Inert pair effect

Last updated

The inert pair effect is the tendency of the two 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 outermost s electrons are more tightly bound to the nucleus in these atoms, and therefore more difficult to ionize or share.


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 bond formation .


Consider as an example thallium (Tl) in group 13. The +1 oxidation state of Tl is the most stable, while 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. The heaviest members of each group, i.e. 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 two valence electrons in s orbitals. A partial explanation is 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] If the total ionization potentials (IP) (see below) of the two electrons in s orbitals (the 2nd + 3rd ionization potentials), are examined it can be seen that there is an expected decrease from B to Al associated with increased atomic size but the values for Ga, In and Tl are higher than expected.

Ionization potentials for group 13 elements
IP Boron Aluminium Gallium Indium Thallium
(2nd + 3rd)6,0864,5604,9424,5244,849

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] The higher value for thallium compared to indium is partly attributable to the influence of the lanthanide contraction and the ensuing poor shielding from the nuclear charge by the intervening filled 4d and 5f subshells. [5]

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 be taken into account. 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. [6] 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. [7]

In the case of Groups 13 to 15 the inert pair effect has been further attributed to, "the decrease in bond energy with the increase in size from Al to Tl so that the energy required to involve the s electron in bonding is not compensated by the energy released in forming the two additional bonds." [2] That said, the authors note several factors are at play, including relativistic effects in the case of gold, and that, "a quantitative rationalisation of all the data has not been achieved." [2]

Steric activity of the lone pair

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. [8] 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. [9] Another example are some thallium(I) salts where the asymmetry has been ascribed to s electrons on Tl interacting with antibonding orbitals. [10]

Related Research Articles

Electronegativity, symbol χ, is a chemical property that describes the tendency of an atom to attract a shared pair of electrons towards itself. An atom's electronegativity is affected by both its atomic number and the distance at which its valence electrons reside from the charged nucleus. The higher the associated electronegativity number, the more an atom or a substituent group attracts electrons towards itself.

Ligand molecule or functional group that binds or can bind to the central atom in a coordination complex

In coordination chemistry, a ligand is an ion or molecule that binds to a central metal atom to form a coordination complex. The bonding with the metal generally involves formal donation of one or more of the ligand's electron pairs. The nature of metal–ligand bonding can range from covalent to ionic. Furthermore, the metal–ligand bond order can range from one to three. Ligands are viewed as Lewis bases, although rare cases are known to involve Lewis acidic "ligands".

Thallium Chemical element with atomic number 81

Thallium is a chemical element with the symbol Tl and atomic number 81. It is a gray post-transition metal that is not found free in nature. When isolated, thallium resembles tin, but discolors when exposed to air. Chemists William Crookes and Claude-Auguste Lamy discovered thallium independently in 1861, in residues of sulfuric acid production. Both used the newly developed method of flame spectroscopy, in which thallium produces a notable green spectral line. Thallium, from Greek θαλλός, thallós, meaning "a green shoot or twig", was named by Crookes. It was isolated by both Lamy and Crookes in 1862; Lamy by electrolysis, and Crookes by precipitation and melting of the resultant powder. Crookes exhibited it as a powder precipitated by zinc at the International exhibition, which opened on 1 May that year.

In chemistry, the term transition metal has three possible meanings:

The oxidation state, sometimes referred to as oxidation number, describes the degree of oxidation of an atom in a chemical compound. Conceptually, the oxidation state, which may be positive, negative or zero, is the hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic, with no covalent component. This is never exactly true for real bonds.

Moscovium is a synthetic chemical element with the symbol Mc and atomic number 115. It was first synthesized in 2003 by a joint team of Russian and American scientists at the Joint Institute for Nuclear Research (JINR) in Dubna, Russia. In December 2015, it was recognized as one of four new elements by the Joint Working Party of international scientific bodies IUPAC and IUPAP. On 28 November 2016, it was officially named after the Moscow Oblast, in which the JINR is situated.

Double bond Chemical bond involving four bonding electrons; has one sigma plus one pi bond

A double bond in chemistry is a covalent bond between two atoms involving four bonding electrons instead of the usual two. Double bonds occur most commonly between two carbon atoms, for example in alkenes. Many double bonds exist between two different elements: for example, in a carbonyl group between a carbon atom and an oxygen atom. Other common double bonds are found in azo compounds (N=N), imines (C=N) and sulfoxides (S=O). In a skeletal formula, a double bond is drawn as two parallel lines (=) between the two connected atoms; typographically, the equals sign is used for this. Double bonds were first introduced in chemical notation by Russian chemist Alexander Butlerov.

Octet rule Chemical rule of thumblets

The octet rule is a chemical rule of thumb that reflects the observation that elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. The rule is especially applicable to carbon, nitrogen, oxygen, and the halogens, but also to metals such as sodium or magnesium.

Lone pair pair of valence electrons that are not shared with another atom; concept used in valence shell electron pair repulsion theory (VSEPR theory) which explains the shapes of molecules

In chemistry, a lone pair refers to a pair of valence electrons that are not shared with another atom in a covalent bond and is sometimes called an unshared pair or non-bonding pair. Lone pairs are found in the outermost electron shell of atoms. They can be identified by using a Lewis structure. Electron pairs are therefore considered lone pairs if two electrons are paired but are not used in chemical bonding. Thus, the number of lone pair electrons plus the number of bonding electrons equals the total number of valence electrons around an atom.

VSEPR theory theoretical model used in chemistry

Valence shell electron pair repulsion theory, or VSEPR theory, is a model used in chemistry to predict the geometry of individual molecules from the number of electron pairs surrounding their central atoms. It is also named the Gillespie-Nyholm theory after its two main developers, Ronald Gillespie and Ronald Nyholm. The premise of VSEPR is that the valence electron pairs surrounding an atom tend to repel each other and will, therefore, adopt an arrangement that minimizes this repulsion, thus determining the molecule's geometry. Gillespie has emphasized that the electron-electron repulsion due to the Pauli exclusion principle is more important in determining molecular geometry than the electrostatic repulsion.

In chemistry, orbital hybridisation is the concept of mixing atomic orbitals into new hybrid orbitals suitable for the pairing of electrons to form chemical bonds in valence bond theory. Hybrid orbitals are very useful in the explanation of molecular geometry and atomic bonding properties and are symmetrically disposed in space. Although sometimes taught together with the valence shell electron-pair repulsion (VSEPR) theory, valence bond and hybridisation are in fact not related to the VSEPR model.

Relativistic quantum chemistry combines relativistic mechanics with quantum chemistry to explain elemental properties and structure, especially for the heavier elements of the periodic table. A prominent example of such an explanation is the color of gold: due to relativistic effects, it is not silvery like most other metals.

In chemistry, the valence or valency of an element is a measure of its combining power with other atoms when it forms chemical compounds or molecules. The concept of valence was developed in the second half of the 19th century and helped successfully explain the molecular structure of inorganic and organic compounds. The quest for the underlying causes of valence led to the modern theories of chemical bonding, including the cubical atom (1902), Lewis structures (1916), valence bond theory (1927), molecular orbitals (1928), valence shell electron pair repulsion theory (1958), and all of the advanced methods of quantum chemistry.

The 3-center 4-electron (3c–4e) bond is a model used to explain bonding in certain hypervalent molecules such as tetratomic and hexatomic interhalogen compounds, sulfur tetrafluoride, the xenon fluorides, and the bifluoride ion. It is also known as the Pimentel–Rundle three-center model after the work published by George C. Pimentel in 1951, which built on concepts developed earlier by Robert E. Rundle for electron-deficient bonding. An extended version of this model is used to describe the whole class of hypervalent molecules such as phosphorus pentafluoride and sulfur hexafluoride as well as multi-center π-bonding such as ozone and sulfur trioxide.

d-block contraction term used in chemistry to describe the effect of having full d orbitals on the period 4 elements

The d-block contraction is a term used in chemistry to describe the effect of having full d orbitals on the period 4 elements. The elements in question are Ga, Ge, As, Se, Br, and Kr. Their electronic configurations include completely filled d orbitals (d10). d-block contraction is best illustrated by comparing some properties of the group 13 elements to highlight the effect on gallium.

An electronic effect influences the structure, reactivity, or properties of molecule but is neither a traditional bond nor a steric effect. In organic chemistry, the term stereoelectronic effect is also used to emphasize the relation between the electronic structure and the geometry (stereochemistry) of a molecule.

In inorganic chemistry, the cis effect is defined as the labilization (making unstable) of CO ligands that are cis to other ligands. CO is a well-known strong pi-accepting ligand in organometallic chemistry that will labilize in the cis position when adjacent to ligands due to steric and electronic effects. The system most often studied for the cis effect is an octahedral complex M(CO)5X where X is the ligand that will labilize a CO ligand cis to it. Unlike trans effect, where this property is most often observed in 4-coordinate square planar complexes, the cis effect is observed in 6-coordinate octahedral transition metal complexes. It has been determined that ligands that are weak sigma donors and non-pi acceptors seem to have the strongest cis-labilizing effects. Therefore, the cis effect has the opposite trend of the trans-effect, which effectively labilizes ligands that are trans to strong pi accepting and sigma donating ligands.

Stereoelectronic effect effect on molecular structures, physical properties and reactivities due to the molecules electronic structures, in particular the interaction between atomic and/or molecular orbitals

In chemistry, primarily organic and computational chemistry, a stereoelectronic effect is an effect on molecular geometry, reactivity, or physical properties due to spatial relationships in the molecules' electronic structure, in particular the interaction between atomic and/or molecular orbitals. Phrased differently, stereoelectronic effects can also be defined as the geometric constraints placed on the ground and/or transition states of molecules that arise from considerations of orbital overlap. Thus, a stereoelectronic effect explains a particular molecular property or reactivity by invoking stabilizing or destabilizing interactions that depend on the relative orientations of electrons in space.

Plumbylene class of chemical compounds

Plumbylenes (or plumbylidenes) are divalent organolead(II) analogues of carbenes, with the general chemical formula, R2Pb, where R denotes a substituent. Plumbylenes possess 6 electrons in their valence shell, and are considered open shell species.

Nontrigonal pnictogen compounds

Nontrigonal pnictogen compounds refer to tricoordinate trivalent pnictogen compounds that are not of typical trigonal pyramidal molecular geometry. By virtue of their geometric constraint, these compounds exhibit distinct electronic structures and reactivities, which bestow on them potential to provide unique nonmetal platforms for bond cleavage reactions.


  1. Sidgwick, Nevil Vincent (1927). The Electronic Theory of Valency. Oxford: Clarendon. pp. 178–81.
  2. 1 2 3 4 Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN   978-0-08-037941-8.
  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. Rodgers, G.; E. (2014). "A visually attractive "Interconnected network of ideas" for organizing the teaching and learning of descriptive inorganic chemistry". Journal of Chemical Education. 91 (2): 216−224 (219). doi:10.1021/ed3003258.
  6. Russell S. Drago (1958). "Thermodynamic Evaluation of the Inert Pair Effect". J Phys Chem. 62 (3): 353–357. doi:10.1021/j150561a027.
  7. 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.
  8. 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.
  9. 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.
  10. 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.