Enthalpy Calculator
Enthalpy change (ΔH) is the heat a reaction or process exchanges at constant pressure. Compute it four ways: from measured heat, from standard formation enthalpies (Hess), from heating a mass (m·c·ΔT), or from a phase change (m·L). See the worked steps, a per mole value, and the energy in kJ, kcal and BTU.
Formula
Worked example
Combustion of methane, CH₄ + 2O₂ → CO₂ + 2H₂O. Products: ΔHf(CO₂) = −393.5 and 2·ΔHf(H₂O) = −571.6. Reactants: ΔHf(CH₄) = −74.6 and 2·ΔHf(O₂) = 0. ΔH = (−393.5 − 571.6) − (−74.6 + 0) = −965.1 − (−74.6) = −890.5 kJ, strongly exothermic. Heating 1 kg of water from 20 to 100 °C instead gives ΔH = 1 × 4.184 × 80 = 334.7 kJ absorbed.
What enthalpy change measures
Enthalpy (H) is defined as the internal energy of a system plus the product of its pressure and volume, H = U + PV. We rarely care about the absolute enthalpy of a substance; what matters is the enthalpy change, ΔH, during a reaction or physical process. At constant pressure the work done against the atmosphere is automatically accounted for, so the enthalpy change equals the heat exchanged: ΔH = q_p. A negative ΔH marks an exothermic process that releases heat and warms the surroundings, while a positive ΔH marks an endothermic process that absorbs heat. Because most laboratory and biological chemistry happens in open vessels at roughly constant atmospheric pressure, ΔH is the energy quantity chemists reach for most often.
Four ways to find ΔH
This calculator covers the four routes that appear in real problems. The measured heat method takes the heat exchanged at constant pressure directly as ΔH, and accepts kJ, J, kcal, cal or BTU so you can match your meter or table. The Hess method sums standard formation enthalpies of products minus reactants, the standard chemistry approach. The sensible heat method uses ΔH = m·c·ΔT to find the heat to warm or cool a mass through a temperature change, the workhorse equation behind HVAC loads, calorimetry and cooking. The phase change method uses ΔH = m·L for melting, freezing, boiling or condensing, where temperature holds steady while latent heat flows. Every result is reported in kJ, kcal and BTU at once, and an optional per mole toggle divides by the moles processed so you can compare to tabulated molar enthalpies.
Hess's law and standard formation enthalpies
Enthalpy is a state function, meaning its change depends only on the starting and ending states and not on the route between them. Hess's law follows directly: the enthalpy of an overall reaction equals the sum of the enthalpies of any set of steps that add up to it. The most convenient steps use standard enthalpies of formation, ΔHf°, the heat change when one mole of a compound forms from its elements in their standard states. The reaction enthalpy is then ΔH = Σ ΔHf°(products) − Σ ΔHf°(reactants), with each value multiplied by its stoichiometric coefficient. Elements in their standard state, such as O₂ gas or solid carbon, have ΔHf° = 0 by definition, which is why they often drop out of the sum.
Sensible heat, latent heat and the role of units
Sensible heat changes the temperature of a substance without changing its phase, and follows q = m·c·ΔT, where c is the specific heat capacity. Water has an unusually high c of 4.184 kJ/(kg·K), which is why it stores and moves so much energy in heating systems and oceans. Latent heat, by contrast, drives a phase change at constant temperature: melting and boiling absorb it, freezing and condensing release it, and the amount is q = m·L. Water needs 334 kJ/kg to melt and a much larger 2257 kJ/kg to boil. Because lab data, appliance ratings and engineering tables mix kilojoules, calories and BTU, the calculator converts your input from any of those units and reports the answer in all three, so a 100 BTU/h coil and a 2000 calorie diet land on the same kJ scale.
Enthalpy versus internal energy
Internal energy change ΔU is the heat exchanged at constant volume, where no expansion work is done, whereas ΔH is the heat exchanged at constant pressure, where the system may push back the surroundings as it expands or contracts. The two are linked by ΔH = ΔU + Δ(PV); for reactions involving ideal gases this simplifies to ΔH = ΔU + Δn·R·T, where Δn is the change in the number of moles of gas. When no gas moles change, ΔH and ΔU are nearly equal. The difference matters most for reactions that produce or consume gas, such as combustion, where the volume work can shift the measured heat by a few kilojoules. This calculator reports ΔH, the constant-pressure heat, which is the quantity tabulated for almost all thermochemical data.
Standard enthalpies of formation and key constants
| Substance | Property | Value |
|---|---|---|
| Carbon dioxide (CO₂, g) | ΔHf° | −393.5 kJ/mol |
| Water (H₂O, l) | ΔHf° | −285.8 kJ/mol |
| Water (H₂O, g) | ΔHf° | −241.8 kJ/mol |
| Methane (CH₄, g) | ΔHf° | −74.6 kJ/mol |
| Glucose (C₆H₁₂O₆, s) | ΔHf° | −1273.3 kJ/mol |
| Liquid water | Specific heat c | 4.184 kJ/(kg·K) |
| Air (dry) | Specific heat c | 1.005 kJ/(kg·K) |
| Water, melting | Latent heat L | 334 kJ/kg |
| Water, boiling | Latent heat L | 2257 kJ/kg |
Formation enthalpies in kJ/mol at 298 K; specific and latent heats for the sensible and phase modes. Elements in their standard state are 0 by definition.
Frequently asked questions
Does a negative ΔH mean the reaction is exothermic?
Yes. A negative enthalpy change means the system releases heat to its surroundings, so the reaction is exothermic and the surroundings warm up. A positive ΔH means the system absorbs heat, so the reaction is endothermic and the surroundings cool down.
Why does ΔH equal the heat of the reaction?
At constant pressure, the first law gives ΔH = q_p, the heat exchanged. The pressure-volume work of any expansion or contraction is already folded into the definition H = U + PV, so the measured heat at constant pressure is exactly the enthalpy change, no separate work term is needed.
How do I handle stoichiometric coefficients?
Multiply each substance’s standard formation enthalpy by its coefficient in the balanced equation before adding it to the sum. For example, 2 H₂O contributes 2 × (−285.8) = −571.6 kJ. Enter each already-scaled value in the products or reactants list.
When should I use the sensible heat (m·c·ΔT) mode?
Use it whenever you are heating or cooling a substance without changing its phase, for example warming water, sizing a heating coil, or a calorimetry experiment. Enter the mass, the specific heat capacity (4.184 kJ/(kg·K) for liquid water), and the start and end temperatures. The result is the heat absorbed at constant pressure, which equals ΔH.
How do I calculate the enthalpy of a phase change?
Use ΔH = m·L, the phase change mode. Multiply the mass by the latent heat of the transition: 334 kJ/kg to melt ice or 2257 kJ/kg to boil water, for example. Melting and boiling absorb heat (positive ΔH); freezing and condensing release it (negative ΔH). The temperature stays constant during the change.