pair of ions comes close enough for their [[valence shell|outer]] [[electron shell]]s (most simple ions have [[closed shell]]s) to overlap, a short-ranged repulsive force occurs,{{sfn|Pauling|1960|p=507}} due to the [[Pauli exclusion principle]].{{sfn|Ashcroft|Mermin|1977|p=379}} The balance between these forces leads to a potential energy well with minimum energy when the nuclei are separated by a specific equilibrium distance.{{sfn|Pauling|1960|p=507}}

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.{{sfn|Pauling|1960|p=65}} Ionic compoundsSalts 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|covalency]].<ref>{{cite journal|last1=Hannay|first1=N. Bruce|last2=Smyth|first2=Charles P.|title=The Dipole Moment of Hydrogen Fluoride and the Ionic Character of Bonds|journal=Journal of the American Chemical Society|date=February 1946|volume=68|issue=2|pages=171–173|doi=10.1021/ja01206a003}}</ref><ref>{{cite journal|last1=Pauling|first1=Linus|title=The modern theory of valency|journal=Journal of the Chemical Society (Resumed)|date=1948|volume=17|pages=1461–1467|doi=10.1039/JR9480001461|pmid=18893624|url=https://authors.library.caltech.edu/59671/|access-date=2021-12-01|archive-date=2021-12-07|archive-url=https://web.archive.org/web/20211207153730/https://authors.library.caltech.edu/59671/|url-status=dead}}</ref> Conversely, covalent bonds between unlike atoms often exhibit some charge separation and can be considered to have a partial ionic character.{{sfn|Pauling|1960|p=65}} 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>{{cite book|first1=John. N.|last1=Lalena|first2=David. A.|last2=Cleary|title=Principles of inorganic materials design|date=2010|publisher=John Wiley|location=Hoboken, N.J|isbn=978-0-470-56753-1|edition=2nd}}</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>{{cite journal|last1=Pearson|first1=Ralph G.|title=Hard and Soft Acids and Bases|journal=Journal of the American Chemical Society|date=November 1963|volume=85|issue=22|pages=3533–3539|doi=10.1021/ja00905a001}}</ref><ref>{{cite journal|last1=Pearson|first1=Ralph G.|title=Hard and soft acids and bases, HSAB, part II: Underlying theories|journal=Journal of Chemical Education|date=October 1968|volume=45|issue=10|page=643|doi=10.1021/ed045p643|bibcode=1968JChEd..45..643P}}</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).{{sfn|Barrow|1988|p=676}}

Using an even simpler approximation of the ions as impenetrable hard spheres, the arrangement of anions in these systems are often related to [[Close-packing of equal spheres|close-packed]] arrangements of spheres, with the cations occupying tetrahedral or octahedral [[interstitial site|interstice]]s.{{sfn|Ashcroft|Mermin|1977|p=383}}{{sfn|Zumdahl|1989|p=444–445}} Depending on the [[stoichiometry]] of the ionic compoundsalt, and the [[Coordination sphere|coordination]] (principally determined by the [[Cation-anion radius ratio|radius ratio]]) of cations and anions, a variety of structures are commonly observed,<ref name=Moore>{{cite book|last1=Moore|first1=Lesley E. Smart; Elaine A.|title=Solid state chemistry: an introduction|date=2005|publisher=Taylor & Francis, CRC|location=Boca Raton, Fla. [u.a.]|isbn=978-0-7487-7516-3|page=44|edition=3.}}</ref> and theoretically rationalized by [[Pauling's rules]].{{sfn|Ashcroft|Mermin|1977|pp=382–387}}

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 [[ionic glass]] is formed (with no long-range order).<ref name=":0">{{cite journal|last1=Souquet|first1=J|title=Electrochemical properties of ionically conductive glasses|journal=Solid State Ionics|date=October 1981|volume=5|pages=77–82|doi=10.1016/0167-2738(81)90198-3}}</ref>

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 point|melting]] and [[boiling point]]s and a low [[Vapor pressure|vapour pressure]].{{sfn|McQuarrie|Rock|1991|p = 503}} Trends in melting points can be even better explained when the structure and ionic size ratio is taken into account.<ref>{{Cite journal|title = The Influence of Relative Ionic Sizes on the Properties of Ionic Compounds|journal = Journal of the American Chemical Society|date = 1928-04-01|issn = 0002-7863|pages = 1036–1045|volume = 50|issue = 4|doi = 10.1021/ja01391a014|first = Linus|last = Pauling}}</ref> Above their melting point, ionic solidssalts melt and become [[molten salt]]s (although some ionic compoundssalts such as [[aluminium chloride]] and [[iron(III) chloride]] show molecule-like structures in the liquid phase).<ref>{{cite book|last1=Tosi|first1=M. P.|editor1-last=Gaune-Escard|editor1-first=Marcelle|title=Molten Salts: From Fundamentals to Applications|date=2002|publisher=Springer Netherlands|location=Dordrecht|isbn=978-94-010-0458-9|page=1|url=https://books.google.com/books?id=ft9sCQAAQBAJ&pg=PA1|url-status=live|archive-url=https://web.archive.org/web/20171203204320/https://books.google.com/books?id=ft9sCQAAQBAJ&lpg=PA11&pg=PA1|archive-date=2017-12-03}}</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&nbsp;°C), and are termed [[ionic liquid]]s.{{sfn|Freemantle|2009|p=1}} Ions in ionic liquids often have uneven charge distributions, or bulky [[substituent]]s like hydrocarbon chains, which also play a role in determining the strength of the interactions and propensity to melt.{{sfn|Freemantle|2009|pages=3–4}}

Most ionic compoundssalts are very [[brittle]]. Once they reach the limit of their strength, they cannot deform [[malleability|malleably]], because the strict alignment of positive and negative ions must be maintained. Instead the material undergoes [[fracture]] via [[cleavage (crystal)|cleavage]].<ref name=":2">{{cite magazine|last1=Johnston|first1=T. L.|last2=Stokes|first2=R. J.|last3=Li|first3=C. H.|title=The ductile–brittle transition in ionic solids|magazine=Philosophical Magazine|date=December 1959|volume=4|issue=48|pages=1316–1324|doi=10.1080/14786435908233367|bibcode=1959PMag....4.1316J}}</ref> As the temperature is elevated (usually close to the melting point) a [[Ductile-brittle transition temperature|ductile–brittle transition]] occurs, and [[plastic flow]] becomes possible by the motion of [[dislocation]]s.<ref name=":2" /><ref>{{Cite magazine|title = Ductile and brittle crystals|magazine=Philosophical Magazine|date = 1967-03-01|issn = 0031-8086|pages = 567–586|volume = 15|issue = 135|doi = 10.1080/14786436708220903|first1 = A.|last1 = Kelly|first2 = W. R.|last2 = Tyson|first3 = A. H.|last3 = Cottrell|bibcode = 1967PMag...15..567K}}</ref>

The [[compressibility]] of ana ionic compoundsalt 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>{{cite journal|last1=Stillwell|first1=Charles W.|title=Crystal chemistry. V. The properties of binary compounds|journal=Journal of Chemical Education|date=January 1937|volume=14|issue=1|page=34|doi=10.1021/ed014p34|bibcode=1937JChEd..14...34S}}</ref>

[[File:SolubilityVsTemperature.png|thumb|right|317px|The aqueous solubility of a variety of ionic compoundssalts as a function of temperature. Some compounds exhibiting unusual solubility behavior have been included.]]

The [[solubility]] of salts is highest in [[polar solvent]]s (such as [[water]]) or [[ionic liquid]]s, but tends to be low in [[nonpolar solvent]]s (such as [[petrol]]/[[gasoline]]).{{sfn|Brown|2009|pages=413–415}} This contrast is principally because the resulting [[Intermolecular force#Ion–dipole and ion–induced dipole forces|ion–dipole interactions]] are significantly stronger than ion-induced dipole interactions, so the [[enthalpy change of solution|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 of mixing|entropy change of solution]] is usually positive for most solid solutes like ionic compoundssalts, which means that their solubility increases when the temperature increases.{{sfn|Brown|2009|p = 422}} There are some unusual ionic compoundssalts 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.{{sfn|Brown|2009|p = 422}}

Salts are characteristically [[Insulator (electricity)|insulators]]. Although they contain charged atoms or clusters, these materials do not typically [[electrical conductivity|conduct electricity]] to any significant extent when the substance is solid. In order to conduct, the charged particles must be [[Electrical mobility|mobile]] rather than stationary in a [[Crystal structure|crystal lattice]]. This is achieved to some degree at high temperatures when the defect concentration increases the ionic mobility and [[solid state ionic conductivity]] is observed. When the ionic compoundssalts are [[Solution (chemistry)|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 [[electrolyte]]s.<ref>{{cite web|title=Electrical Conductivity of Ionic Compound|url=http://cikguwong.blogspot.com/2011/05/chemistry-form-4-chapter-5-electrical.html|access-date=2 December 2012|url-status=live|archive-url=https://web.archive.org/web/20140521205809/http://cikguwong.blogspot.com/2011/05/chemistry-form-4-chapter-5-electrical.html|archive-date=21 May 2014|date=2011-05-22}}</ref> This conductivity gain upon dissolving or melting is sometimes used as a defining characteristic of ionic compoundssalts.{{sfn|Zumdahl|1989|p=341}}

In some unusual ionic compoundssalts: [[fast -ion conductor]]s, and [[ionic glassesglass]]es,<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">{{Cite book|title = An Introduction to Electronic and Ionic Materials|last1 = Gao|first1 = Wei|publisher = World Scientific|year = 1999|isbn = 978-981-02-3473-7|page = 261|url = https://books.google.com/books?id=fxH3N_7L0LwC&pg=PA261|last2 = Sammes|first2 = Nigel M|url-status = live|archive-url = https://web.archive.org/web/20171203204320/https://books.google.com/books?id=fxH3N_7L0LwC&lpg=PR7&ots=MR0Sj2c4x9&pg=PA261#v=onepage&f=false|archive-date = 2017-12-03}}</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 [[supercapacitor]]s, [[battery (electricity)|batteries]], and [[fuel cell]]s, and in various kinds of [[chemical sensor]]s.<ref>{{cite journal|last1=West|first1=Anthony R.|title=Solid electrolytes and mixed ionic?electronic conductors: an applications overview|journal=Journal of Materials Chemistry|date=1991|volume=1|issue=2|page=157|doi=10.1039/JM9910100157}}</ref><ref>{{cite journal|last1=Boivin|first1=J. C.|last2=Mairesse|first2=G.|title=Recent Material Developments in Fast Oxide Ion Conductors|journal=Chemistry of Materials|date=October 1998|volume=10|issue=10|pages=2870–2888|doi=10.1021/cm980236q}}</ref>

{{see also|ColorColour of chemicals}}

The [[colourColor of chemicals#saltsSalts|colour of an ionica compoundsalt]] is often different from the [[colour of chemicals#ions in aqueous solution|colour of an aqueous solution]] containing the constituent ions,{{Sfn|Pauling|1960|p=105}} or the [[hydrate]]d form of the same compound.{{sfn|Brown|2009|page=417}}

The absorption band of simple cations shifts toward a shorter wavelength when they are involved in more covalent interactions.{{Sfn|Pauling|1960|p=107}} This occurs during [[solvation|hydration]] of metal ions, so colorless [[anhydrous]] ionic compoundssalts with an anion absorbing in the infrared can become colorful in solution.{{Sfn|Pauling|1960|p=107}}

Ionic compoundsSalts have long had a wide variety of uses and applications. Many [[minerals]] are ionic.{{sfn|Wenk|Bulakh|2004|page=774}} 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>{{cite book|last1=Kurlansky|first1=Mark|title=Salt: a world history|date=2003|publisher=Vintage|location=London|isbn=978-0-09-928199-3|edition=1st }}</ref> Many ionic compoundssalts 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>{{cite web|last1=Lower|first1=Simon|title=Naming Chemical Substances|url=http://www.chem1.com/acad/webtext/intro/int-5.html|website=Chem₁ General Chemistry Virtual Textbook|access-date=14 January 2016|date=2014|url-status=live|archive-url=https://web.archive.org/web/20160116000437/http://www.chem1.com/acad/webtext/intro/int-5.html|archive-date=16 January 2016}}</ref>

Soluble ionic compounds like saltsalts 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]].{{sfn|Atkins|de Paula|2006|pages=150–157}} Because the solutes are charged ions they also increase the electrical conductivity of the solution.{{sfn|Atkins|de Paula|2006|pages=761–770}} The increased ionic strength reduces the thickness of the [[electrical double layer]] around [[colloid]]al particles, and therefore the stability of [[emulsion]]s and [[Suspension (chemistry)|suspensions]].{{sfn|Atkins|de Paula|2006|pages=163–169}}

Solid ionic compoundssalts have long been used as paint pigments, and are resistant to organic solvents, but are sensitive to acidity or basicity.<ref>{{cite book|last1=Satake|first1=M|last2=Mido|first2=Y|title=Chemistry of Colour|date=1995|publisher=Discovery Publishing House|isbn=978-81-7141-276-1|page=230|url=https://books.google.com/books?id=FA4hOk5KJBgC&pg=PA230|url-status=live|archive-url=https://web.archive.org/web/20171203204320/https://books.google.com/books?id=FA4hOk5KJBgC&lpg=PA230&ots=4wpC5lAywl&pg=PA230|archive-date=2017-12-03}}</ref> Since 1801 [[pyrotechnician]]s have described and widely used metal-containing ionic compoundssalts as sources of colour in fireworks.{{sfn|Russell|2009|page=14}} Under intense heat, the electrons in the metal ions or small molecules can be excited.{{sfn|Russell|2009|page=82}} These electrons later return to lower energy states, and release light with a colour spectrum characteristic of the species present.{{sfn|Russell|2009|pages=108–117}}{{sfn|Russell|2009|pages=129–133}}

In chemistry[[chemical synthesis]], ionic compoundssalts are often used as [[Precursor (chemistry)|precursors]] for high-temperature solid-state synthesis.<ref>{{cite book|last1=Xu|first1=Ruren|first2=Wenqin|last2=Pang|first3=Qisheng|last3=Huo|title=Modern inorganic synthetic chemistry|url=https://archive.org/details/moderninorganics00xuru|url-access=limited|date=2011|publisher=Elsevier|location=Amsterdam|isbn=978-0-444-53599-3|page=[https://archive.org/details/moderninorganics00xuru/page/n27 22]}}</ref>

Many metals are geologically most abundant as ionic compoundssalts within [[ore]]s.{{sfn|Zumdahl|Zumdahl|2015|pages=822}} To obtain the [[Chemical element|elemental]] materials, these ores are processed by [[smelting]] or [[electrolysis]], in which [[redox reaction]]s occur (often with a reducing agent such as carbon) such that the metal ions gain electrons to become neutral atoms.{{sfn|Zumdahl|Zumdahl|2015|pages=823}}<ref>{{cite book|last1=Gupta|first1=Chiranjib Kumar|title=Chemical metallurgy principles and practice|url=https://archive.org/details/chemicalmetallur00gupt|url-access=limited|date=2003|publisher=Wiley-VCH|location=Weinheim|isbn=978-3-527-60525-5|pages=[https://archive.org/details/chemicalmetallur00gupt/page/n376 359]–365}}</ref>

According to the [[nomenclature]] recommended by [[IUPAC]], ionic compoundssalts are named according to their composition, not their structure.{{sfn|IUPAC|2005|p=68}} In the most simple case of a binary ionic compoundsalt with no possible ambiguity about the charges and thus the [[stoichiometry]], the common name is written using two words.{{sfn|IUPAC|2005|p=70}} The name of the cation (the unmodified element name for monatomic cations) comes first, followed by the name of the anion.{{sfn|IUPAC|2005|p=69}}<ref name=Kotz>{{cite book |last1=Kotz |first1= John C.|last2= Treichel|first2= Paul M|last3 = Weaver|first3 = Gabriela C.|title= Chemistry and Chemical Reactivity|edition=Sixth|date= 2006|publisher= Thomson Brooks/Cole|location= Belmont, CA|isbn=978-0-534-99766-3 |page= 111}}</ref> For example, MgCl₂ is named [[magnesium chloride]], and Na₂SO₄ is named [[sodium sulfate]] ({{chem|SO|4|2−}}, [[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.{{sfn|Brown|2009|pp=36-37}}