Octet rule

The octet theory, enunciated in 1916 by the physical chemist Gilbert Newton Lewis, says that the ions of the elements of the periodic system have the tendency to complete their last energy levels with a amount of 8 electrons, in such a way that they acquire a very stable configuration. This configuration is similar to that of a noble gas, the elements located to the far right of the periodic table. The noble gases are electrochemically stable elements, since they comply with the Lewis structure, they are inert, that is, it is very difficult for them to react with any other element. This conclusion is known as the octet rule.

This rule is applicable to the creation of bonds between atoms, the nature of these bonds will determine the behavior and properties of the molecules. These properties will therefore depend on the type of bond, the number of bonds per atom, and the intermolecular forces.

There are different types of chemical bonds, all of them based, as explained before, on the special stability of the electronic configuration of noble gases, tending to be surrounded by eight electrons at its outermost level. This electronic octet can be acquired by an atom in different ways, depending on the electronegativity:

• Ionic link
• covalent link
• metallic link
• intermolecular links
• Coordinated link

It's important to know that the octet rule is a rough rule of thumb with many exceptions, but it can predict the behavior of many substances.

CO2, with two double links.

The figure shows the four valence electrons of carbon, creating two covalent bonds, with the six electrons in the last energy level of each of the oxygens, whose valence is 2. The sum of the electrons of each one of the atoms is 8, which leads to the octet. Note that there are cases of molecules with atoms that do not meet the octet and are also stable.

Exceptions

There are exceptions to this rule. The atoms that do not comply with the octet rule in some compounds are: Phosphorus, Sulfur, Selenium, Silicon and Helium. It generally occurs in elements of the main group from the third period(nP, n≥3). These elements have the availability to accommodate a greater number of electrons in the (n+1)P orbitals, this behavior is called hypervalence. First introduced in 1969 by Jeremy Musher.

Hydrogen has a single orbital in its valence shell which can accept a maximum of two electrons, along with beryllium which is completed with a quantity of four electrons and boron which requires 6 electrons to carry out this function, so that the regulation that specifies that every element is completed with 8 electrons at its disposal is evaded. On the other hand, non-metallic atoms from the third period can form "expanded octets" that is, they can hold more than eight electrons in their valence shell, usually placing the extra electrons in subshells.

Some highly reactive molecules or ions have atoms with less than eight electrons in their outer layer. An example is the trifluoride of boron (BF3{displaystyle {ce {BF3}}}). In the molecule BF3{displaystyle {ce {BF3}}} the central boron atom only has six electrons around it.

The clearest way to see graphically the functioning of the "arrangement" is Lewis' representation of molecules. Before you can write some structures of Lewis, you should know how atoms are united with each other. Consider nitric acid for example. Although the nitric acid formula is often represented as HNO3{displaystyle {ce {HNO3}}, actually hydrogen is attached to oxygen, not nitrogen. The structure is MAN2{displaystyle {ce {HONO2}}} and no HNO3{displaystyle {ce {HNO3}}.

It can also occur when there are odd molecules, hypovalent molecules and hypervalent molecules. It is when atoms form compounds by losing, gaining or sharing electrons to acquire 8 valence electrons. Hydrogen achieves the stability of helium, with 2 valence electrons.

The noble gas atoms are characterized by having all their energy levels and sublevels completely filled. The stability of the noble gases is associated with the electronic structure of its last shell that remains with the octet of electrons, this is tried to be explained by the noble gas hypothesis.

History

At the end of the 19th century, it was known that coordination compounds (formerly called "molecular compounds") were formed by combining atoms or molecules in such a way that the valences of the atoms involved were apparently tight. In 1893, Alfred Werner showed that the number of atoms or groups associated with a central atom (the so-called "coordination number") is usually 4 or 6; up to a maximum of 8 coordination numbers were known, but they were less frequent. In 1904, Richard Abegg was one of the first to expand the concept of coordination number and begin to recognize the importance of valence in distinguishing atoms as electron donors or acceptors, which would lead to the consideration of positive valence states. and negative valence, which closely resembles the modern concept of oxidation state. Abegg observed that the difference between the maximum positive and negative valences of a chemical element is, according to his model, often eight. Gilbert N. Lewis called this concept "Abegg's rule", and the used to formulate his own model of the cubic atom and the 'rule of eight', to make a distinction between valence and valence electrons. In 1919, Irving Langmuir refined these concepts further, calling them "cubic octet atom and "octet theory, which evolved to become the "octet rule".