Water and ethanol are common liquids, yet they exhibit a significant difference in their boiling points. Water boils at \(100^\circ\text{C}\) at standard atmospheric pressure, while ethanol begins to boil at the notably lower temperature of approximately \(78^\circ\text{C}\). This difference is not a matter of molecular mass, as ethanol molecules are larger and heavier than water molecules. Instead, the disparity in boiling temperatures is a direct result of the varying strengths and extents of the attractive forces that exist between the molecules in each liquid.
Defining Boiling Point and Intermolecular Forces
Boiling is the physical process where a liquid transitions into a gas, which requires molecules to overcome the attractive forces holding them together. These attractions are known as intermolecular forces (IMFs). The boiling point is the temperature at which the energy supplied is sufficient to break these forces, allowing the molecules to escape as vapor. All molecules experience London dispersion forces, which are weak, temporary attractions arising from random electron movement. Polar molecules, like water and ethanol, also experience dipole-dipole attractions.
The strongest type of dipole-dipole interaction is hydrogen bonding. This occurs when a hydrogen atom is covalently bonded to a highly electronegative atom, such as oxygen, nitrogen, or fluorine. This creates a strong partial positive charge on the hydrogen atom, which is then strongly attracted to the lone pair of electrons on a neighboring molecule. Hydrogen bonds are significantly stronger than other IMFs, meaning substances that form them require a larger energy input to reach their boiling point. The presence and strength of hydrogen bonding is the primary factor determining the difference between the boiling points of water and ethanol.
Molecular Structure and Hydrogen Bonding in Water
The water molecule (\(H_2O\)) possesses a bent shape due to the oxygen atom’s two lone pairs of electrons. Oxygen is highly electronegative, strongly pulling the shared electrons toward itself, giving the oxygen atom a large partial negative charge and the hydrogen atoms a significant partial positive charge. This highly polar structure is designed for extensive hydrogen bonding.
Each water molecule contains two hydrogen atoms that act as bond donors and two lone pairs of electrons that act as bond acceptors. This unique geometry allows a single water molecule to participate in a maximum of four strong hydrogen bonds simultaneously with its neighbors. This extensive, three-dimensional network creates a highly cohesive structure in liquid water. Breaking this network requires a substantial amount of thermal energy, which is why water has a relatively high boiling point compared to many other small molecules.
The Limiting Factors in Ethanol’s Structure
Ethanol (\(C_2H_5OH\)) is an alcohol that also contains a hydroxyl (\(–OH\)) group, allowing it to form hydrogen bonds just like water. However, the structure of ethanol is significantly different because it includes an ethyl group (\(C_2H_5\)), which is a non-polar hydrocarbon chain. This non-polar “tail” is unable to participate in hydrogen bonding and limits the overall polarity of the molecule.
The larger ethyl group introduces a steric hindrance that physically limits the close approach and optimal arrangement of molecules needed for a dense hydrogen bond network. More importantly, the ethanol molecule only has one hydrogen atom attached to the oxygen atom that can act as a hydrogen bond donor, in contrast to the two donor hydrogen atoms in a water molecule.
While the oxygen atom in ethanol still has two lone pairs to accept hydrogen bonds, the single donor site means each ethanol molecule can participate in a maximum of three hydrogen bonds, which is one less than water’s four. The presence of the large, non-polar ethyl group means a greater portion of the attractions between ethanol molecules are the weaker London dispersion forces. Consequently, less energy is required to separate the ethanol molecules and turn the liquid into a gas, resulting in a boiling point nearly \(22^\circ\text{C}\) lower than that of water.