Enthalpy Calculator
Enter coefficients and standard enthalpies of formation (ΔHf°) for up to two reactants and two products to compute ΔH_rxn, and see whether the reaction is exothermic or endothermic.
Reviewed by the ToolNestr Editorial Team — July 2026
Leave a species' ΔHf° blank to exclude it (useful for single-reactant or single-product reactions). Elements in their standard state have ΔHf° = 0.
Reactants
Products
Two ideas that trip students up
1. Exothermic: products end up lower
The reactant platform sits higher than the product platform. The drop in enthalpy is released as heat — ΔH is negative.
2. Endothermic: products end up higher
Here the product platform sits above the reactant platform — the reaction must absorb heat from its surroundings to climb that rise. ΔH is positive.
Enthalpy graphs
How it works
The core idea in one line: reaction enthalpy is just bookkeeping — add up the formation enthalpies of everything you end with, subtract the formation enthalpies of everything you started with.
ΔHrxn = ΣΔHf°(products) − ΣΔHf°(reactants)
each ΔHf° weighted by its stoichiometric coefficient
ΔH < 0 → exothermic
ΔH > 0 → endothermic
Each ΔHf° is weighted by its stoichiometric coefficient before summing. A negative ΔH_rxn means the products sit lower in enthalpy than the reactants — the excess energy is released as heat (exothermic). A positive ΔH_rxn means the products sit higher — the reaction must absorb heat to proceed (endothermic).
Worked example 1 — combustion of methane
Given: CH₄(g) + 2 O₂(g) → CO₂(g) + 2 H₂O(l). ΔHf°: CH₄ = −74.8, O₂ = 0 (element), CO₂ = −393.5, H₂O(l) = −285.8 kJ/mol.
ΔH_rxn ≈ −890.3 kJ — strongly exothermic, matching the standard textbook value for methane combustion.
Worked example 2 — simple A → B reaction
Given: A → B. ΔHf°(A) = −20 kJ/mol, ΔHf°(B) = +15 kJ/mol, both with coefficient 1, no second reactant or product.
ΔH_rxn = +35 kJ > 0, so this reaction is endothermic — it absorbs heat as A converts to B.
Standard enthalpies of formation (ΔHf°) for common substances
Reference values at 25°C, 1 atm, used in the worked examples above. Elements in their standard state are exactly zero by convention.
| Substance | ΔHf° (kJ/mol) |
|---|---|
| CO₂(g) | −393.5 |
| H₂O(l) | −285.8 |
| H₂O(g) | −241.8 |
| CH₄(g) | −74.8 |
| NH₃(g) | −46.1 |
| O₂(g), N₂(g), H₂(g) | 0 (elements) |
Values are standard reference data (approximate, rounded to one decimal) — always confirm against a current data table for precision work.
Where reaction enthalpy actually matters
🔥 Fuel combustion energy content
The enthalpy of combustion tells engineers how much heat a fuel releases per mole or per kilogram — the basis for comparing natural gas, gasoline, and hydrogen as energy sources, and for sizing burners and engines.
🏭 Industrial process heat management
Chemical plants use ΔH_rxn to predict whether a reactor needs active cooling (strongly exothermic reactions) or heating (endothermic reactions) to stay at a safe, efficient operating temperature.
🍎 Food calorie content
Nutritional Calories are essentially the enthalpy released when food is fully oxidized (a related but distinct measurement, done via bomb calorimetry) — the same underlying idea of heat released per mole of reaction.
❄️ Refrigerant & HVAC thermodynamics
Refrigeration cycles rely on the enthalpy change of a refrigerant as it evaporates and condenses, absorbing heat from inside a fridge or room and releasing it outside — engineers select refrigerants partly by their enthalpy of vaporization.
Common misconceptions
"Exothermic reactions always happen spontaneously and quickly."
ΔH says nothing about reaction rate (kinetics) or spontaneity. A reaction can be strongly exothermic yet extremely slow without a catalyst (e.g. rusting), or need an initial energy input (activation energy) before it proceeds. True spontaneity requires ΔG (Gibbs free energy), which also accounts for entropy and temperature.
"The ΔHf° of any pure element is always some large number."
It is defined as exactly zero for an element in its standard state (e.g. O2(g), N2(g), graphite), by convention — this is the reference point every other ΔHf° is measured against, not a coincidence or a measured-to-be-small value.
"A positive ΔH means the reaction cannot happen."
Endothermic reactions (ΔH > 0) happen all the time — melting ice, photosynthesis, many dissolution processes. They simply require a net heat input from the surroundings rather than releasing heat.
"ΔH and ΔHf° are always the same value."
ΔHf° is the enthalpy of formation of one specific compound from its elements. ΔH_rxn is the overall enthalpy change of a full reaction, calculated by combining the ΔHf° values of every reactant and product weighted by their coefficients.
Formula sources & further reading
The formulas here are standard, traceable to:
- • OpenStax, Chemistry 2e — Chapter 5, Thermochemistry (free, peer-reviewed). openstax.org
- • Brown, LeMay & Bursten, Chemistry: The Central Science — Chapter 5, Thermochemistry.
- • Zumdahl & Zumdahl, Chemistry — standard enthalpies of formation and Hess's law.
ΔH_rxn = Σ(coefficient × ΔHf°)products − Σ(coefficient × ΔHf°)reactants. Results are rounded for display.
How to use this calculator
Enter reactants
Coefficient and ΔHf° for reactant 1; optionally reactant 2. Leave ΔHf° at 0 for elements in their standard state.
Enter products
Coefficient and ΔHf° for product 1; optionally product 2.
Read the result
ΔH_rxn solves live, with an exothermic/endothermic label and the 3D energy-level scene.
Related tools
Frequently asked questions
What is enthalpy?
Enthalpy (H) is a measure of the total heat content of a system at constant pressure. Chemists usually care about the change in enthalpy, ΔH, during a reaction — the heat absorbed or released as reactants turn into products. It is measured in kJ or kJ/mol.
What does a negative or positive ΔH mean?
ΔH < 0 (negative) means the reaction is exothermic — it releases heat to the surroundings, and the products end up lower in enthalpy than the reactants. ΔH > 0 (positive) means the reaction is endothermic — it absorbs heat from the surroundings, and the products end up higher in enthalpy than the reactants.
Why is the standard enthalpy of formation of an element zero?
By convention, the standard enthalpy of formation (ΔHf°) of any element in its most stable form at standard conditions (25°C, 1 atm) — such as O2(g), N2(g), H2(g), or graphite for carbon — is defined as exactly zero. This gives every other ΔHf° value a consistent, comparable reference point, the same way sea level is the zero point for elevation.
How does ΔH relate to bond energies?
Breaking bonds always requires energy input (endothermic), while forming bonds always releases energy (exothermic). The overall ΔH of a reaction is approximately the energy needed to break the reactant bonds minus the energy released forming the product bonds. This bond-energy approach and the ΔHf°-based approach here are two ways to estimate the same ΔH_rxn.
Does ΔH tell you how fast a reaction happens?
No. ΔH is a thermodynamic quantity — it tells you the net heat exchanged and the relative stability of products versus reactants, not how quickly the reaction proceeds. Reaction rate is governed by kinetics (activation energy, temperature, catalysts), which is an entirely separate question from the sign or size of ΔH.