Thermochemistry Calculator
Heat from calorimetry, Hess's law reaction enthalpy, and Gibbs free energy with the equilibrium constant.
q = m·c·ΔT (J); ΔH = −q/n. Hess: ΔHrxn = Σ(coef·ΔHf)products − Σ(coef·ΔHf)reactants. Gibbs: ΔG = ΔH − T·ΔS/1000 (kJ/mol); K = exp(−ΔG·1000/(R·T)) with R = 8.314 J/(mol·K).
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What it does
Thermochemistry answers three questions that look different but share the same energy bookkeeping: how much heat a process exchanges, what the enthalpy change of a reaction is when it cannot be measured directly, and whether a reaction is spontaneous at a given temperature. The calculator handles each with the standard textbook relations, so it is useful for checking homework, planning a calorimetry experiment, or sanity-checking a reported value before you trust it.
When to use it
Use calorimetry mode to turn a measured temperature rise into joules and, when the amount of reactant is known, into a molar enthalpy. Use Hess's law mode to add up tabulated formation enthalpies for a reaction that is hard to run cleanly in a calorimeter, such as a combustion or a synthesis with side products. Use Gibbs mode to see how a reaction's spontaneity changes with temperature, or to estimate the equilibrium constant behind a ΔG° value from a database.
FAQ
- What does the sign of q mean in calorimetry?
- A positive q is heat absorbed by the system, a negative q is heat released. The tool uses q = m·c·ΔT with mass in grams, specific heat in J/(g·K) and ΔT in kelvin or degrees Celsius, which are numerically the same interval. If the heat came from a reaction, the reaction enthalpy is ΔH = −q/n: the reaction releases what the solution absorbed, so the sign flips.
- How is Hess's law applied here?
- Each Hess row is one species with its stoichiometric coefficient and standard enthalpy of formation. The tool sums coefficient × ΔHf over the products, subtracts the same sum over the reactants, and the difference is ΔHrxn. Formation enthalpies of elements in their standard states are zero, so an O₂ row contributes nothing.
- How do ΔG, ΔH, ΔS and temperature relate?
- ΔG = ΔH − T·ΔS, with ΔH in kJ/mol, ΔS in J/(mol·K) and T in kelvin, so the entropy term is divided by 1000 to match units. A negative ΔG means the forward reaction is spontaneous at that temperature, and because T multiplies ΔS, a reaction can switch direction as the temperature changes.
- Why is the equilibrium constant so large or so small?
- K = exp(−ΔG°/(R·T)) is exponential: at 298 K a ΔG of −20 kJ/mol gives K ≈ 3×10³, while +20 kJ/mol gives about 3×10⁻⁴. A few kJ/mol moves K by orders of magnitude, and values beyond roughly 10³⁰⁰ are shown as ∞ because the reaction is effectively quantitative.