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Chapter 22: Problem 69

Oxygen forms double bonds in \(\mathrm{O}_{2}\), but sulfur forms single bonds in \(\mathrm{S}_{8}\). Explain.

Short Answer

Expert verified
Oxygen in \(\mathrm{O}_{2}\) forms double bonds due to its need to share two electrons, but its inability to expand its valence shell beyond eight electrons. Sulphur in \(\mathrm{S}_{8}\) forms single bonds in a crown-shaped ring structure, being able to expand its valence shell using d-orbitals, and attains stability through bonding and non-bonding pairs of electrons.

Step by step solution

01

Explore The Electron Configuration of Oxygen

Oxygen has 6 electrons in its outermost shell, needing 2 more to achieve the stable electron configuration of the nearest noble gas, Neon. This makes oxygen a group 16 element in the periodic table, which comprises nonmetals and chalcogens. In the formation of \(\mathrm{O}_{2}\), one oxygen atom can share two of its valence electrons with another oxygen atom. This sharing of electrons leads to the formation of a double bond.
02

Explore The Electron Configuration of Sulphur

Sulfur, like oxygen, is also a group 16 element on the periodic table, but it differs from oxygen in that it has a larger atomic radius and lower electronegativity. Sulfur has 6 electrons in its outer shell and also requires 2 more to attain a stable electron configuration. But in its most stable molecular form, \(\mathrm{S}_{8}\), each sulfur atom forms a single bond with two other sulfur atoms, utilizing two of its six valence electrons. The remaining four electrons are nonbonding or lone pair electrons.
03

Summary of Orbital Overlap and Bonding

In \(\mathrm{O}_{2}\), oxygen’s inability to expand its valence shell limits it to forming double bonds, using p-orbital overlap. In \(\mathrm{S}_{8}\), sulfur atoms can utilize d-orbitals to expand its valence shell allowing for greater flexibility in bonding configurations, leading to formation of a single bond in a crown-shaped ring for increased stability.

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