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Example 7.11: Predicting Electron-pair Geometry and Molecular Structure: \(\ce{CO2}\) and \(\ce{BCl3}\)

Predict the electron-pair geometry and molecular structure for each of the following:
(a) carbon dioxide, \(\ce{CO2}\), a molecule produced by the combustion of fossil fuels
(b) boron trichloride, \(\ce{BCl3}\), an important industrial chemical

Solution

(a) We write the Lewis structure of \(\ce{CO2}\) as:

This shows us two regions of high electron density around the carbon atom—each double bond counts as one region, and there are no lone pairs on the carbon atom. Using \(\ce{VSEPR}\) theory, we predict that the two regions of electron density arrange themselves on opposite sides of the central atom with a bond angle of 180°. The electron-pair geometry and molecular structure are identical, and \(\ce{CO2}\) molecules are linear.


(b) We write the Lewis structure of \(\ce{BCl3}\) as:

Thus we see that \(\ce{BCl3}\) contains three bonds, and there are no lone pairs of electrons on boron. The arrangement of three regions of high electron density gives a trigonal planar electron-pair geometry. The B–Cl bonds lie in a plane with 120° angles between them. \(\ce{BCl3}\) also has a trigonal planar molecular structure ( Figure 21).

The electron-pair geometry and molecular structure of \(\ce{BCl3}\) are both trigonal planar. Note that the \(\ce{VSEPR}\) geometry indicates the correct bond angles (120°), unlike the Lewis structure shown above.