18-electron rule

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The 18-electron rule (18-valence electron rule) in chemistry states that complexes with 18 total valence electrons are particularly stable. The rule can be applied to many complexes of transition metals . It results from the molecular orbital theory and applies to elements from the fourth period of the periodic table .
The 18-electron rule is thus the counterpart to the well-known " octet rule " of the main group elements.

Examples

The transition metal chromium forms the compound chromium hexacarbonyl [Cr (CO) 6 ], iron forms the stable iron pentacarbonyl [Fe (CO) 5 ], while nickel forms the particularly stable nickel tetracarbonyl [Ni (CO) 4 ]. In all three cases the corresponding metal center ( oxidation state = 0) has a total valence electron number of 18 and thus the noble gas configuration of krypton : chromium itself has six valence electrons, iron eight and nickel ten. Since each CO molecule contributes two electrons to the coordinative bond, twelve valence electrons must be added for [Cr (CO) 6 ], ten for [Fe (CO) 5 ] and eight for [Ni (CO) 4 ].

Limits of the model

With the 18-electron rule, e.g. B. the stability of ferrocene (18 electrons) and the reducing character of the metallocene compounds cobaltocene (19 electrons) and nickelocene (20 electrons) can be explained. Nickelocene is less reactive because there are 2 electrons in a weakly antibonding orbital .

Steric reasons

With the early transition metals, the 18-electron rule is often not fulfilled for steric reasons. This means that there is not enough space around the central particle to store enough ligands - and thus missing electrons. For example, vanadium hexacarbonyl [V (CO) 6 ] has only 17 electrons on the vanadium atom . A conceivable way out of this electron deficiency would be dimerization with the formation of a covalent V – V bond (gain of a common electron). However, this reaction is no longer possible for reasons of space. However, [V (CO) 6 ] acts as a moderately strong oxidizing agent, since it is converted into the anion [V (CO) 6 ] - with 18 total valence electrons by taking up an electron .

Case of predominantly electrostatic interactions

Another reason for not complying with the rule is the presence of more electrostatic (electrovalent, ionic) bond relationships. Orbital overlaps are not critical here. So there is no need to follow any rules that come from orbital theory. The bond is mainly based on classic electrostatics . A typical example is the well-known stable copper tetrammine complex [Cu (NH 3 ) 4 ] 2+ , which has 17 total valence electrons (Cu 2+ : 9 e - + 4 × NH 3 : 4 × 2 e - ) according to the usual counting method should.

The complexes of early transition metals mentioned should also be mentioned again from this point of view. Among the transition metals, these metals have the smallest electronegativities and therefore, with many complex ligands, tend to form electrostatically bound complexes that do not (need) meet the 18-electron rule. The same applies to complexes of alkali and alkaline earth metals - they too do not have to meet either the 8-electron or the 18-electron rule, since covalent bond components are negligibly small here too.

Molecular orbital theoretical consideration

With the help of molecular orbital theory , the energetic position of the orbitals in the complex can be determined, whereby a distinction is made between binding, non-binding and anti-binding orbitals. The occupation of binding orbitals contributes to the stability of the complex, the occupation of anti-binding orbitals is associated with reduced stability. All octahedral complexes have in common that the a 1g , the two e g and the three t 1u orbitals are strongly binding. The stability of the complex resulting from the occupation of these orbitals means that practically all octahedral complexes have a minimum number of 12 valence electrons. Further frontier orbitals are the three non-bonding t 2g and the two antibonding e g * orbitals, the character of which depends on the ligands:

  • Weak ligands (pure σ donors and σ / π donors) cause only a weak antibonding character of the e g * orbitals in elements of the 3d block and, in the case of π donors, raise the t 2g orbitals to a slightly antibonding character on. However, the antibonding character is so small that it hardly makes any difference how the orbitals are occupied. Here 12 (minimum number from above) to 22 (full occupancy of the t 2g and e g * orbitals) valence electrons are observed.
  • In the case of elements of the 4d / 5d block, even weak ligands have a strong antibonding character in the e g * orbitals, which makes their occupation unfavorable. The degree of occupation of the t 2g orbitals, however, still does not play an essential role, so that 12 to 18 valence electrons are observed.
  • Strong ligands (σ donors and π acceptors) generally cause a strong increase in the energy of the e g * orbitals and, via the π backbonding, lower the t 2g orbitals into a binding character. The occupation of the t 2g orbitals is thus very favorable, that of the e g * orbitals is unfavorable. This leads to a particularly high stability of complexes with 18 valence electrons.

The 18-electron rule therefore applies in particular to complexes with strong ligands such as CO and CN - .

The above considerations apply in this form only to octahedral complexes. However, tetrahedral complexes can be investigated in a similar way, again with the result that the 18-electron rule only applies to complexes with strong ligands. In contrast, in square-planar complexes there is a 16-electron rule for strong ligands.

See also

literature

Individual evidence

  1. ^ Richard Göttlich, Siegfried Schindler, Parham Rooshenas: Basic chemical internship in a minor subject , Pearson Verlag, 2011, ISBN 978-3-86894-030-5 , p. 95.