Ionic compound

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The crystal structure of sodium chloride, NaCl, a typical ionic compound. The purple spheres represent sodium cations, Na+, and the green spheres represent chloride anions, Cl.

In chemistry, an ionic compound is a chemical compound in which ions are held together in a structure by electrostatic forces termed ionic bonds. The positively charged ions are called cations and the negatively charged ions are called anions. These can be simple ions such as the sodium (Na+) and chloride (Cl) in sodium chloride, or polyatomic species such as the carbonate ion (CO32−) in calcium carbonate. Individual ions usually have multiple nearest neighbours, so are not considered to be part of molecules, but instead part of a continuous three-dimensional network, usually in a crystalline structure.

A salt is a compound which contains a cation other than H+ and an anion other than hydroxide ion, OH-, or oxide ion, O2-.[1] Compounds containing hydrogen ions are classified as acids and compounds containing hydroxide or oxide ions are classified as bases. In general, salts can be formed by acid-base reactions.

Ionic compounds typically have high melting and boiling points, and are hard and brittle. As solids they are almost always electrically insulating, but when melted or dissolved they become highly conductive, because the ions are mobilized.

History of discovery[edit]

The word ion is the Greek ἰόν, ion, "going", the present participle of ἰέναι, ienai, "to go". This term was introduced by English physicist and chemist Michael Faraday in 1834 for the then-unknown species that goes from one electrode to the other through an aqueous medium.[2][3]

Principal contributors to the development of a theoretical treatment of ionic crystal structure were Max Born, Fritz Haber, Alfred Landé, Erwin Madelung, Paul Peter Ewald, and Kazimierz Fajans.[4]


Ionic compounds can be produced by evaporation, precipitation, or freezing. If the solvent from an electrolyte solution is evaporated, the ions do not go into the vapour, but stay in the remaining solution, and ultimately crystallize into the ionic compound. Insoluble ionic compounds can be precipitated by mixing two solutions, one with the cation and one with the anion in it. Molten salts will solidify on cooling to below their freezing point.


Ions in ionic compounds are primarily held together by the electrostatic forces between the charge distribution of these bodies, and in particular the ionic bond resulting from the Coulomb attraction between the net charge on anions and cations.[5] When a pair of ions comes close enough for their outer electron shells (most simple ions have closed shells) to overlap, a repulsive force occurs,[6] due to the Pauli exclusion principle.[7] 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.[8] The balance between these forces leads to a potential energy well with a minimum energy when the nuclei are separated by a specific equilibrium distance.[6]

If the electronic structure of the two interacting bodies is affected by the presence of one another, covalent interactions (non-ionic) also contribute.[9] Ionic compounds are rarely purely ionic, i.e. held together only by electrostatic forces. The bonds between even the most electronegative/electropositive pairs such as those in caesium fluoride exhibit a small degree of covalency.[10][11] Conversely, covalent bonds between unlike atoms often exhibit some charge separation and can be considered to have a partial ionic character.[9] The circumstances under which a compound will have ionic or covalent character can typically be understood using HSAB theory, 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.[12][13] This difference in electronegativities means that the charge separation is maintained even when the ions are in contact (the excess electrons on the anions are not transferred to neutralize the cations).


Ions typically pack into extremely regular crystalline structures, in an arrangement that minimizes the lattice energy (maximizing attractions and minimizing repulsions). The lattice energy is the summation of the interaction of all sites with all other sites. For spherical ions (including all simple ions) only the charges and distances are required to determine the interaction energy. For any particular crystal structure, all distances are related to the smallest internuclear distance. So for each possible crystal structure, the total Coulomb energy can be related to the Coulomb energy of unit charges at the nearest neighbour distance by a multiplicative contant called the Madelung constant[6] that can be efficiently computed using an Ewald sum.[14] When a reasonable form is assumed for the additional repulsive energy, the total lattice energy can be modelled using the Born–Landé equation.[15]

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 interstices.[16] 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,[17] and theoretically rationalized by Pauling's rules.[18]

Common ionic compound structures with close-packed anions[17]
Stoichiometry Cation:anion coordination Interstitial sites occupied Cubic close packing of anions Hexagonal close packing of anions
name critical radius ratio Madelung constant name critical radius ratio Madelung constant
MX 6:6 all octahedral sodium chloride 0.4142[16] 1.747565[19] nickel arsenide
4:4 alternate tetrahedral zinc blende 0.2247[20] 1.6381[19] wurtzite 0.2247[20]
MX2 8:4 all tetrahedral fluorite 5.03878[21]
6:3 half octahedral (alternate layers fully occupied) cadmium chloride 5.61[22] cadmium iodide 4.71[21]
MX3 6:2 one-third octahedral chromium(III) chloride[23] bismuth iodide
M2X3 6:4 two-thirds octahedral corundum 25.0312[21]
ABO3 two-thirds octahedral ilmenite
AB2O4 one-eighth tetrahedral and one-half octahedral spinel, inverse spinel depends on cation site distributions[24][25][26] olivine depends on cation site distributions[27]

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[23]
Stoichiometry Cation:anion coordination Interstitial sites occupied Example structure
name critical radius ratio Madelung constant
MX 8:8 entirely filled cesium chloride 0.7321[28] 1.762675[19]
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).[29]


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, they generally have very high melting and boiling points and a low vapour pressure.[30] Inorganic compounds with simple ions typically have small ions, and thus have high melting points, so are solids at room temperature.

At high temperatures ionic solids melt and become molten salts. Some substances with larger ions, however, do form ionic liquids at room temperature.


Ionic compounds are typically very brittle. Once they reach the limit of their strength, they cannot deform mealleably, because the strict alignment of positive and negative ions must be maintained. Instead the material undergoes brittle fracture.


When ionic compounds dissolve, the individual ions dissociate and are solvated by the solvent and dispersed throughout the resulting solution.[31] The solubility is highest in polar solvents (such as water) or ionic liquids, but tends to be low in nonpolar solvents (such as petrol/gasoline). This 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 energy released in solvation is used to overcome the lattice energy so that the ions are freed from their positions in the crystal and dissolve in the liquid. Because the ions are released into solution when dissolved, soluble salts are a common class of electrolytes.

Electrical conductivity[edit]

Although ionic compounds contain charged atoms or clusters, these materials do not typically conduct electricity when the substance is solid. In order to conduct, the charged particles must be mobile rather than stationary in a crystal lattice. When the ionic compounds are dissolved in a liquid or are themselves melted into a liquid, they can conduct electricity because the ions become mobile.[32]

In some unusual ionic compounds: fast ion conductors, and ionic glasses,[29] one or more of the ionic components has a significant mobility, allowing conductivity even while the material as a whole remains solid.


The colour of an ionic compound is often different to the colour of an aqueous solution containing the constituent ions.[33]

The anions in compounds with bonds with the most ionic character tend to be colourless (with an absorption band in the ultraviolet part of the spectrum).[34] In compounds with less ionic character, their colour deepens through yellow, orange, red and black (as the absorption band shifts to longer wavelengths into the visible spectrum).[34]

The absorption band of simple cations shift toward shorter wavelength when they are involved in more covalent interactions.[34] This occurs during hydration of metal ions, so colourless ionic compounds with anion absorbing in the infrared can become colourful in solution.[34]


Ionic compounds have long had a wide variety of uses and applications. Many minerals are ionic. Humans have processed sodium chloride for over 8000 years, using it first as a food seasoning and preservative, and now also in manufacturing, agriculture, water conditioning, and for de-icing roads.


According to the nomenclature recommended by IUPAC, ionic compounds are named according to their composition, not their structure.[35] 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.[36] The name of the cation (the unmodified element name for monatomic cations) comes first, followed by the name of the anion.[37][38] For example, MgCl2 is named magnesium chloride, and Na2SO4 is named sodium sulfate (SO42−, 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.

If there are multiple cations and/or anions, multiplicative prefixes (di, tri, tetra, ...) are often required to indicate the relative compositions,[39] and cations then anions are listed in alphabetical order.[40] For example, KMgCl3 is named magnesium potassium trichloride (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).[41] When one of the ions already has a multiplicative prefix in its name, the alternate multiplicative prefixes (bis, tris, tetrakis, ...) are used.[42] For example, Ba(BrF4)2 is named barium bis(tetrafluoridobromate).[43]

Compounds containing one or more elements which can exist in a variety of charge/oxidation states will have a stoichiometry that depends on which oxidation states are present, to ensure overall neutrality. This can be indicated in the name by specifying either the oxidation state of the elements present, or the charge on the ions.[43] Because of the risk of ambiguity in allocating oxidation states, IUPAC prefers direct indication of the ionic charge numbers.[43] These are written as an arabic integer followed by the sign (..., 2−, 1−, 1+, 2+, ...) in parentheses directly after the name of the cation (without a space separating them).[43] 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).[43] 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.[44] An even older naming system for metal cations, also still widely used, appended the suffixes "ous" and "ic" to the Latin roots the name, to give special names for the low and high oxidation states.[45] For example this scheme uses "ferrous" and "ferric", for iron(II) and iron(III) respectively,[45] so the examples given above were classically named ferrous sulfate and ferric sulfate.

See also[edit]


  1. ^ Whitten, Kenneth W.; Galley, Kenneth D.; Davis, Raymond E. (1992). General Chemistry (4th ed.). Saunders. p. 128. ISBN 0-03-072373-6. 
  2. ^ Michael Faraday (1791-1867). UK: BBC. 
  3. ^ "Online etymology dictionary". Retrieved 2011-01-07. 
  4. ^ Pauling 1960, p. 505.
  5. ^ Pauling 1960, p. 6.
  6. ^ a b c Pauling 1960, p. 507.
  7. ^ Ashcroft & Mermin 1977, p. 379.
  8. ^ Kittel 2005, p. 61.
  9. ^ a b Pauling 1960, p. 65.
  10. ^ Hannay, N. Bruce; Smyth, Charles P. (February 1946). "The Dipole Moment of Hydrogen Fluoride and the Ionic Character of Bonds". Journal of the American Chemical Society 68 (2): 171–173. doi:10.1021/ja01206a003. 
  11. ^ Pauling, Linus (1948). "The modern theory of valency". Journal of the Chemical Society (Resumed): 1461–1467. doi:10.1039/JR9480001461. 
  12. ^ Pearson, Ralph G. (November 1963). "Hard and Soft Acids and Bases". Journal of the American Chemical Society 85 (22): 3533–3539. doi:10.1021/ja00905a001. 
  13. ^ Pearson, Ralph G. (October 1968). "Hard and soft acids and bases, HSAB, part II: Underlying theories". Journal of Chemical Education 45 (10): 643. doi:10.1021/ed045p643. 
  14. ^ Kittel 2005, p. 64.
  15. ^ Pauling 1960, p. 509.
  16. ^ a b Ashcroft & Mermin 1977, p. 383.
  17. ^ a b Moore, Lesley E. Smart; Elaine A. (2005). Solid state chemistry : an introduction (3. ed.). Boca Raton, Fla. [u.a.]: Taylor & Francis, CRC. p. 44. ISBN 9780748775163. 
  18. ^ Ashcroft & Mermin 1977, pp. 382-387.
  19. ^ a b c Kittel 2005, p. 65.
  20. ^ a b Ashcroft & Mermin 1977, p. 386.
  21. ^ a b c Dienes, Richard J. Borg, G.J. (1992). The physical chemistry of solids. Boston: Academic Press. p. 123. ISBN 9780121184209. 
  22. ^ Brackett, Thomas E.; Brackett, Elizabeth B. (1965). "The Lattice Energies of the Alkaline Earth Halides". Journal of Physical Chemistry 69 (10): 3611–3614. doi:10.1021/j100894a062. 
  23. ^ a b Ellis, Arthur B. [] et al. (1995). Teaching general chemistry : a materials science companion (3. print ed.). Washington: American Chemical Soceity. p. 121. ISBN 084122725X. 
  24. ^ Verwey, E. J. W. (1947). "Physical Properties and Cation Arrangement of Oxides with Spinel Structures I. Cation Arrangement in Spinels". Journal of Chemical Physics 15 (4): 174–180. doi:10.1063/1.1746464. 
  25. ^ Verwey, E. J. W.; de Boer, F.; van Santen, J. H. (1948). "Cation Arrangement in Spinels". The Journal of Chemical Physics 16 (12): 1091. doi:10.1063/1.1746736. 
  26. ^ Thompson, P.; Grimes, N. W. (27 September 2006). "Madelung calculations for the spinel structure". Philosophical Magazine 36 (3): 501–505. doi:10.1080/14786437708239734. 
  27. ^ Alberti, A.; Vezzalini, G. (1978). "Madelung energies and cation distributions in olivine-type structures". Zeitschrift für Kristallographie - Crystalline Materials 147 (1–4): 167–176. doi:10.1524/zkri.1978.147.14.167. 
  28. ^ Ashcroft & Mermin 1977, p. 384.
  29. ^ a b Souquet, J (October 1981). "Electrochemical properties of ionically conductive glasses". Solid State Ionics 5: 77–82. doi:10.1016/0167-2738(81)90198-3. 
  30. ^ McQuarrie & Rock 1991, p. 503.
  31. ^ Brown 2009, pp. 89-91.
  32. ^ "Electrical Conductivity of Ionic Compound". Retrieved 2 December 2012. 
  33. ^ Pauling 1960, p. 105.
  34. ^ a b c d Pauling 1960, p. 107.
  35. ^ IUPAC 2005, p. 68.
  36. ^ IUPAC 2005, p. 70.
  37. ^ IUPAC 2005, p. 69.
  38. ^ Kotz, John C.; Treichel, Paul M; Weaver, Gabriela C. (2006). Chemistry and Chemical Reactivity (Sixth ed.). Belmont, CA: Thomson Brooks/Cole. p. 111. ISBN 0-534-99766-X. 
  39. ^ IUPAC 2005, pp. 75-76.
  40. ^ IUPAC 2005, p. 75.
  41. ^ IUPAC 2005, p. 76.
  42. ^ IUPAC 2005, pp. 76-77.
  43. ^ a b c d e IUPAC 2005, p. 77.
  44. ^ IUPAC 2005, pp. 77-78.
  45. ^ a b Brown 2009, p. 38.


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