Inert pair effect

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The inert pair effect is the tendency of the electrons in the outermost atomic s orbital to remain unionized or unshared in compounds of post-transition metals. The term inert pair effect is often used in relation to the increasing stability of oxidation states that are two less than the group valency for the heavier elements of groups 13, 14, 15 and 16. The term "inert pair" was first proposed by Nevil Sidgwick in 1927.[1] The name suggests that the s electrons are more tightly bound to the nucleus and therefore more difficult to ionize.

For example, the p-block elements of the 4th, 5th and 6th period come after d-block elements, but the electrons present in the intervening d- (and f-) orbitals do not effectively shield the s-electrons of the valence shell. As a result, the inert pair of ns electrons remains more tightly held by the nucleus and hence participates less in bonding.


As an example in group 13 the +1 oxidation state of Tl is the most stable and TlIII compounds are comparatively rare. The stability of the +1 oxidation state increases in the following sequence:[2]

AlI < GaI < InI < TlI.

The same trend in stability is noted in groups 14, 15 and 16. As such the heaviest members of the groups, e.g. lead, bismuth and polonium are comparatively stable in oxidation states +2, +3, and +4 respectively.
The lower oxidation state in each of the elements in question has 2 valence electrons in s - orbitals. On the face of it, a simple explanation could be that the valence electrons in an s orbital are more tightly bound and are of lower energy than electrons in p orbitals and therefore less likely to be involved in bonding.[3] Unfortunately this explanation does not stand up. If the total ionization potentials (IP) (see below) of the 2 electrons in s orbitals (the 2nd + 3rd ionization potentials), are examined it can be seen that they increase in the sequence:

In < Al < Tl < Ga.
Ionization potentials for group 13 elements
IP Boron Aluminium Gallium Indium Thallium
1st 800.6 577.5 578.8 558.3 589.4
2nd 2427.1 1816.7 1979.3 1820.6 1971
3rd 3659.7 2744.8 2963 2704 2878
(2nd + 3rd) 6086.8 4561.5 4942.3 4524.6 4849

The high ionization potential (IP) (2nd + 3rd) of gallium is explained by d-block contraction, and the higher IP (2nd + 3rd) of thallium relative to indium, has been explained by relativistic effects.[4]

An important consideration is that compounds in the lower oxidation state are ionic, whereas the compounds in the higher oxidation state tend to be covalent. Therefore, covalency effects must also be taken into account. In fact an alternative explanation of the inert pair effect by Drago in 1958 attributed the effect to low M-X bond enthalpies for the heavy p-block elements and the fact that it requires less energy to oxidize an element to a low oxidation state than to a higher oxidation state.[5] This energy has to be supplied by ionic or covalent bonds, so if bonding to a particular element is weak, the high oxidation state may be inaccessible. Further work involving relativistic effects confirms this.[6] In view of this it has been suggested that the term inert pair effect should be viewed as a description rather than as an explanation.[2]

Steric activity of the lone pair[edit]

The chemical inertness of the s electrons in the lower oxidation state is not always married to steric inertness (where steric inertness means that the presence of the s electron lone pair has little or no influence on the geometry of the molecule or crystal). A simple example of steric activity is that of SnCl2 which is bent in accordance with VSEPR. Some examples where the lone pair appears to be inactive are bismuth(III) iodide, BiI3, and the BiI3−
anion. In both of these the central Bi atom is octahedrally coordinated with little or no distortion, in contravention to VSEPR theory.[7] The steric activity of the lone pair has long been assumed to be due to the orbital having some p character, i.e. the orbital is not spherically symmetric.[2] More recent theoretical work shows that this is not always necessarily the case. For example, the litharge structure of PbO contrasts to the more symmetric and simpler rock salt structure of PbS and this has been explained in terms of PbII − anion interactions in PbO leading to an asymmetry in electron density. Similar interactions do not occur in PbS.[8] Another example are some thallium(I) salts where the asymmetry has been ascribed to s electrons on Tl interacting with antibonding orbitals.[9]

Failure of the theory[edit]

Part of the rationale for describing this as an effect was the fact that, at the time when it was proposed, there were no known compounds of Group 13 elements with the intermediate, +2, oxidation state. This is no longer true since the discovery of certain complexes of Ga(II) and In(II), such as halides of the form [M2X6]2−.[10][11] These complex ions are stabilized by the formation of a covalent M–M bond. It follows that the instability of simple complexes of ions such as Ga2+ is due to kinetic factors, namely that the Ga2+, having an unpaired electron, behaves as a free radical and is rapidly destroyed by reaction with another free radical. Drago's explanation[5] as described above was correct, these compounds are thermodynamically stable by virtue of the formation of a covalent bond between the gallium ions.


  1. ^ Sidgwick, Nevil Vincent (1927). The Electronic Theory of Valency. Oxford: Clarendon. pp. 178–81. 
  2. ^ a b c Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9. 
  3. ^ Electronegativity UC Davis ChemWiki by University of California, Davis
  4. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  5. ^ a b Russell S. Drago (1958). "Thermodynamic Evaluation of the Inert Pair Effect.". J Phys Chem. 62 (3): 353–357. doi:10.1021/j150561a027. 
  6. ^ Schwerdtfeger P, Heath GA, Dolg M, Bennet MA (1992). "Low valencies and periodic trends in heavy element chemistry. A theoretical study of relativistic effects and electron correlation effects in Group 13 and Period 6 hydrides and halides". Journal of the American Chemical Society. 114 (19): 7518–7527. doi:10.1021/ja00045a027. 
  7. ^ Ralph A. Wheeler and P. N. V. Pavan Kumar (1992). "Stereochemically active or inactive lone pair electrons in some six-coordinate, group 15 halides". Journal of the American Chemical Society. 114 (12): 4776–4784. doi:10.1021/ja00038a049. 
  8. ^ Walsh A, Watson GW (2005). "The origin of the stereochemically active Pb(II) lone pair: DFT calculations on PbO and PbS". Journal of Solid State Chemistry. 178 (5): 1422–1428. Bibcode:2005JSSCh.178.1422W. doi:10.1016/j.jssc.2005.01.030. 
  9. ^ Mudring AJ, Rieger F (2005). "Lone Pair Effect in Thallium(I) Macrocyclic Compounds". Inorg. Chem. 44 (18): 6240–6243. doi:10.1021/ic050547k. PMID 16124801. 
  10. ^ Freeland, B.H., Hencher, J.L.,Tuck, D.G. and Contreras, J,G. (1976). "Coordination compounds of indium. 32. Preparation and properties of hexahalogenatodiindate(II) anions". Inorganic Chemistry. 15 (9): 2144–2146. doi:10.1021/ic50163a028. 
  11. ^ Drake, J.E., Hencher, Khasrou, L.N., ,Tuck, D.G. and Victoriano, L. (1980). "Coordination compounds of indium. 34. Preparative and spectroscopic studies of InX3Y and InX2Y2 anions (X .noteq. Y = chlorine, bromine, iodine)". Inorganic Chemistry. 19 (1): 34–38. doi:10.1021/ic50203a008. 

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