Question

Explain from structural considerations why the dipole moment of the \(\mathrm{NH}_{3}\) molecule is lower than that of the \(\mathrm{H}_{2} \mathrm{O}\) molecule.

Short answer

The dipole moment of H2O is higher than that of NH3 because water's bent shape and oxygen's high electronegativity result in a stronger separation of charges than the trigonal pyramidal shape and lower electronegativity of nitrogen in ammonia.

Step-by-step solution

Understanding Molecular Geometry of NH3

The ammonia molecule (NH3) has a trigonal pyramidal shape with the nitrogen atom at the apex and the three hydrogen atoms at the corners of a trigonal base. The lone pair of electrons on the nitrogen atom contributes to the polarity but the symmetry reduces the net dipole moment.

Understanding Molecular Geometry of H2O

The water molecule (H2O) has a bent shape due to the two lone pairs of electrons on the oxygen atom. These lone pairs strongly influence the molecule's polarity, along with the oxygen atom's higher electronegativity compared to hydrogen.

Comparing Electronegativity and Molecular Geometry

Because of oxygen’s higher electronegativity in H2O, and the bent shape enhancing the separation of charge, water has a higher dipole moment. In contrast, NH3’s trigonal pyramidal shape with a less electronegative nitrogen leads to a smaller separation of charges and thus a lower dipole moment.

Contribution of Lone Electron Pairs

The lone pairs in NH3 and H2O both contribute to each molecule's dipole moment, but the effect is more pronounced in H2O due to its geometry, which does not cancel the dipole moments as effectively as in NH3.

Key concepts

Molecular Geometry

Molecular geometry plays a central role in determining the physical and chemical properties of a molecule, including its polarity and hence the molecular dipole moment. These geometries arise from the VSEPR (Valence Shell Electron Pair Repulsion) theory which suggests that the shape of a molecule is largely determined by the repulsion between the valence electron pairs surrounding the central atom.

For instance, in ammonia (\textbf{NH}\(_3\)), the geometry is trigonal pyramidal because of the three hydrogen atoms bonded to the nitrogen and one pair of non-bonding electrons (lone pair). This lone pair pushes down the hydrogen atoms, resulting in a pyramid-like structure with a non-linear arrangement, which contributes to a net dipole moment. However, the molecular geometry of water (\textbf{H}\(_2\)O) is bent or angular, where the two hydrogen atoms and two lone pairs are around the oxygen atom. This arrangement leads to an asymmetrical distribution of charges and a stronger dipole moment in water compared to ammonia.

Electronegativity

Understanding Electronegativity

Electronegativity is a measure of an atom's ability to attract and hold onto electrons within a chemical bond. The greater the difference in electronegativity between bonded atoms, the more polarized the electron distribution is, and the larger the dipole moment of the molecule.

In a water molecule, the oxygen atom is highly electronegative, attracting the electrons more than the hydrogen atoms. This creates a significant difference in charge across the molecule, enhancing its polarity. The nitrogen atom in ammonia, although electronegative, is not as strongly so as oxygen. Consequently, even though the geometry of both molecules is asymmetric, the larger difference in electronegativity between oxygen and hydrogen in water results in a stronger dipole moment than between nitrogen and hydrogen in ammonia.

Lone Pair Electrons

The Role of Lone Pair Electrons in Molecular Polarity

Lone pair electrons are valence electrons that are not shared with other atoms and are located on the central atom. These electrons contribute significantly to molecular polarity because they occupy more space than bonding electron pairs, leading to changes in the molecular geometry that can increase asymmetry in the electron distribution.

In our comparison, the lone pair electrons on nitrogen in ammonia push the hydrogen atoms downwards, causing a separation of charges but with some symmetry still intact. In water, however, the two lone pairs create a more pronounced bent shape, leading to a stronger asymmetrical charge distribution and, thus, a larger dipole moment. This highlights the crucial influence lone pairs have on the geometry and dipole moments of molecules.