Water's higher boiling point is due to strong hydrogen bonding, while methane's lower boiling point is attributed to the weaker van der Waals forces, particularly London dispersion forces.
The difference in boiling points between water (H₂O) and methane (CH₄) can be attributed to variations in their intermolecular forces. Water exhibits hydrogen bonding, a strong intermolecular force resulting from the significant electronegativity difference between hydrogen and oxygen atoms.
Hydrogen bonding leads to cohesive forces that require more energy to break during the transition from liquid to gas, resulting in a higher boiling point for water at 100 degrees Celsius.
On the other hand, methane relies on weaker van der Waals forces, specifically London dispersion forces, as its predominant intermolecular force. These forces arise from temporary fluctuations in electron distribution, inducing temporary dipoles.
Because London dispersion forces are generally weaker than hydrogen bonding, methane boils at a much lower temperature, approximately -161.5 degrees Celsius.