Valence Electrons and the Octet Rule

The electrons that occupy the outermost shell of an atom are called valence electrons, and it is these electrons — not the inner, tightly-held core electrons — that participate in chemical bonding. When atoms combine, they do so in a way that leaves each of them with a more stable, lower-energy electron arrangement, most often the eight-electron valence shell of the nearest noble gas.

This observation is generalised as the octet rule: atoms tend to gain, lose, or share electrons so as to acquire eight electrons in their valence shell, matching the stable $ns^2np^6$ configuration of a noble gas. Sodium ($1s^2 2s^2 2p^6 3s^1$) achieves this by losing its single $3s$ electron to become $\text{Na}^+$ (matching neon's configuration); chlorine ($1s^2 2s^2 2p^6 3s^2 3p^5$) achieves it by gaining one electron to become $\text{Cl}^-$ (matching argon's configuration). Hydrogen and helium are exceptions from the start — they are satisfied with just two electrons (a duplet), matching helium's $1s^2$ shell, since their valence shell is the first shell and can hold no more than two electrons.

The octet rule is an extremely useful first approximation, but Class 11 chemistry meets several genuine exceptions to it later in this very chapter:

• Electron-deficient molecules, where the central atom has fewer than eight valence electrons around it even after bonding. Beryllium chloride, $\text{BeCl}_2$, has only four electrons around beryllium, and boron trifluoride, $\text{BF}_3$, has only six around boron.
• Expanded-octet molecules, where the central atom, by virtue of having accessible $d$ orbitals, accommodates more than eight electrons. Phosphorus pentachloride, $\text{PCl}_5$ (ten electrons around P), and sulphur hexafluoride, $\text{SF}_6$ (twelve electrons around S), are the standard examples.
• Odd-electron species, such as nitric oxide, $\text{NO}$, which has an odd total number of valence electrons and so can never be arranged into complete pairs at all.

None of these exceptions overturns the octet rule as a guiding principle — they simply show that the rule describes a strong tendency, governed by the underlying goal of achieving the lowest-energy, most stable electron arrangement, rather than an absolute law of nature. Both the electron-deficient and expanded-octet cases reappear when this chapter discusses VSEPR theory and hybridisation, where the true shapes of $\text{BeCl}_2$, $\text{BF}_3$, $\text{PCl}_5$ and $\text{SF}_6$ are worked out in detail.

Understanding valence electrons and the octet rule is the essential first step before the chapter's two main bonding models — the electron-transfer picture of the ionic bond and the electron-sharing picture of the covalent bond — can be developed.