The shape of the ammonia molecule ( NH3) is Tetrahedral bent trigonal pyramidal. Trigonal planar:
The Correct Answer and Explanation is:
The correct answer is trigonal pyramidal.
The shape of the ammonia molecule (NH₃) is determined using the Valence Shell Electron Pair Repulsion (VSEPR) theory. This model predicts the three-dimensional geometry of molecules based on the idea that electron pairs surrounding a central atom will arrange themselves to minimize electrostatic repulsion.
First, we analyze the central atom, which is nitrogen (N). Nitrogen is in Group 15 of the periodic table, meaning it has five valence electrons. It forms single covalent bonds with three hydrogen (H) atoms, each contributing one valence electron. This uses three of nitrogen’s valence electrons for bonding. The remaining two valence electrons on the nitrogen atom form a non-bonding lone pair.
Consequently, the central nitrogen atom is surrounded by four distinct regions of electron density: three bonding pairs (the N-H bonds) and one lone pair. According to VSEPR theory, four electron domains arrange themselves in a tetrahedral electron geometry to maximize the distance between them.
However, the molecular shape or geometry describes the arrangement of only the atoms, not the lone pairs. While the electron pairs have a tetrahedral arrangement, the lone pair is not “seen” in the final molecular shape. With three atoms bonded to a central atom and one lone pair, the resulting shape is a trigonal pyramid. The nitrogen atom sits at the apex of this pyramid, with the three hydrogen atoms forming the triangular base below it.
The presence of the lone pair is crucial. It exerts a greater repulsive force than bonding pairs, pushing the N-H bonds down and closer together. This reduces the H-N-H bond angle from the ideal tetrahedral angle of 109.5° to approximately 107°. This shape is distinct from trigonal planar, which has no lone pair and 120° bond angles, and from a true tetrahedral shape like methane (CH₄), which has four bonding pairs and no lone pairs.
