The octet rule is a chemistry rule of thumb that says that atoms combine in a way that gives them eight electrons in their valence shells. This achieves a stable electron configuration similar to that of noble gases. The octet rule is not universal and has many exceptions, but it helps with predicting and understanding the bonding behavior of many elements.
History
American chemist Gilbert N. Lewis proposed the octet rule in 1916. Lewis observed that noble gases, with their full valence shells of eight electrons, were especially stable and unreactive. He hypothesized that other elements achieve similar stability by sharing, gaining, or losing electrons to reach a filled shell. This led to his formulation of the octet rule, which was later expanded into Lewis structures and valence bond theory.
Octet Rule Examples
Atoms follow the octet rule by either donating/accepting electrons or by sharing electrons.
- Donating/Accepting Electrons: Sodium, a member of the alkali metals, has one electron in its outermost shell and eight electrons in the next shell. To achieve a noble gas configuration, it donates the one electron, resulting in a positive sodium ion (Na+) and an octet valence electron shell.
- Accepting Electrons: Chlorine has seven electrons in its valence shell. It needs one more for a stable noble gas configuration, which it gets by accepting an electron from another atom, thus forming a negative chloride ion (Cl–).
- Sharing Electrons: Oxygen has six electrons in its valence shell and needs two more to satisfy the octet rule. In the formation of water (H2O), each hydrogen atom shares its single electron with oxygen, which in turn shares one electron with each hydrogen atom. This forms two covalent bonds and fills the oxygen’s valence shell with eight electrons, while each hydrogen atom attains the noble gas configuration of helium.
Noble gases are relatively inert because they already have an octet electron configuration. So, examples of the octet rule involve other atoms that do not have a noble gas configuration. Note that the octet rule really only applies to s and p electrons, so it works for main group elements.
Why the Octet Rule Works
The octet rule works because of the nature of electron configuration in atoms, specifically in relation to the stability provided by a full valence shell.
Electrons in atoms are organized into energy levels, or shells, and each shell has a maximum capacity of electrons it holds. The first energy level holds up to 2 electrons, the second holds up to 8, and so on. These energy levels correspond to the periods (rows) on the periodic table.
The most stable, lowest-energy electron configuration for an atom is one where its outermost shell (the valence shell) is full. This occurs naturally in the noble gases, which reside at the far right of the periodic table and are known for their stability and low reactivity. Their stability comes from their full valence shells: helium has a full first shell with 2 electrons, while the rest (neon, argon, krypton, xenon, radon) have full shells with 8 electrons. Atoms of other elements try to achieve this stable configuration by gaining, losing, or sharing electrons to fill their valence shell.
Exceptions to the Octet Rule
There are exceptions to the octet rule, particularly for elements in the third period and beyond on the periodic table. These elements accommodate more than eight electrons because they have d and f orbitals in their valence shells.
Here are a few examples of elements that do not strictly follow the octet rule:
- Hydrogen: It only accommodates 2 electrons in its valence shell (to achieve the configuration of helium), so it does not follow the octet rule.
- Helium: Similarly, helium’s valence shell is complete with just two electrons.
- Lithium and Beryllium: In the second period of the periodic table, lithium and beryllium often have less than eight electrons in their compounds.
- Boron: Boron often forms compounds in which it has only six electrons around it.
- Elements in and beyond the third period: These elements often have more than eight electrons in their valence shells in compounds. Examples include phosphorus in PCl5 (phosphorus pentachloride) or sulfur in SF6 (sulfur hexafluoride), both of which exceed the octet.
- Transition metals: Many transition metals do not follow the octet rule. For instance, iron (Fe) in FeCl2 has more than eight electrons in its valence shell.
It’s important to note that these “violations” of the octet rule don’t invalidate the rule. Instead, they highlight its limitations and point towards the more complex and nuanced reality of atomic structure and bonding.
Uses of the Octet Rule
The primary benefit of the octet rule is its simplicity and broad applicability. It allows for a straightforward understanding of molecular structures and chemical reactions, making it a powerful tool in the early stages of chemical education.
Alternatives to the Octet Rule
However, the rule is not all-encompassing. The octet rule doesn’t apply well to many molecules, including those with an odd number of electrons like nitric oxide (NO), and compounds of transition metals. Furthermore, it does not account for the relative strengths of covalent bonds and the variation in bond lengths. So, there are alternatives to the rule that cover more situations.
One significant alternative is the molecular orbital (MO) theory, which provides a more complete and detailed description of the behavior of electrons in molecules. MO theory considers the entire molecule as a whole rather than focusing on individual atoms and their electrons. It explains phenomena that the octet rule cannot, such as the color of compounds, the magnetism of molecules, and why some substances are electrical conductors while others are not.
Another alternative is the valence bond (VB) theory, which is a more complex extension of the octet rule. The VB theory involves hybridization of atomic orbitals to explain the shapes of molecules.
References
- Abegg, R. (1904). “Die Valenz und das periodische System. Versuch einer Theorie der Molekularverbindungen (Valency and the periodic system – Attempt at a theory of molecular compounds)”. Zeitschrift für anorganische Chemie. 39 (1): 330–380. doi:10.1002/zaac.19040390125
- Frenking, Gernot; Fröhlich, Nikolaus (2000). “The Nature of the Bonding in Transition-Metal Compounds”. Chem. Rev. 100 (2): 717–774. doi:10.1021/cr980401l
- Housecroft, Catherine E.; Sharpe, Alan G. (2005). Inorganic Chemistry (2nd ed.). Pearson Education Limited. ISBN 0130-39913-2.
- Langmuir, Irving (1919). “The Arrangement of Electrons in Atoms and Molecules”. Journal of the American Chemical Society. 41 (6): 868–934. doi:10.1021/ja02227a002
- Lewis, Gilbert N. (1916). “The Atom and the Molecule”. Journal of the American Chemical Society. 38 (4): 762–785. doi:10.1021/ja02261a002