Boiling Point Calculator
Calculate the boiling point of a substance at any pressure using the Clausius-Clapeyron equation. Select a common substance or enter custom values to find how boiling temperature changes with pressure.
ln(P₁/P₂) = −(ΔHvap / R) × (1/T₁ − 1/T₂)
The Clausius-Clapeyron equation assumes ΔHvap is constant over the temperature range. Results are most accurate near the normal boiling point.
What Is the Boiling Point?
The boiling point of a substance is the temperature at which its vapor pressure equals the external pressure surrounding the liquid. At this temperature, the liquid undergoes a phase transition from liquid to gas, forming bubbles of vapor throughout the bulk of the liquid rather than just at the surface.
Unlike evaporation, which occurs at any temperature from the liquid's surface, boiling is a vigorous process that happens throughout the entire liquid when sufficient energy is supplied. For water at standard atmospheric pressure (1 atm or 101.325 kPa), the boiling point is 100°C (212°F or 373.15 K).
It is important to understand that the boiling point is not a fixed property of a substance in isolation -- it depends on the surrounding pressure. Reducing the pressure lowers the boiling point, and increasing the pressure raises it. This is why water boils at lower temperatures at high altitudes, where atmospheric pressure is reduced.
What Determines the Boiling Point of a Substance?
Several factors influence a substance's boiling point:
- Intermolecular forces: Substances with stronger intermolecular forces (hydrogen bonding, dipole-dipole interactions, London dispersion forces) require more energy to overcome these attractions, resulting in higher boiling points. Water, with its extensive hydrogen bonding, has a notably high boiling point for its molecular weight.
- Molecular weight: Within a family of similar compounds, heavier molecules generally have higher boiling points because they have greater London dispersion forces.
- Molecular shape: Compact, spherical molecules have lower boiling points than elongated ones of similar weight, because linear molecules have more surface area for intermolecular contact.
- Polarity: Polar molecules typically have higher boiling points than nonpolar molecules of similar size due to dipole-dipole interactions.
- External pressure: Higher ambient pressure forces more energy to be required for molecules to escape into the gas phase, raising the boiling point. Lower pressure has the opposite effect.
The Clausius-Clapeyron Equation Explained
The Clausius-Clapeyron equation is a fundamental thermodynamic relationship that describes how the vapor pressure of a substance changes with temperature. Named after Rudolf Clausius and Benoit Paul Emile Clapeyron, it provides a way to calculate the boiling point of a substance at any given pressure, provided you know its boiling point at a reference pressure and its enthalpy of vaporization.
Where:
- P₁ and P₂ are the vapor pressures at temperatures T₁ and T₂
- T₁ and T₂ are the absolute temperatures (in Kelvin)
- ΔHvap is the molar enthalpy (heat) of vaporization in J/mol
- R is the universal gas constant = 8.314 J/(mol·K)
- ln is the natural logarithm
This equation assumes that the enthalpy of vaporization remains constant over the temperature range considered and that the vapor behaves as an ideal gas. These assumptions hold reasonably well for moderate temperature ranges near the normal boiling point.
Key insight: The Clausius-Clapeyron equation tells us that a plot of ln(P) versus 1/T yields a straight line with a slope of −ΔHvap/R. This linear relationship is the basis for determining enthalpies of vaporization from experimental vapor pressure data.
Step-by-Step: How to Calculate Boiling Point at Different Pressures
Follow these steps to find the boiling point (T₂) at a new pressure (P₂), given a known boiling point (T₁) at a reference pressure (P₁):
Gather your known values
Identify P₁ (reference pressure, often 1 atm), T₁ (known boiling point in Kelvin), P₂ (target pressure), and ΔHvap (heat of vaporization in J/mol). For example, water: P₁ = 1 atm, T₁ = 373.15 K, ΔHvap = 40,660 J/mol.
Convert all units to SI
Make sure temperature is in Kelvin (K = °C + 273.15) and pressures are in the same units. The Clausius-Clapeyron equation uses a ratio P₁/P₂, so as long as both pressures use the same unit, they cancel out.
Rearrange the equation to solve for T₂
Starting from ln(P₁/P₂) = −(ΔHvap/R) × (1/T₁ − 1/T₂), solve for 1/T₂:
1/T₂ = 1/T₁ − (R / ΔHvap) × ln(P₁/P₂)
Plug in your values and compute
For water at P₂ = 0.5 atm: 1/T₂ = 1/373.15 − (8.314/40660) × ln(1/0.5) = 0.002680 − 0.0001418 = 0.002538. Therefore T₂ = 1/0.002538 = 394.0... Wait -- let us recalculate more carefully: 1/373.15 = 0.002680, (8.314/40660) = 0.0002045, ln(2) = 0.6931, so the correction = 0.0002045 × 0.6931 = 0.0001418. Thus 1/T₂ = 0.002680 + 0.0001418 = 0.002822 (note the sign), giving T₂ = 354.5 K = 81.4°C.
Convert back to desired units
Convert the result from Kelvin to Celsius (°C = K − 273.15) or Fahrenheit (°F = K × 9/5 − 459.67) as needed. In our example, water boils at about 81.4°C (178.5°F) at 0.5 atm.
Boiling Points of Common Substances
The table below shows the normal boiling points (at 1 atm) and heats of vaporization for commonly used substances:
| Substance | Formula | Boiling Point (°C) | Boiling Point (K) | ΔHvap (J/mol) |
|---|---|---|---|---|
| Water | H₂O | 100.00 | 373.15 | 40,660 |
| Ethanol | C₂H₅OH | 78.29 | 351.44 | 38,560 |
| Methanol | CH₃OH | 64.70 | 337.85 | 35,210 |
| Acetone | C₃H₆O | 56.05 | 329.20 | 31,300 |
| Benzene | C₆H₆ | 80.10 | 353.25 | 30,720 |
| Diethyl Ether | C₄H₁₀O | 34.60 | 307.75 | 26,520 |
| Chloroform | CHCl₃ | 61.20 | 334.35 | 29,240 |
| Hexane | C₆H₁₄ | 69.00 | 342.15 | 28,850 |
Effect of Pressure on Boiling Point
Pressure has a direct and significant effect on the boiling point of any liquid. The relationship is governed by the Clausius-Clapeyron equation, and understanding it has profound practical implications.
Lower pressure = lower boiling point
At higher altitudes, atmospheric pressure decreases. For every 300 meters (roughly 1,000 feet) of elevation gain, atmospheric pressure drops by about 3.5%. This means water at the summit of Mount Everest (approximately 8,849 meters, where pressure is about 0.33 atm) boils at approximately 70°C (158°F) -- too cool to cook many foods properly.
Higher pressure = higher boiling point
A pressure cooker works by sealing steam inside the vessel, raising the internal pressure to about 1.7-2.0 atm. At 2 atm, water boils at about 121°C (250°F), which dramatically speeds up cooking times and is also the basis for autoclaving (sterilization) in medical and laboratory settings.
Why Does Salt Increase Water's Boiling Point?
When salt (or any non-volatile solute) is dissolved in water, it causes a phenomenon known as boiling point elevation. This is a colligative property, meaning it depends on the number of dissolved particles rather than their identity.
The explanation lies in vapor pressure lowering: dissolved salt particles occupy space at the liquid's surface, reducing the number of water molecules that can escape into the gas phase. With lower vapor pressure, a higher temperature is needed for the vapor pressure to reach atmospheric pressure -- hence the boiling point rises.
The boiling point elevation is calculated using:
Where i is the van't Hoff factor (number of particles the solute dissociates into; for NaCl, i = 2), Kb is the ebullioscopic constant of the solvent (0.512 °C·kg/mol for water), and m is the molality of the solution.
For typical cooking concentrations of salt (about 1-2 tablespoons per liter), the boiling point of water increases by only about 0.5-1°C. While this is scientifically real, it has a negligible effect on cooking times. The primary reason chefs add salt to boiling water is for flavor, not to meaningfully change the boiling point.
Real-World Applications
Cooking at High Altitude
At altitudes above 600 meters (2,000 feet), the reduced atmospheric pressure significantly lowers the boiling point of water. In Denver, Colorado (1,609 m), water boils at about 95°C (203°F). In La Paz, Bolivia (3,640 m), it boils at roughly 87°C (189°F). This means food takes longer to cook in boiling water at high altitudes because the water temperature is lower. Recipes often recommend increasing cooking times or using pressure cookers to compensate.
Pressure Cookers
A pressure cooker is a sealed vessel that traps steam, increasing the internal pressure above atmospheric. At approximately 15 psi above atmospheric pressure (about 2 atm total), water boils at around 121°C (250°F). This higher temperature cooks food significantly faster -- often 3 to 10 times quicker than conventional boiling. Pressure cooking also preserves more nutrients because of shorter cooking times and a sealed environment.
Distillation
Distillation relies on the different boiling points of components in a liquid mixture to separate them. In petroleum refining, crude oil is heated and its components (gasoline, diesel, kerosene, etc.) are separated based on their boiling point ranges. Vacuum distillation uses reduced pressure to lower boiling points, which is essential for separating heat-sensitive compounds that would decompose at their normal boiling temperatures.
Vacuum Evaporation in Food Processing
In the production of concentrated fruit juices, condensed milk, and sugar, vacuum evaporation is used to boil off water at reduced temperatures (often 40-60°C). This preserves the flavor, color, and nutritional value of the product that would be damaged by higher temperatures at normal pressure.
Industrial Chemical Processing
Understanding how boiling points change with pressure is critical in designing chemical reactors, heat exchangers, and separation equipment. Engineers routinely use the Clausius-Clapeyron equation to predict phase behavior and optimize industrial processes.
Frequently Asked Questions
The normal boiling point (also called the standard boiling point) is the temperature at which a liquid boils when the external pressure is exactly 1 atmosphere (101.325 kPa). For example, the normal boiling point of water is 100°C (212°F, 373.15 K). When people refer to "the boiling point" of a substance without specifying a pressure, they typically mean the normal boiling point.
Yes. If you reduce the pressure sufficiently, water can boil at room temperature or even colder. At a pressure of about 0.023 atm (2.34 kPa or 17.5 mmHg), water boils at 20°C. This is the principle behind vacuum chambers used in laboratories and industrial processes. In the vacuum of space, exposed water would boil almost instantly due to the near-zero pressure.
The Clausius-Clapeyron equation provides good approximations near the normal boiling point, typically within 1-2% accuracy for modest pressure changes. However, its accuracy decreases for very large pressure or temperature ranges because it assumes that ΔHvap remains constant (which it does not -- it decreases as temperature increases toward the critical point). For high-precision work or extreme conditions, more sophisticated equations of state (like the Antoine equation or Peng-Robinson equation) are used.
Evaporation occurs at any temperature and only from the surface of a liquid. Molecules at the surface with enough kinetic energy can escape into the gas phase. Boiling occurs at a specific temperature (the boiling point) throughout the entire volume of the liquid, when the vapor pressure equals the external pressure. Boiling is much more vigorous than evaporation and requires continuous heat input to sustain. A puddle of water evaporates on a warm day, but it does not boil because the temperature is well below 100°C at 1 atm.
Ethanol (78.3°C) and methanol (64.7°C) have lower boiling points than water (100°C) primarily because their intermolecular hydrogen bonding is weaker than water's. Water has two hydrogen bond donor sites (O-H) and two lone pairs that can accept hydrogen bonds, creating an extensive three-dimensional hydrogen bonding network. Alcohols have only one O-H group per molecule, resulting in fewer and weaker hydrogen bonds. With less intermolecular attraction to overcome, alcohol molecules escape into the gas phase at lower temperatures.
At higher altitudes, reduced atmospheric pressure lowers the boiling point of water. Since water boils at a lower temperature, food cooks more slowly because it is being heated to a lower maximum temperature. As a general rule, for every 150 meters (500 feet) above 600 meters (2,000 feet), add about 1 minute of cooking time per 5 minutes specified in a recipe. For baking, adjustments to leavening, liquid, sugar, and oven temperature may also be needed. Using a pressure cooker is the most effective way to restore normal cooking times at high altitude.
As pressure increases, the boiling point continues to rise -- but only up to a point. Every substance has a critical point, a specific temperature and pressure above which the distinction between liquid and gas ceases to exist. For water, the critical point is 374°C at 218 atm (22.1 MPa). Above this critical temperature, no amount of pressure can condense the gas into a liquid. The substance instead becomes a supercritical fluid with properties intermediate between those of a liquid and a gas.