Question

The following illustration shows the orbitals used to form the bonds in carbon dioxide.

Each colour represents a different orbital. Label each orbital, draw the Lewis structure for carbon dioxide, and explain how the localized electron model describes the bonding in $\mathit{C}{\mathit{O}}_{\mathbf{2}}$.

Short answer

Dark green has $sp$orbital, light green has $s{p}^2$and orange $p$orbitals.

Step-by-step solution

The given question and Lewis structure

A Lewis Structure is a general representation of a molecule's valence shell electrons. It's used to explain how electrons in a molecule are distributed around specific atoms. Electrons are depicted as "dots" or as a line between two atoms when they are bonded.

Draw the lewis structure give the explanation of the description of the localized electron model about the bonding in ${\mathrm{CO}}_2$.

The description of the localized electron model

The light green orbitals surrounding oxygen are $s{p}^2$hybrid orbitals, while the dark green orbitals around carbon are $sp$hybrid orbitals. Unhybridized $p$atomic orbitals are represented by the remaining orange orbitals.

The pi bond is created by the side-to-side overlap of unhybridized $p$atomic orbitals from carbon with an unhybridized$p$atomic orbital from each oxygen, as shown in the diagram below.

In addition, the overlap of hybrid orbitals from carbon with $s{p}^2$hybrid orbitals from each oxygen results in the formation of two carbon-oxygen sigma bonds.

Two pi bonds perpendicular to each other and two sigma bonds make up a molecule ${\mathrm{CO}}_2$.

So, dark green has $sp$orbital, light green has $s{p}^2$and orange$p$orbitals