Reaction Enthalpy Calculator
Reaction enthalpy (ΔH) quantifies the heat released or absorbed during a chemical reaction at constant pressure, and Hess's Law allows it to be calculated from standard enthalpies of formation by subtracting the sum of reactant enthalpies from the sum of product enthalpies. A negative ΔH indicates an exothermic reaction that releases heat, while a positive ΔH indicates an endothermic reaction that absorbs heat.
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Formula
ΔH°rxn = Σ ΔHf°(products) − Σ ΔHf°(reactants)
ΔH°rxn is the standard enthalpy of reaction. Σ ΔHf°(products) is the sum of standard enthalpies of formation of all products, each multiplied by its stoichiometric coefficient. Σ ΔHf°(reactants) is the equivalent sum for all reactants. Standard enthalpies of formation are referenced to elements in their standard states, which have ΔHf° = 0 by definition. A negative result indicates heat is released (exothermic); a positive result indicates heat is absorbed (endothermic).
How to use the Reaction Enthalpy Calculator
- 1
Enter your total enthalpy of products (σδhf products)
Value should be in kJ/mol.
- 2
Enter your total enthalpy of reactants (σδhf reactants)
Value should be in kJ/mol.
- 3
Read your results instantly
Results update in real time as you type.
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Hess's Law: enthalpy is a state function
Hess's Law states that the total enthalpy change for a reaction is the same regardless of the pathway taken, because enthalpy is a state function — it depends only on the initial and final states. This makes it possible to calculate ΔH for reactions that are difficult or impossible to measure directly by combining enthalpy values for related reactions. For example, the combustion of carbon to CO₂ is hard to control to stop at CO, so ΔHf of CO is calculated indirectly. Hess's Law also means that if you reverse a reaction, you change the sign of ΔH, and if you scale a reaction by a factor, you scale ΔH by the same factor.
Exothermic vs. endothermic reactions
An exothermic reaction has ΔH < 0 and releases heat to the surroundings — combustion, neutralization of strong acids with bases, and most oxidation reactions are familiar examples. The energy stored in chemical bonds of the products is lower than in the reactants, and the difference is released as heat or light. An endothermic reaction has ΔH > 0 and absorbs heat from the surroundings, making them feel cold to the touch. Dissolving ammonium nitrate in water, photosynthesis, and cooking an egg are all endothermic. The sign of ΔH is one of the two factors (along with entropy change ΔS) that determine whether a reaction is thermodynamically spontaneous under the Gibbs free energy framework.
Tips & Insights
Multiply ΔHf by stoichiometric coefficients
Standard enthalpies of formation are per mole of substance. If your balanced equation has 2 mol CO₂ as a product, you must multiply ΔHf(CO₂) by 2 before entering it as the products enthalpy.
Elements in standard state have ΔHf = 0
The standard enthalpy of formation of elemental substances in their standard state (O₂(g), C(graphite), H₂(g), etc.) is zero by definition. You do not need to include them in your calculation.
Negative ΔH = exothermic, positive = endothermic
A negative reaction enthalpy means the reaction releases heat (exothermic). A positive ΔH means it absorbs heat (endothermic). This sign convention follows the system's perspective: losing heat means negative enthalpy.
Worked Examples
Combustion of methane (CH₄ + 2O₂ → CO₂ + 2H₂O)
ΔH = −969.2 kJ/mol — strongly exothermic, consistent with natural gas releasing significant heat when burned.
Endothermic decomposition
ΔH = +90 kJ/mol — a positive ΔH confirms the reaction absorbs heat from the surroundings.
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Frequently Asked Questions
What is the standard enthalpy of formation?
It is the enthalpy change when one mole of a compound is formed from its elements in their standard states at 25°C and 1 bar pressure. Values for thousands of compounds are tabulated in reference databases.
What does a negative ΔH mean?
A negative ΔH means the reaction releases heat to the surroundings — it is exothermic. The products have lower enthalpy (stored energy) than the reactants, and the difference is released.
How is ΔH different from ΔG?
ΔH is the enthalpy change (heat at constant pressure). ΔG is the Gibbs free energy change, which also accounts for entropy: ΔG = ΔH − TΔS. A reaction is spontaneous when ΔG < 0, not just when ΔH < 0.
Can I use this for bond enthalpy calculations?
Yes, with an adjustment: ΔH ≈ Σ(bond energies broken) − Σ(bond energies formed). Enter the total energy of bonds broken as the reactants enthalpy and bonds formed as the products enthalpy. Bond enthalpies give approximate results; formation enthalpies give more accurate values.
What is Hess's Law?
Hess's Law states that the total enthalpy change for a chemical reaction is independent of the pathway taken. This allows ΔH to be calculated from tabulated formation enthalpies without measuring the reaction directly.
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