-965.1
kJ/mol
-74.8
kJ/mol
-890.3
kJ/mol
-1
890.3
kJ/mol
-890.3
kJ/mol
-965.1
kJ/mol
-74.8
kJ/mol
-890.3
kJ/mol
-1
890.3
kJ/mol
-890.3
kJ/mol
The Heat of Formation Calculator applies Hess's law to determine the standard enthalpy of reaction (ΔH°rxn) from tabulated standard enthalpies of formation (ΔH°f). The standard enthalpy of formation is the enthalpy change when one mole of a compound is formed from its elements in their standard states at 25°C and 1 atm. This calculator accepts up to three products and three reactants with their stoichiometric coefficients and ΔH°f values. It is widely used in physical chemistry, chemical engineering, and environmental science to predict whether reactions release or absorb energy without performing experiments.
Hess's law states that ΔH for a reaction is independent of the pathway:
$$\Delta H^\circ_{rxn} = \sum n_i \Delta H^\circ_f(\text{products}) - \sum n_j \Delta H^\circ_f(\text{reactants})$$
where ni and nj are stoichiometric coefficients. Key conventions:
The result tells you whether the reaction is exothermic (ΔH°rxn < 0) or endothermic (ΔH°rxn > 0).
A negative ΔH°rxn indicates an exothermic reaction — products are more stable than reactants. A positive ΔH°rxn means the reaction is endothermic. The magnitude shows how much energy is released or absorbed per mole of reaction as written. This information is critical for assessing reaction feasibility and designing safe industrial processes.
Inputs
Results
Products: 1(−393.5) + 2(−285.8) = −965.1 kJ. Reactants: 1(−74.8) + 2(0) = −74.8 kJ. ΔH° = −965.1 − (−74.8) = −890.3 kJ. Highly exothermic combustion.
Inputs
Results
CaCO₃ → CaO + CO₂. Products: (−635.1) + (−393.5) = −1028.6. Reactants: −1206.9. ΔH° = −1028.6 − (−1206.9) = +178.3 kJ. Endothermic — requires heating.
ΔH°f is the enthalpy change when exactly one mole of a compound is formed from its elements, each in their standard state (25°C, 1 atm). For example, ΔH°f(H₂O(l)) = −285.8 kJ/mol.
By convention, elements in their most stable form at 25°C and 1 atm are defined as the reference state with ΔH°f = 0. This includes O₂(g), N₂(g), C(graphite), Fe(s), etc.
Hess's law states that the total enthalpy change for a reaction is the same regardless of the number of steps. This allows calculation of ΔH from tabulated formation enthalpies without running the actual reaction.
Standard enthalpies of formation are tabulated in chemistry textbooks, the CRC Handbook, and the NIST Chemistry WebBook. Values are well-established for thousands of compounds.
Enter ΔH°f = 0 for any element in its standard state. This includes diatomic gases (O₂, N₂, H₂, F₂, Cl₂), graphite, and metals in their most stable crystal form.
Yes, significantly. ΔH°f(H₂O(g)) = −241.8 kJ/mol while ΔH°f(H₂O(l)) = −285.8 kJ/mol. The 44 kJ/mol difference is the enthalpy of vaporization. Always use the correct phase.
Yes, if you use ΔH°f values for aqueous species (e.g., ΔH°f(Na⁺(aq)) = −240.1 kJ/mol). These account for solvation enthalpies.
Very accurate — Hess's law is exact since enthalpy is a state function. The accuracy depends on the quality of tabulated ΔH°f values, which are typically known to ±0.1−1 kJ/mol.
Negative ΔH°rxn = exothermic (releases heat). Positive ΔH°rxn = endothermic (absorbs heat). This follows from the products being lower or higher in enthalpy than reactants.
Yes. Some compounds have positive ΔH°f, meaning they are less stable than their elements. Examples: NO(g) = +90.3 kJ/mol, C₂H₂(g) = +226.7 kJ/mol. These are endothermic to form.
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