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 complex polyatomic species such as the carbonate (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 network. Ionic compounds have high melting and boiling points, and they are hard and very brittle. As solids they are almost always electrically insulating, but when melted or dissolved they become highly conductive, because the ions are mobilized. For example, these traits are apparent in salts, which are a major class of ionic compounds.


27 Caesium Auride.gif
Caesium auride (CsAu), an ionic compound between two metals, shown dissolved in liquid ammonia and in solid crystal form.

Ions in ionic compounds are held together by the electrostatic forces between oppositely charged bodies. Ionic compounds are rarely purely ionic, i.e. held together only by electrostatic forces. The bonds between the most electronegative/electropositive pairs such as caesium fluoride exhibit a degree of covalency. Similarly, covalent compounds often exhibit charge separations. See also HSAB theory.

Ionic compounds have very strong electrostatic bonds between particles. As a result, they generally have very high melting and boiling points and a low vapour pressure.[1] They also have good electrical conductivity when molten or in an aqueous solution.[2] Ionic inorganic compounds typically have high melting points so are solids at room temperature and usually form crystals. Unlike organic compounds they do not char nor ignite. On the other hand organic compounds have low melting points, most of them are insoluble in water, and characteristically they ignite quite easily.[3]

The ions produced by electron transfer attract each other by electrostatic attraction and this creates an ionic bond.


Ions typically pack into extremely regular crystalline structures, in an arrangement that minimizes the Coulomb energy (maximizing attractions and minimizing repulsions). For spherical ions (including all simple ions), the arrangement of anions in these systems are often related to close-packed arrangements of spheres, with the cations occupying interstices. Depending on the stoichiometry of the ionic compound, and the coordination (principally determined by the size ratio) of cations and anions, a variety of structures are commonly observed.[4]

Common ionic compound structures with close-packed anions[4]
Stoichiometry Cation:anion coordination Interstitial sites occupied Cubic close packing Hexagonal close packing
MX 6:6 all octahedral sodium chloride nickel arsenide
4:4 alternate tetrahedral zinc blende wurtzite
MX2 8:4 all tetrahedral fluorite
6:3 half octahedral (alternate layers fully occupied) cadmium chloride cadmium iodide
MX3 6:2 one-third octahedral chromium(III) chloride[5] bismuth iodide
M2X3 6:4 two-thirds octahedral corundum
ABO3 two-thirds octahedral ilmenite
AB2O4 one-eighth tetrahedral and one-half octahedral spinel, inverse spinel olivine

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[5]
Stoichiometry Cation:anion coordination Interstitial sites occupied Example structure
MX 8:8 entirely filled cesium chloride
MX2 8:4 half filled calcium fluoride
M2X 4:8 half filled lithium oxide


Ionic compounds dissolve most readily in polar solvents (such as water) or ionic liquids, but tend to have a low solubility 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.

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


According to the nomenclature recommended by IUPAC, ionic compounds are named according to their composition, not their structure.[7] 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.[8] The name of the cation (the unmodified element name for monatomic cations) comes first, followed by the name of the anion.[9][10] 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,[11] and cations then anions are listed in alphabetical order.[12] 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).[13] When one of the ions already has a multiplicative prefix in its name, the alternate multiplicative prefixes (bis, tris, tetrakis, ...) are used.[14] For example, Ba(BrF4)2 is named barium bis(tetrafluoridobromate).[15]

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.[15] Because of the risk of ambiguity in allocating oxidation states, IUPAC prefers direct indication of the ionic charge numbers.[15] 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).[15] 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), and iron(3+) sulfate respectively.[15] 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.[16] The Classical naming system, no longer in common use, gave some ionic oxidation states special names, such as "ferrous" and "ferric", for iron(II) and iron(III) respectively, so the examples given above were classically named ferrous sulfate and ferric sulfate.

See also[edit]


  1. ^ McQuarrie & Rock 1991, p. 503.
  2. ^ Dutta, Priti. "What are the characteristics of Ionic Compunds". Retrieved 2 December 2012. 
  3. ^ "A comparison of organic and inorganic compounds". Retrieved 2 December 2012. 
  4. ^ 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. 
  5. ^ 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. 
  6. ^ "Electrical Conductivity of Ionic Compound". Retrieved 2 December 2012. 
  7. ^ IUPAC 2005, p. 68.
  8. ^ IUPAC 2005, p. 70.
  9. ^ IUPAC 2005, p. 69.
  10. ^ 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. 
  11. ^ IUPAC 2005, pp. 75-76.
  12. ^ IUPAC 2005, p. 75.
  13. ^ IUPAC 2005, p. 76.
  14. ^ IUPAC 2005, pp. 76-77.
  15. ^ a b c d e IUPAC 2005, p. 77.
  16. ^ IUPAC 2005, pp. 77-78.