4,184
J
4.184
kJ
4.184
kJ/mol
1
4.184
kJ
4,184
J
4.184
kJ
4.184
kJ/mol
1
4.184
kJ
The Enthalpy Calculator determines the heat change (ΔH) for chemical reactions and physical processes using two methods: calorimetry (q = mcΔT) and Hess's law (ΔH°rxn = ΣΔH°f(products) − ΣΔH°f(reactants)). Enthalpy is a key thermodynamic quantity that measures the total heat content of a system at constant pressure. A negative ΔH indicates an exothermic reaction (releases heat), while a positive ΔH indicates an endothermic reaction (absorbs heat). This calculator is essential for energy balance calculations, process design, food science, and understanding the energetics of chemical transformations.
Method 1 — Calorimetry:
$$q = mc\Delta T$$
where m = mass (g), c = specific heat capacity (J/(g·°C)), ΔT = temperature change (°C). The enthalpy change per mole is:
$$\Delta H = \frac{q}{n}$$
Method 2 — Hess's Law:
$$\Delta H^\circ_{rxn} = \sum \Delta H^\circ_f(\text{products}) - \sum \Delta H^\circ_f(\text{reactants})$$
This uses tabulated standard enthalpies of formation. By definition, ΔH°f of elements in their standard states is zero. The sign convention: negative ΔH = exothermic (heat released), positive ΔH = endothermic (heat absorbed).
A negative ΔH (exothermic) means the reaction releases heat to the surroundings — the products are more stable than reactants. A positive ΔH (endothermic) means the reaction absorbs heat — energy input is required. The magnitude tells you how much energy is involved per mole. Large negative values (like combustion reactions, −800 kJ/mol for methane) indicate highly energetic processes.
Inputs
Results
q = 200 × 4.184 × 5.5 = 4602.4 J = 4.60 kJ. Per mole: 4.60/0.1 = 46.02 kJ/mol. Temperature increased, so the dissolution is exothermic (from the system's perspective, q is positive for surroundings).
Inputs
Results
CH₄ + 2O₂ → CO₂ + 2H₂O. ΔH° = [−393.5 + 2(−285.8)] − [−74.8 + 0] = −1182.6 − (−74.8) = −890.3 kJ/mol (approximate with rounded values).
Enthalpy (H) is a thermodynamic quantity equal to the internal energy plus the product of pressure and volume: H = U + PV. The enthalpy change (ΔH) at constant pressure equals the heat exchanged with surroundings.
Exothermic reactions release heat (ΔH < 0) and the surroundings warm up. Endothermic reactions absorb heat (ΔH > 0) and the surroundings cool down.
Hess's law states that the total enthalpy change for a reaction is independent of the pathway. You can calculate ΔH by summing formation enthalpies of products minus reactants, regardless of intermediate steps.
ΔH°f is the enthalpy change when one mole of a compound is formed from its elements in their standard states (25°C, 1 atm). By convention, ΔH°f of pure elements is zero.
Yes. ΔH varies with temperature according to Kirchhoff's equation: ΔH(T₂) = ΔH(T₁) + ∫ΔCp dT. For small temperature ranges, ΔH is approximately constant.
In the system convention, q > 0 means heat flows into the system (endothermic), and q < 0 means heat flows out (exothermic). Calorimetry measures the surroundings' temperature change.
For phase changes, use q = nΔHfus or q = nΔHvap instead of q = mcΔT. The calorimetry formula only applies when temperature changes, not during phase transitions.
Water has a specific heat of 4.184 J/(g·°C), which is unusually high compared to most substances. This makes water an excellent heat reservoir and calorimetric medium.
Hess's law is exact in principle since enthalpy is a state function. Accuracy depends on the quality of tabulated ΔH°f values, which are typically known to within ±0.1–1 kJ/mol for common compounds.
Enthalpy is measured in joules (J) or kilojoules (kJ). Molar enthalpy changes are in kJ/mol. Older literature may use calories (1 cal = 4.184 J).
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