Calculate the enthalpy change (ΔH) of a chemical reaction using standard enthalpies of formation. Enter the stoichiometric coefficients and ΔHf° values for up to three reactants and three products — you get back the reaction enthalpy (ΔH rxn) in kJ/mol, plus whether the reaction is endothermic or exothermic.
Enthalpy Change (ΔH rxn)
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Sum of Products ΔHf°
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Sum of Reactants ΔHf°
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Reaction Type
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Enthalpy (H) is a thermodynamic quantity that represents the total heat content of a system. It is defined as H = U + pV, where U is internal energy, p is pressure, and V is volume. In chemistry, we typically work with the change in enthalpy (ΔH) rather than absolute values, which tells us how much heat is absorbed or released during a reaction at constant pressure.
The enthalpy change of a reaction is calculated using Hess's Law: ΔH rxn = Σ(n × ΔHf° of products) − Σ(n × ΔHf° of reactants), where n is the stoichiometric coefficient of each substance. Simply multiply each compound's standard enthalpy of formation by its coefficient, sum the products side, sum the reactants side, and subtract.
An exothermic reaction releases heat to the surroundings and has a negative ΔH value (ΔH < 0). An endothermic reaction absorbs heat from the surroundings and has a positive ΔH value (ΔH > 0). Combustion reactions are classic examples of exothermic processes, while dissolving ammonium nitrate in water is a familiar endothermic reaction.
Yes. A negative ΔH indicates that the products have less energy than the reactants, meaning the system releases energy as heat to its surroundings — which is the definition of an exothermic reaction. A positive ΔH means energy is absorbed, making it endothermic.
By definition, the standard enthalpy of formation (ΔHf°) is zero for any element in its most stable form at standard conditions (298 K, 1 atm). Examples include O₂(g), N₂(g), H₂(g), C(graphite), Fe(s), and Na(s). These are the reference states from which all other formation enthalpies are measured.
The standard enthalpy of formation of liquid water H₂O(l) is −285.83 kJ/mol, and for water vapor H₂O(g) it is −241.82 kJ/mol. The difference of about 44 kJ/mol represents the enthalpy of vaporization of water at standard conditions.
ΔHf° is the standard enthalpy of formation — the heat change when one mole of a compound is formed from its elements in their standard states at 298 K and 1 atm. These values are tabulated in chemistry textbooks, the NIST Chemistry WebBook, and standard thermodynamic tables. Common values include CO₂(g) = −393.51 kJ/mol and CH₄(g) = −74.81 kJ/mol.
Enthalpy change is one factor in determining spontaneity, but not the only one. The Gibbs free energy equation ΔG = ΔH − TΔS combines enthalpy (ΔH) and entropy (ΔS) to predict spontaneity. A reaction is spontaneous when ΔG < 0, which can occur even for endothermic reactions if there is a large enough increase in entropy.