In ammonia, a central nitrogen atom is bonded to three hydrogen atoms. Similarly, boron trifluoride has a central boron atom bonded to three fluorine atoms. However, ammonia is pyramidal and boron trifluoride is trigonal planar in shape. Which statement justifies this difference in their structure?

Question

A.
The lone pair of electrons in boron trifluoride increases the bond angle between the bonding pairs, giving it a planar structure.
B.
The lone pair of electrons in boron trifluoride decreases the bond angle between the bonding pairs, giving it a planar structure.
C.
The lone pair of electrons in boron trifluoride stabilizes the bond angle between the bonding pairs, giving it a planar structure.
D.
The lone pair of electrons in ammonia increases the bond angle between the bonding pairs, giving it a pyramidal structure.
E.
The lone pair of electrons on the nitrogen in ammonia decreases the bond angle between the bonding pairs, giving it a pyramidal structure.

Answer 1

The correct answer is:

D. The lone pair of electrons in ammonia increases the bond angle between the bonding pairs, giving it a pyramidal structure.

Here's a brief explanation of the relevant concepts:

In ammonia (NH₃), the nitrogen atom has one lone pair of electrons, which repels the bonding pairs of electrons (the bonds to the hydrogen atoms) more than bonding pairs repel each other. This leads to a pyramidal structure where the bond angles are slightly less than 109.5° (the typical angle in a tetrahedral arrangement).

In contrast, boron trifluoride (BF₃) does not have any lone pairs on the boron atom; instead, it has three bonding pairs. This arrangement places the fluorine atoms in a trigonal planar configuration with bond angles of 120°, as there is no lone pair to alter the geometry.

So the key difference is the presence of the lone pair of electrons in ammonia, which gives it a pyramidal shape.