Ammonia (\(\text{NH}_3\)) is a fundamental molecule used widely in biological systems and industrial processes. Its structure consists of one nitrogen atom bonded to three hydrogen atoms. A lone pair is a pair of valence electrons associated with a single atom that are not involved in forming a covalent bond. The central nitrogen atom in \(\text{NH}_3\) carries exactly one lone pair of electrons. This single pair of non-bonding electrons dictates the molecule’s three-dimensional shape and its chemical reactivity.
Determining the Valence Electrons
The presence of the lone pair is determined by accounting for the valence electrons contributed by each atom. Nitrogen (Group 15) contributes five valence electrons, and each of the three hydrogen atoms (Group 1) contributes one. This results in a total of eight valence electrons available for bonding (\(5 + 3 \times 1\)).
The three single covalent bonds between nitrogen and hydrogen utilize six of these eight electrons. The remaining two electrons are placed on the central nitrogen atom, forming the single lone pair. This arrangement satisfies the octet rule for nitrogen, which now has eight electrons in its valence shell.
How the Lone Pair Defines Molecular Geometry
The three-dimensional arrangement of \(\text{NH}_3\) is predicted by the Valence Shell Electron Pair Repulsion (VSEPR) theory. VSEPR theory dictates that electron pairs around the central nitrogen atom position themselves as far apart as possible to minimize repulsive forces. The nitrogen atom has four electron domains: three bonding pairs and one lone pair. These four domains arrange themselves in a tetrahedral electron geometry.
The molecular geometry, which describes only the position of the atoms, is trigonal pyramidal. The lone pair occupies space but is not an atom, resulting in the nitrogen atom resting at the apex above the three hydrogen atoms. The lone pair exerts a stronger repulsive force on the adjacent bonding pairs than the bonding pairs exert on each other. This repulsion compresses the \(\text{H}-\text{N}-\text{H}\) bond angle from the ideal tetrahedral angle of \(109.5^\circ\) to approximately \(107^\circ\).
Chemical Behavior Resulting from the Lone Pair
The lone pair is responsible for ammonia’s defining chemical characteristics, including its polarity and basicity. Nitrogen is more electronegative than hydrogen, making the three \(\text{N}-\text{H}\) bonds polar. The lone pair enhances this charge imbalance, increasing electron density around the nitrogen nucleus.
Because the trigonal pyramidal shape is asymmetrical, the individual bond dipoles do not cancel out. The vector sum of the bond dipoles and the lone pair’s dipole results in a strong net dipole moment, making \(\text{NH}_3\) highly polar. This polarity allows ammonia to readily dissolve in polar solvents like water.
The most significant chemical role of the lone pair is its ability to act as an electron pair donor. This makes ammonia an effective Lewis base, capable of donating its non-bonding electron pair to form a new bond. Ammonia also acts as a Brønsted-Lowry base by using the lone pair to accept a proton (\(\text{H}^+\)), forming the ammonium ion (\(\text{NH}_4^+\)).