- What is the octet rule?
- Examples of the octet rule
- Exceptions to the octet rule
- Octet rule and Lewis structure
We explain what the octet rule is in chemistry, who was its creator, examples and exceptions. Also, the Lewis structure.
What is the octet rule?
In chemistry, is known as the octet rule or octet theory to the explanation of the way in which the atoms of the chemical elements it combines.
This theory was enunciated in 1917 by the American chemical physicist Gilbert N. Lewis (1875-1946) and explains that the atoms of the different elements usually always maintain a stable electronic configuration by locating eight electrons in your last energy levels.
The octet rule states that the ions of the various chemical elements found in the Periodic Table usually complete their last energy levels with 8 electrons. Because of this, molecules can acquire a stability similar to that of Noble gases (located at the extreme right of the periodic table), whose electronic structure (with its last full energy level) makes them very stable, that is, not very reactive.
Thus, elements with high electronegativity (such as halogens and amphigens, that is, elements of group 16 of the Table) tend to “gain” electrons up to the octet, while those with low electronegativity (such as alkaline or alkaline earth) tend to "lose" electrons to reach the octet.
This rule explains one of the ways in which atoms form their bonds, and the behavior and chemical properties of the resulting molecules will depend on their nature. Thus, the octet rule is a practical principle that serves to predict the behavior of many substances, although it also presents different exceptions.
Examples of the octet rule
Consider a CO2 molecule whose atoms have valences of 4 (carbon) and 2 (oxygen), joined by chemical links double. (It is important to clarify that valence are the electrons that a chemical element must give up or accept to achieve its last energy level to be complete. Chemical valence should not be confused with valence electrons, since the latter are the electrons that are are located in the last energy level).
This molecule is stable if each atom has 8 electrons in total at its last energy level, reaching the stable octet, which is fulfilled with the 2-electron compartment between carbon and oxygen atoms:
- Carbon shares two electrons with each oxygen, increasing the electrons at the last energy level of each oxygen from 6 to 8.
- At the same time, each oxygen shares two electrons with carbon, increasing the electrons from 4 to 8 in the last energy level of carbon.
Another way of looking at it would be that the total of the electrons transferred and taken must always be eight.
That is the case for other stable molecules, such as sodium chloride (NaCl).Sodium contributes its single electron (valence 1) to chlorine (valence 7) to complete the octet. Thus, we would have Na1 + Cl1- (that is, sodium gave up an electron, and gained a positive charge, and chlorine accepted an electron and with it a negative charge).
Exceptions to the octet rule
The octet rule has several exceptions, that is, compounds that achieve stability without being governed by the electron octet. Atoms such as phosphorus (P), sulfur (S), selenium (Se), silicon (Si) or helium (He) can accommodate more electrons than suggested by Lewis (hypervalence).
In contrast, hydrogen (H), which has a single electron in a single atomic orbital (the region of space where an electron is most likely to be found around the atomic nucleus), can accept up to two electrons in a chemical bond. Other exceptions are beryllium (Be), which acquires stability with just four electrons, or boron (B), which does so with six.
Octet rule and Lewis structure
Another of Lewis's great contributions to chemistry was his famous way of representing atomic bonds, today known as the “Lewis structure” or “Lewis formula”.
It consists of placing dots or dashes to represent the shared electrons in a molecule and the electrons that are free on each atom.
This type of two-dimensional graphical representation allows to know the valence of an atom that interacts with others in a compound and whether it forms single, double, or triple bonds, all of which will affect molecular geometry.
To represent a molecule in this way we need to choose a central atom, which will be surrounded by the others (called terminals) establishing bonds until reaching the valences of all those involved. The former are usually the least electronegative and the latter the most electronegative.
For example, the representation of Water (H2O) shows the free electrons that the oxygen atom has, in addition you can visualize the simple bonds between the oxygen atom and the hydrogen atoms (the electrons that belong to the oxygen atom are represented in red and those of the atoms hydrogen in black). The acetylene molecule (C2H2) is also represented, where you can visualize the triple bond between the two carbon atoms and the single bonds between each carbon atom and a hydrogen atom (the electrons that belong to the carbon atoms are represented in red and those of hydrogen atoms in black).