Properties

Vapor pressure is the pressure exerted by a vapor in dynamic equilibrium with its liquid at a given temperature. Not all molecules in a liquid have the same kinetic energy — they follow a Maxwell–Boltzmann distribution. At any instant, a fraction of surface molecules have enough energy to overcome intermolecular attractions and escape into the gas phase. In a closed container, evaporation and condensation reach a balance, and the resulting equilibrium vapor pressure depends entirely on temperature and IMF strength. Liquids with weak IMFs (e.g., diethyl ether, $\ce{C2H5OC2H5}$) are volatile and have high vapor pressures at room temperature. Liquids with strong hydrogen bonding (e.g., water, $\ce{H2O}$) are less volatile and have lower vapor pressures. Raising temperature increases the fraction of molecules with sufficient energy to escape, so vapor pressure always increases with temperature — a relationship captured quantitatively by the Clausius–Clapeyron equation.

Surface tension ($\gamma$) arises because molecules at the surface of a liquid experience a net inward force. A molecule in the bulk liquid is surrounded on all sides by neighbours and experiences balanced attractive forces. A molecule at the surface, however, has no neighbours above it in the vapour phase, so it is pulled inward and sideways. The liquid therefore minimises its surface area, behaving as though it has a thin elastic 'skin'. Surface tension is measured in N/m (or equivalently J/m²). Water has an unusually high surface tension (≈ 0.072 N/m at 25 °C) due to its extensive hydrogen-bond network, allowing small insects such as water striders to walk on its surface. Surfactants such as soap (sodium stearate, $\ce{C17H35COONa}$) disrupt hydrogen bonding at the surface and dramatically reduce surface tension. In general, stronger IMFs → higher surface tension.

Viscosity ($\eta$) is a measure of a liquid's resistance to flow. When one layer of liquid slides past an adjacent layer, intermolecular attractions must be partially overcome; the stronger those attractions, the more slowly the liquid flows. Viscosity is measured in pascal-seconds (Pa·s) or the older unit centipoise (cP). Water at 20 °C has $\eta \approx 1.0$ cP, while honey has $\eta \approx 10{,}000$ cP. Two structural factors determine viscosity: (1) IMF strength — stronger forces increase viscosity; and (2) molecular shape and size — long, chain-like molecules can become entangled (e.g., motor oil polymers), greatly increasing viscosity. Temperature has a strong inverse effect on viscosity: heating a liquid increases molecular kinetic energy, allowing layers to slide more easily, so viscosity decreases. This is why engine oil flows more readily when the engine warms up. Glycerol ($\ce{HOCH2CH(OH)CH2OH}$) is highly viscous because its three –OH groups form an extensive network of hydrogen bonds.

A unifying principle runs through all three properties: the stronger the intermolecular forces in a liquid, the lower the vapor pressure, the higher the surface tension, and the higher the viscosity. For example, comparing ethanol ($\ce{C2H5OH}$, capable of hydrogen bonding) with dimethyl ether ($\ce{CH3OCH3}$, same molecular formula but only dipole–dipole forces): ethanol has a lower vapor pressure (5.95 kPa vs 530 kPa at 20 °C), higher surface tension, and higher viscosity. Within a homologous series such as the alkanes (only London dispersion forces), all three properties increase with molar mass because larger, more polarisable molecules have stronger dispersion forces. Recognising the type and relative strength of IMFs from a structural formula is therefore the single most powerful predictive tool for comparing liquid properties.

• P_1,\; P_2 — Vapor pressures (Pa or atm): Equilibrium vapor pressure of the liquid at temperatures T₁ and T₂ respectively.
• \Delta H_{\text{vap}} — Enthalpy of vaporization (J mol⁻¹): Energy required to convert one mole of liquid to vapor at constant pressure; reflects IMF strength.
• R — Universal gas constant (8.314 J mol⁻¹ K⁻¹): Relates energy scale to temperature scale in thermodynamic equations.
• T_1,\; T_2 — Absolute temperatures (K): Temperatures in kelvin at which the corresponding vapor pressures P₁ and P₂ are measured.

Problem. The vapor pressure of water is 2.338 kPa at 20 °C and its enthalpy of vaporization is 44,000 J mol⁻¹. Estimate the vapor pressure of water at 35 °C. (R = 8.314 J mol⁻¹ K⁻¹)

Answer. The vapor pressure of water at 35 °C is approximately 5.65 kPa.

• Intermolecular Forces — Vapor pressure, surface tension, and viscosity are all direct consequences of the type and strength of IMFs; this concept is the essential prerequisite.
• Kinetic Molecular Theory and Maxwell–Boltzmann Distribution — The Maxwell–Boltzmann energy distribution explains why only a fraction of surface molecules escape to form vapour, underpinning the concept of vapor pressure.
• Phase Diagrams and Changes of State — Vapor pressure curves form the liquid–gas boundary on a phase diagram; understanding vapor pressure is essential for reading and interpreting phase diagrams.
• Enthalpy of Vaporization — ΔH_vap appears directly in the Clausius–Clapeyron equation and quantifies the energy cost of overcoming IMFs when a liquid evaporates.
• Molecular Polarity and Dipole Moments — Polarity determines whether a molecule engages in dipole–dipole interactions or only London dispersion forces, directly affecting all three liquid properties discussed here.