Exothermic vs Endothermic Enthalpy Changes
An exothermic process releases heat to the surroundings and has a negative enthalpy change (ΔH < 0), while an endothermic process absorbs heat from the surroundings and has a positive enthalpy change (ΔH > 0). In brief: burning methane releases 890 kJ per mole (exothermic), while melting ice absorbs 6.01 kJ per mole (endothermic).
Whether a reaction feels hot or cold depends on the sign of its enthalpy change. The distinction matters in everyday applications such as hand warmers and instant cold packs, as well as in the design of chemical reactors. This article defines enthalpy change, compares exothermic and endothermic processes, explains energy diagrams, gives data for typical examples and works through calculations of heat from reaction amounts.
What is enthalpy change?
Enthalpy (H) is defined as H = U + PV. At constant pressure, the change in enthalpy is equal to the heat exchanged with the surroundings: ΔH = qp. For a chemical reaction, ΔH is the difference between the enthalpy of the products and that of the reactants:
ΔH = Hproducts − Hreactants
It is expressed in kJ for the reaction as written, or in kJ/mol. If the products have lower enthalpy than the reactants, the difference is released as heat. If the products have higher enthalpy, heat must be absorbed.
Exothermic and endothermic processes compared
| Feature | Exothermic | Endothermic |
|---|---|---|
| Heat flow | System to surroundings | Surroundings to system |
| Sign of ΔH | Negative | Positive |
| Surroundings | Get warmer | Get colder |
| Enthalpy of products | Lower than reactants | Higher than reactants |
| Bonds | Stronger bonds formed than broken | Weaker bonds formed than broken |
What do energy diagrams show?
In an enthalpy diagram, the vertical axis is enthalpy and the horizontal axis is the progress of the reaction. For an exothermic reaction the line of the products lies below that of the reactants, and the vertical difference is |ΔH|. For an endothermic reaction the products lie above the reactants. In both cases the curve rises to a peak between reactants and products, and the height of the peak above the reactants is the activation energy. A catalyst lowers the peak but does not change ΔH.
Enthalpy changes of common processes
| Process | Equation | ΔH (kJ/mol) | Type |
|---|---|---|---|
| Combustion of methane | CH4 + 2O2 → CO2 + 2H2O(l) | −890 | Exothermic |
| Neutralization | H+(aq) + OH−(aq) → H2O(l) | −56 | Exothermic |
| Freezing of water | H2O(l) → H2O(s) | −6.01 | Exothermic |
| Melting of ice | H2O(s) → H2O(l) | +6.01 | Endothermic |
| Vaporization of water | H2O(l) → H2O(g) | +40.7 | Endothermic |
| Dissolving ammonium nitrate | NH4NO3(s) → NH4+(aq) + NO3−(aq) | +25.7 | Endothermic |
| Decomposition of calcium carbonate | CaCO3 → CaO + CO2 | +178 | Endothermic |
How are thermochemical equations written?
A thermochemical equation is a balanced equation with states of matter and the value of ΔH, for example:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔH = −890 kJ
Three rules apply:
- ΔH is proportional to the amounts: doubling the equation doubles ΔH.
- The sign of ΔH changes when the reaction is reversed.
- The state of each substance must be stated, since ΔH depends on it: forming H2O(g) releases 44 kJ per mole less than forming H2O(l).
Worked examples
Example 1: heat from burning methane
How much heat is released by burning 10.0 g of methane? n = 10.0 / 16.04 = 0.6234 mol, and q = 0.6234 × (−890) = −555 kJ. So 555 kJ of heat is released.
Example 2: melting ice
How much heat is needed to melt 100 g of ice at 0 °C? n = 100 / 18.015 = 5.551 mol, and q = 5.551 × 6.01 = 33.4 kJ. In terms of specific values, this is 334 J per gram.
Example 3: evaporating water
How much heat is needed to evaporate 100 g of water at 100 °C? q = 5.551 × 40.7 = 226 kJ. This is about 2260 J/g, about seven times the energy needed to melt the same mass of ice. This is why steam burns are so severe: when steam condenses on the skin, it releases this energy.
Example 4: reversing a reaction
The reaction 2H2(g) + O2(g) → 2H2O(l) has ΔH = −571.6 kJ. Splitting 2 mol of liquid water into its elements, the reverse reaction, needs +571.6 kJ. For 1 mol of water, +285.8 kJ.
Where do you meet these changes in daily life?
- Hand warmers: the oxidation of iron powder in air is exothermic and generates heat for hours. Reusable gel warmers use the crystallization of sodium acetate.
- Instant cold packs: ammonium nitrate or urea dissolving in water absorbs heat, so the pack becomes cold.
- Sweating and refrigeration: evaporation absorbs heat and cools the skin or the fridge.
- Cooking and fuels: burning of fuels and the metabolism of food are exothermic.
- Photosynthesis: endothermic, driven by light energy.
Where does the word enthalpy come from?
The quantity H = U + PV was used earlier by Josiah Willard Gibbs as the “heat function at constant pressure”. The name enthalpy, from the Greek enthalpein (“to warm in”), is usually credited to the Dutch physicist Heike Kamerlingh Onnes in the early 20th century. Thermochemistry itself was developed from the 1840s by Germain Hess and, later, Marcellin Berthelot and Julius Thomsen, who measured the heats of many reactions.
Which mistakes are common?
- Confusing the sign: exothermic means negative ΔH, even though the system loses heat.
- Thinking that exothermic reactions are always spontaneous: the entropy change also matters.
- Forgetting to scale ΔH with the amount: ΔH is per equation as written.
- Ignoring states: liquid and gaseous water have different enthalpies.
- Assuming that a catalyst changes ΔH: it only changes the rate.