How do you know if a reaction is endothermic or exothermic

In thermochemistry the heat exchanged between reacting substances and their surroundings is quantified by the enthalpy change, ΔH. A negative ΔH indicates that the reaction releases heat to the surroundings, so the temperature of the surroundings rises – this is an exothermic process. A positive ΔH means heat is taken from the surroundings, causing a temperature drop – an endothermic process.

Sign of ΔH and Heat Flow

Calorimetry directly measures the temperature change of a known mass of water (or another solvent) that absorbs or releases heat during the reaction. Using $q = m c \Delta T$ you calculate the heat q transferred, then divide by the number of moles reacting to obtain ΔH per mole. The sign of q (and therefore ΔH) tells you whether the reaction is endothermic or exothermic.

Three quick ways to decide whether a reaction is endothermic or exothermic

• Measure the temperature change of the surroundings
• Look up ΔH in a thermodynamic table
• Calculate ΔH from bond‑energy differences

Procedure to obtain ΔH from a simple calorimetry experiment

1. 1Weigh a known mass of water and record its initial temperature.
2. 2Carry out the reaction in the water and record the final temperature.
3. 3Compute $q = m c \Delta T$ with $c = 4.18\ \text{J g}^{-1}\text{K}^{-1}$.
4. 4Divide q by the moles of reactant to get ΔH per mole and note the sign.

Typical ΔH values for common reactions

Example: 5.00 g of ammonium nitrate (NH₄NO₃) dissolves in 100 g of water. The temperature falls from 25.0 °C to 22.3 °C. Using $q = m c \Delta T$ with $m = 100\ \text{g}$ and $c = 4.18\ \text{J g}^{-1}\text{K}^{-1}$ gives $q = 100 \times 4.18 \times (22.3-25.0) = -1.13\times10^{3}\ \text{J}$. The negative q indicates heat was absorbed by the system, so ΔH = +1.13 kJ for the 5.00 g sample. Converting to per‑mole basis (molar mass 80.04 g mol⁻¹) yields ΔH ≈ +18 kJ mol⁻¹, confirming an endothermic process.