A Lewis dot structure is a diagram used in chemistry to visually represent the valence electrons of a molecule, showing how atoms bond together and where non-bonding electrons reside. This structure provides a two-dimensional map of electron arrangement, which is foundational for predicting a molecule’s properties and three-dimensional shape. Ammonia (\(\text{NH}_3\)) is a common molecule with a pungent odor, used widely in the production of fertilizers and household cleaners.
Calculating Total Valence Electrons for \(\text{NH}_3\)
The first step in drawing any Lewis structure is to determine the total number of valence electrons available in the molecule. The number of valence electrons an atom contributes is based on its group number on the periodic table. Nitrogen (\(\text{N}\)) is found in Group 15, possessing five valence electrons, and Hydrogen (\(\text{H}\)) is in Group 1, contributing a single valence electron per atom. Since the ammonia molecule contains one nitrogen atom and three hydrogen atoms, the total count is calculated as \(5 + (3 \times 1)\). This results in a total of eight valence electrons, corresponding to four electron pairs that must be arranged around the central atom.
Step-by-Step Construction of the Lewis Structure
The nitrogen atom is identified as the central atom in ammonia because hydrogen atoms can only form one bond and are always terminal in a molecular structure. Single covalent bonds are drawn between the central nitrogen atom and each of the three surrounding hydrogen atoms.
Each single bond represents a shared pair of two electrons, meaning that \(3 \times 2 = 6\) of the total eight valence electrons have been used to form the three \(\text{N}-\text{H}\) bonds. The remaining \(8 – 6 = 2\) valence electrons must be placed on the atoms as lone pairs to satisfy the octet rule. Hydrogen atoms are stable with only two electrons, so the remaining two electrons must be placed on the central nitrogen atom. Placing the final two electrons as a single lone pair on the nitrogen atom completes the Lewis structure. This placement gives the nitrogen atom a full octet, totaling eight electrons (six bonding and two non-bonding).
Understanding the Final Structure and Molecular Geometry
The final Lewis structure for \(\text{NH}_3\) shows the central nitrogen atom is surrounded by four groups of electrons: three bonding pairs and one lone pair. This arrangement of electron groups is directly responsible for the molecule’s three-dimensional shape, which is determined by the Valence Shell Electron Pair Repulsion (VSEPR) theory. VSEPR theory states that electron groups around a central atom will arrange themselves to minimize repulsion. The four electron groups try to occupy the corners of a tetrahedron, giving ammonia a tetrahedral electron geometry.
However, the molecular geometry, which only considers the position of the atoms, is different due to the presence of the lone pair. The lone pair exerts a greater repulsive force than the bonding pairs, pushing the three hydrogen atoms closer together. This increased repulsion compresses the \(\text{H}-\text{N}-\text{H}\) bond angles from the ideal tetrahedral angle of \(109.5^\circ\) down to approximately \(107^\circ\). The resulting arrangement of the three hydrogen atoms below the nitrogen atom, with the lone pair sitting at the top, gives the ammonia molecule its characteristic trigonal pyramidal shape.