A reaction with a positive enthalpy absorbs heat from its surroundings. This means the products end up holding more energy than the reactants started with, and the difference has to come from somewhere: the environment. You can often feel this directly, as the surroundings get noticeably colder when the reaction takes place. In thermochemistry, a positive ΔH (enthalpy change) is the defining signature of an endothermic reaction.
Why the Enthalpy Change Is Positive
Every chemical reaction involves breaking bonds in the starting materials and forming new bonds in the products. Breaking bonds always requires energy, while forming bonds always releases it. In a reaction with a positive enthalpy, the bonds broken in the reactants are stronger than the bonds formed in the products. The system absorbs more energy during bond-breaking than it releases during bond-forming, so the net flow of energy is inward, from the surroundings into the reaction mixture.
This is the opposite of what happens in combustion or other exothermic reactions, where strong product bonds release more energy than it cost to break the reactant bonds. In an endothermic reaction, the combined enthalpy of the products is greater than the combined enthalpy of the reactants, making ΔH positive.
Common Examples
Positive enthalpy changes show up in everyday processes and large-scale chemistry alike. Some of the most familiar examples are phase transitions. Melting ice absorbs 6.02 kJ per mole of water, and boiling water absorbs 40.7 kJ per mole. Both are endothermic: the system pulls heat from the surroundings to overcome the forces holding molecules together in the solid or liquid state.
Photosynthesis is a major biological example. Plants absorb sunlight to convert carbon dioxide and water into sugars and oxygen, with an enthalpy change of roughly +469 kJ per mole. Without that continuous input of solar energy, the reaction simply would not proceed.
In industry, the thermal decomposition of calcium carbonate (limestone) into calcium oxide and carbon dioxide requires +177.8 kJ per mole and only occurs at temperatures above about 723°C. Kilns must supply heat continuously to keep this reaction going, which is why cement and lime production are energy-intensive processes.
Can a Positive Enthalpy Reaction Happen on Its Own?
A positive ΔH does not automatically mean a reaction is impossible without constant heating. Whether a reaction proceeds spontaneously depends on two factors: enthalpy and entropy (a measure of how much disorder increases). The relationship between them is captured by the Gibbs free energy equation: ΔG = ΔH − TΔS. A reaction is spontaneous when ΔG is negative.
If both ΔH and the entropy change (ΔS) are positive, the reaction is nonspontaneous at low temperatures but becomes spontaneous at high temperatures. That’s because the TΔS term grows with temperature, eventually outweighing the positive ΔH and driving ΔG negative. The decomposition of mercury oxide into liquid mercury and oxygen gas follows this pattern: endothermic but spontaneous once you heat it enough.
If ΔH is positive and ΔS is negative (the products are both higher in energy and more ordered), ΔG is positive at every temperature. The reaction is never spontaneous under any conditions. Converting oxygen gas into ozone is one example: it requires both energy input and results in fewer, more ordered molecules.
Ice melting illustrates temperature dependence neatly. At −10°C, the entropy increase of the system is not large enough to compensate for the heat absorbed, and the total entropy of the universe decreases, so melting is nonspontaneous. At +10°C, the math flips: the entropy of the universe increases by about 0.9 J/K, and melting proceeds on its own.
Energy Diagrams and Reaction Speed
On an energy diagram, a reaction with positive enthalpy looks like an uphill path. The products sit at a higher energy level than the reactants, and the vertical gap between them represents ΔH. Between the two levels, there is an even higher peak called the activation energy barrier, which represents the minimum energy the reacting molecules need to get the process started.
Because the products are already higher in energy than the reactants, the activation energy for an endothermic reaction is always at least as large as the ΔH itself. This tends to make endothermic reactions slower than comparable exothermic ones, since molecules need more energy just to reach the transition state. The equilibrium constant for these reactions is typically less than 1, meaning reactants are favored over products at standard conditions.
How Temperature Shifts the Equilibrium
For reversible reactions with a positive enthalpy, temperature is a powerful lever. Le Chatelier’s principle treats heat as a reactant in an endothermic process. Raising the temperature is like adding more of that reactant: it pushes the equilibrium toward the products, increasing yield. Lowering the temperature does the opposite, favoring the reverse (exothermic) direction and reducing product formation.
This is why many industrial decomposition reactions are run at high temperatures. The combination of overcoming the activation energy barrier and shifting the equilibrium toward products makes heat the essential driving force for positive enthalpy chemistry.
Measuring Positive Enthalpy in the Lab
The standard way to measure enthalpy changes is calorimetry. A simple version uses two nested Styrofoam cups filled with a known volume of water. You carry out the reaction inside the cup and record the temperature change. For an endothermic reaction, the water temperature drops because the reaction is pulling heat from the surrounding liquid.
The heat absorbed by the reaction equals the heat lost by the water, calculated with the formula: q = mass × specific heat × ΔT. Since the calorimeter loses heat while the reaction gains it, the sign is flipped: q for the reaction is the negative of q for the water reservoir. Dividing that value by the number of moles of reactant gives you ΔH per mole. When you dissolve certain salts like ammonium nitrate in water, the temperature drop is dramatic enough to feel through the cup, which is exactly the principle behind instant cold packs.