Vapor Pressure of Water at Various Temperatures
| Temperature (°C) | Temperature (K) | Vapor Pressure (mmHg) | Vapor Pressure (kPa) | Vapor Pressure (atm) |
|---|---|---|---|---|
| 0 | 273.15 | 4.58 | 0.611 | 0.00603 |
| 5 | 278.15 | 6.54 | 0.872 | 0.00861 |
| 10 | 283.15 | 9.21 | 1.228 | 0.01212 |
| 15 | 288.15 | 12.79 | 1.705 | 0.01683 |
| 17 | 290.15 | 14.53 | 1.937 | 0.01912 |
| 18 | 291.15 | 15.48 | 2.064 | 0.02037 |
| 19 | 292.15 | 16.48 | 2.197 | 0.02169 |
| 20 | 293.15 | 17.54 | 2.338 | 0.02308 |
| 21 | 294.15 | 18.65 | 2.487 | 0.02454 |
| 22 | 295.15 | 19.83 | 2.644 | 0.02609 |
| 23 | 296.15 | 21.07 | 2.81 | 0.02772 |
| 24 | 297.15 | 22.38 | 2.984 | 0.02945 |
| 25 | 298.15 | 23.76 | 3.169 | 0.03126 |
| 26 | 299.15 | 25.21 | 3.361 | 0.03317 |
| 27 | 300.15 | 26.74 | 3.565 | 0.03518 |
| 28 | 301.15 | 28.35 | 3.78 | 0.0373 |
| 29 | 302.15 | 30.04 | 4.005 | 0.03953 |
| 30 | 303.15 | 31.82 | 4.243 | 0.04187 |
| 35 | 308.15 | 42.18 | 5.623 | 0.0555 |
| 40 | 313.15 | 55.32 | 7.376 | 0.07279 |
| 45 | 318.15 | 71.88 | 9.583 | 0.09458 |
| 50 | 323.15 | 92.51 | 12.33 | 0.1217 |
| 55 | 328.15 | 118 | 15.73 | 0.1553 |
| 60 | 333.15 | 149.4 | 19.92 | 0.1966 |
| 65 | 338.15 | 187.5 | 25 | 0.2467 |
| 70 | 343.15 | 233.7 | 31.16 | 0.3075 |
| 75 | 348.15 | 289.1 | 38.55 | 0.3804 |
| 80 | 353.15 | 355.1 | 47.34 | 0.4672 |
| 85 | 358.15 | 433.6 | 57.81 | 0.5705 |
| 90 | 363.15 | 525.8 | 70.1 | 0.6918 |
| 95 | 368.15 | 633.9 | 84.53 | 0.8341 |
| 100 | 373.15 | 760 | 101.33 | 1 |
Boiling occurs when vapor pressure equals external pressure: at sea level (760 mmHg) that's 100 °C, but at altitude the boiling point drops — Denver (~630 mmHg) boils water near 95 °C, and the summit of Everest (~250 mmHg) near 70 °C. The Clausius–Clapeyron equation, ln(P₂/P₁) = (ΔH_vap/R)(1/T₁ − 1/T₂), describes the temperature dependence; for water ΔH_vap ≈ 40.7 kJ/mol between 25 °C and 100 °C. Linear interpolation between adjacent rows is fine for temperatures within ~5 °C of a tabulated value, but for larger gaps use Clausius–Clapeyron to avoid systematic error. Sources: CRC Handbook of Chemistry and Physics and NIST Chemistry WebBook.
Frequently Asked Questions
Why is the vapor pressure of water important for gas collection?
When you collect a gas by water displacement, the gas in the inverted bottle is mixed with water vapor saturated at the bath temperature. Dalton's law of partial pressures says P_total = P_gas + P_H₂O, so to get the dry-gas pressure for a PV = nRT calculation you subtract the vapor pressure from the barometric reading: P_dry_gas = P_total − P_H₂O. At 25 °C that's 23.76 mmHg, which is about 3% of atmospheric pressure — small but not negligible. Forgetting this correction is one of the most common errors in introductory gas-law problems.
Why does water boil at 100 °C?
Only at 1 atm. Boiling happens when vapor pressure equals the external pressure pushing down on the liquid surface, allowing bubbles of vapor to form throughout the bulk. At sea level (760 mmHg), water's vapor pressure hits that mark at exactly 100 °C — by definition, that's how the Celsius scale was originally calibrated. Lower the external pressure and water boils cooler: ~95 °C in Denver, ~70 °C atop Everest. A pressure cooker raises the internal pressure to ~2 atm, which pushes the boiling point to about 121 °C and dramatically speeds up cooking.
How do you use the Clausius-Clapeyron equation with this table?
Clausius–Clapeyron, ln(P₂/P₁) = (ΔH_vap/R)(1/T₁ − 1/T₂), gives you two practical moves. First, with two table rows you can solve for ΔH_vap: using P₁ = 23.76 mmHg at 298 K and P₂ = 760 mmHg at 373 K gives ΔH_vap ≈ 42.6 kJ/mol, close to the accepted 40.7 kJ/mol for the 25–100 °C range. Second, with ΔH_vap and one tabulated point you can predict vapor pressure at any other temperature. Use R = 8.314 J/(mol·K), temperatures in Kelvin, and watch the sign convention on the 1/T term.