Abegg’s rule states that the difference between the maximum positive and negative valence of an element usually is eight. Another name for the rule is “Abegg’s law of valence and countervalence.” German chemist Richard Abegg proposed the rule in 1904.
Example of Abegg’s Rule
For example, the negative valence of the element sulfur (S) is -2 in the compound H2S and its positive valence (counter valence) is +6 in H2SO4. The difference between -2 and +6 is 8.
How Abegg’s Rule Works
In the context of Abegg’s rule, valence describes whether an atom acts as an electron donor or receiver. This aligns with the modern concept of the oxidation state. For example, the group 5 elements are pentavalent (have 5 valence electrons). An atom from group 5 (e.g., vanadium, niobium, tantalum) acts as an electron donor (-3) or can also act as an electron acceptor (+5). In either situation, the atom achieves a stable octet when it forms chemical bonds. The difference between the normal valence (-3) and contra valence (+5) is 8.
Exceptions to Abegg’s Rule
Abegg’s “rule” is more of a guideline. It does not work for all elements. The obvious exception is hydrogen, which ranges in valence from +1 to -1. In other words, the hydrogen atom gains or loses a single electron. With a single proton, hydrogen does not have a nucleus that can accommodate enough electrons for an octet.
Other elements that violate the octet rule sometimes violate Abegg’s rule. For example, the elements silicon, phosphorus, sulfur, and chlorine all sometimes bond to more than four atoms. They go beyond satisfying the s2p6 octet. Atoms from these elements have five 5 orbitals that can participate in bonding. Applying an “Even-Odd” rule to Abegg’s rule helps with the expanded octet exceptions.
An atom can violate the octet rule (have an expanded octet) and still satisfy Abegg’s rule. In the case of sulfur hexafluoride (SF6), sulfur has 12 bonding electrons (+6) and bonds to fix fluorine atoms. The normal valence of sulfur is -2, while the contra valence is +6, with a difference of 8.
Some atoms may have oxidation states greater than +8. For example the oxidation state of iridium ranges from -3 to +10 in [PtO4]2+. These atoms are exceptions to Abegg’s rule.
Importance of Abegg’s Rule
Abegg’s rule is important because of its influence on other scientists. Gilbert N. Lewis applied Abegg’s rule in his cubical atom theory (1916), which ultimately led to the development of the octet rule. Linus Pauling’s influential 1938 text, The Nature of the Chemical Bond, drew upon the work of Abegg and Lewis.
- Abegg, R. (1904). “Die Valenz und das periodische System. Versuch einer Theorie der Molekularverbindungen” [Valency and the periodic table. Attempt at a theory of molecular compounds]. Zeitschrift für anorganische Chemie (in German). 39 (1): 330–380. doi:10.1002/zaac.19040390125
- Auvert, Geoffroy (2104). “Improvement of the Lewis-Abegg-Octet Rule Using an “Even-Odd” Rule in Chemical Structural Formulas: Application to Hypo and Hyper-Valences of Stable Uncharged Gaseous Single-Bonded Molecules with Main Group Elements”. Open Journal of Physical Chemistry. 4(2): 60-66. doi:10.4236/ojpc.2014.42009
- Housecroft, Catherine E.; Sharpe, Alan G. (2005). Inorganic Chemistry (2nd ed.). Pearson Education Limited. ISBN 0130-39913-2.
- Lewis, Gilbert N. (1916-04-01). “The atom and the molecule”. Journal of the American Chemical Society. 38 (4): 762–785. doi:10.1021/ja02261a002
- Pauling, Linus (1960). The Nature of the Chemical Bond and the Structure of Molecules and Crystals; An Introduction to Modern Structural Chemistry (3rd ed.). Cornell University Press. ISBN 0-8014-0333-2.
- Ritter, Stephen K. (2016). “Oxidation state +10 a possibility“. C&EN. 94(25).