Understanding Vapor Pressure
Vapor pressure represents the pressure at which a liquid or solid exists in dynamic equilibrium with its vapor phase. Molecules at the surface continuously evaporate and condense; vapor pressure is the pressure at which these rates balance. In a sealed container, this equilibrium establishes itself naturally. The substance behaves differently depending on whether the surrounding pressure exceeds, equals, or falls short of its vapor pressure at that temperature.
Molecules require sufficient kinetic energy to escape from their condensed phase. When temperature rises, more molecules gain this energy, causing vapor pressure to increase exponentially rather than linearly. Conversely, cooling reduces molecular energy and lowers vapor pressure. This relationship holds for all pure substances, though the exact rate of change varies.
The molecular nature of the substance strongly influences its vapor pressure. Water, with its hydrogen bonding, has notably different vapor pressure behavior compared to alcohols or hydrocarbons. At the same temperature, substances with weaker intermolecular forces exhibit higher vapor pressures.
Factors Affecting Water Vapor Pressure
Temperature is the dominant factor controlling vapor pressure. Each degree increase in temperature raises the vapor pressure of water appreciably. At 0°C, water's vapor pressure is approximately 0.61 kPa; at 25°C it reaches 3.17 kPa; and at 100°C it climbs to 101.3 kPa (one atmosphere).
Intermolecular forces determine how readily water molecules escape into the vapor phase. Hydrogen bonds in liquid water are strong enough to keep most molecules in the condensed state at room temperature, yet weak enough to allow significant evaporation. Breaking these bonds requires energy supplied by temperature.
Pressure above the liquid affects how much vapor can exist in equilibrium. In an open system, molecules escape continuously. In a sealed container, vapor accumulates until pressure reaches equilibrium. Increasing external pressure on a liquid suppresses further evaporation.
Below freezing, ice exhibits lower vapor pressure than liquid water at the same temperature because molecules in the solid state have even less freedom to escape. This phenomenon, called sublimation, allows ice to evaporate directly without melting.
Common Vapor Pressure Equations
Five empirical formulas provide accurate vapor pressure estimates across different temperature ranges. The Antoine equation dominates industrial practice, while the Magnus and Tetens approximations perform well for meteorology and general laboratory work. Each formula has specific constants optimized for particular temperature windows.
Antoine Formula:
P = 10^(A − B/(C + T))
Magnus Formula:
P = 0.61094 × exp((17.625 × T)/(T + 243.04))
Tetens Formula (water):
P = 0.61078 × exp((17.27 × T)/(T + 237.3))
Simple Formula:
P = exp(20.386 − 5132/(T + 273))
Buck Formula (water):
P = 0.61121 × exp((18.678 − T/234.5) × T/(257.14 + T))
P— Vapor pressure in kPaT— Temperature in degrees CelsiusA, B, C— Antoine equation constants (values vary depending on the source and temperature range)
Vapor Pressure Data for Water
The following reference values show how dramatically vapor pressure changes across common temperature ranges:
- 0°C: 0.61 kPa (freezing point)
- 10°C: 1.23 kPa
- 20°C: 2.34 kPa (room temperature)
- 30°C: 4.24 kPa
- 50°C: 12.34 kPa
- 75°C: 38.56 kPa
- 100°C: 101.33 kPa (normal boiling point)
- 120°C: 198.32 kPa (above atmospheric pressure)
These values highlight why boiling occurs when vapor pressure equals atmospheric pressure. Below 100°C at sea level, water's vapor pressure remains below one atmosphere, so bubbles of vapor collapse before reaching the surface. At 100°C, vapor pressure matches atmospheric pressure, allowing bubbles to form throughout the liquid and rise freely.
Practical Considerations for Vapor Pressure Work
Several common pitfalls arise when applying vapor pressure data to real-world problems.
- Sealed vs. open systems behave differently — Vapor pressure only develops in closed containers where molecules cannot escape. In an open beaker, water evaporates continuously at any temperature above absolute zero, but vapor pressure never truly equilibrates. Never apply vapor pressure concepts to evaporative cooling or clothesline drying without accounting for continuous molecule loss.
- Formula accuracy varies by temperature range — The Antoine equation (particularly the second variant with coefficients 8.14019, 1810.94, 244.485) excels between 0°C and 100°C but becomes less reliable above 120°C. Magnus and Tetens formulas work well near room temperature but diverge outside the 0–50°C window. Always verify which temperature range your chosen formula covers.
- Dissolved substances alter vapor pressure — The presence of dissolved salts or other solutes lowers water's vapor pressure through colligative effects. Seawater's vapor pressure is measurably lower than pure water at the same temperature. Ignore this effect only for pure water; for solutions, apply Raoult's law or seek substance-specific data.
- Ice sublimation follows different kinetics — Below freezing, ice vapor pressure is always lower than liquid water vapor pressure at the same temperature. The Buck and Tetens formulas include separate coefficients for ice. Using water-liquid coefficients below 0°C will overestimate how readily ice sublimates in a freezer or vacuum chamber.