Water, a compound of two hydrogen atoms and one oxygen atom, boils at \(100^{\circ}\text{C}\) (\(212^{\circ}\text{F}\)) at sea level. The boiling point of a substance is the temperature at which its liquid form turns into a gas. This temperature is surprisingly high for a molecule with such a small molecular mass of approximately 18 grams per mole. The general rule for molecules of similar size is that they should boil at much lower temperatures. The explanation for this unusual thermal stability lies in the powerful forces that exist between individual water molecules.
The Polar Nature of the Water Molecule
The ability of water to form these strong intermolecular forces begins with its internal structure. The oxygen atom within the \(\text{H}_2\text{O}\) molecule is significantly more electronegative than the two hydrogen atoms. Because of this difference, the oxygen atom pulls the shared electrons closer to itself, creating an unequal sharing of charge within the molecule.
This unequal electron distribution causes the oxygen atom to develop a partial negative charge, often symbolized as \(\delta-\). Conversely, the hydrogen atoms, having their electrons pulled away, each acquire a partial positive charge, symbolized as \(\delta+\). This separation of charge is known as a dipole, making the water molecule highly polar. The molecular geometry of water is also a significant factor, as the molecule is not linear but has a bent or V-shape with a bond angle of about \(104.5^{\circ}\).
This bent structure is due to the two unshared pairs of electrons on the oxygen atom, which repel the hydrogen atoms. If the molecule were linear, the two dipoles would pull in opposite directions and cancel each other out, resulting in a nonpolar molecule. However, the asymmetrical, bent shape means the partial charges do not cancel, giving the entire water molecule a net direction of polarity. This structural arrangement is the prerequisite for the strong forces that hold liquid water together.
Hydrogen Bonding: The High Energy Requirement
The polarity of water molecules allows them to form an especially strong type of intermolecular attraction known as a hydrogen bond. A hydrogen bond forms when the partially positive hydrogen atom (\(\delta+\)) of one water molecule is attracted to the partially negative oxygen atom (\(\delta-\)) of an adjacent water molecule. This attraction is significantly stronger than the typical dipole-dipole forces found in most other polar molecules.
The strength of this bond arises because the highly electronegative oxygen atom leaves the hydrogen nucleus almost bare of electron density. This exposes the hydrogen atom’s positive charge, creating a strong localized attraction to the lone pair of electrons on a neighboring oxygen atom. Each water molecule can participate in up to four hydrogen bonds with its neighbors due to its two hydrogen atoms and two lone pairs of electrons on the oxygen atom.
In the liquid state, these individual attractions form an extensive, three-dimensional network of interconnected molecules. For water to boil and transition into a gas, a substantial amount of thermal energy must be supplied to break these numerous and powerful hydrogen bonds. This requirement for significant energy input to overcome the strong forces between molecules directly accounts for the elevated boiling temperature of \(100^{\circ}\text{C}\). The large amount of heat needed to break these bonds is termed the high heat of vaporization.
Comparing Water to Other Molecular Compounds
The anomalous nature of water’s boiling point becomes clear when compared to the hydrides of other elements in the same group of the periodic table, Group 16. These molecules, such as hydrogen sulfide (\(\text{H}_2\text{S}\)), hydrogen selenide (\(\text{H}_2\text{Se}\)), and hydrogen telluride (\(\text{H}_2\text{Te}\)), have similar molecular structures but lack the powerful hydrogen-bonding network. For most compounds, boiling point generally increases as molecular weight and size increase, leading to stronger London dispersion forces.
Following this trend, the boiling points of the Group 16 hydrides increase steadily:
- Hydrogen sulfide boils at \(-60.7^{\circ}\text{C}\).
- Hydrogen selenide boils at \(-41.25^{\circ}\text{C}\).
- Hydrogen telluride boils at approximately \(-2^{\circ}\text{C}\).
Water, being the smallest and lightest of these molecules, should theoretically follow the trend and boil at a temperature somewhere near \(-80^{\circ}\text{C}\) or \(-100^{\circ}\text{C}\). Water’s actual boiling point of \(100^{\circ}\text{C}\) is a deviation of nearly \(200^{\circ}\text{C}\) from the expected value. This dramatic difference confirms that the weak dipole-dipole and London dispersion forces present in these other hydrides are not comparable to the collective strength of the hydrogen bonds in liquid water.