What Is an Endothermic Reaction? Energy Changes Explained

Chemical reactions involve breaking existing bonds and forming new ones. Breaking bonds requires energy, while forming bonds releases energy. In an endothermic reaction, the energy needed for bond breaking and other changes is greater overall than the energy released as new bonds form. The products therefore store more chemical energy than the reactants.

That energy difference is the key to an endothermic reaction. In terms of enthalpy, the overall enthalpy change is positive because the products are at a higher energy level than the reactants. An energy profile diagram shows the products above the reactants, with an activation energy barrier between them.

Because an endothermic process takes energy from its surroundings, the surroundings can lose thermal energy and become cooler. Some instant cold packs use dissolving processes that absorb energy in this way. The temperature change you measure depends on the substances, amounts, insulation, and experimental conditions.

When learning about endothermic reactions, be careful with the direction of energy transfer. The reacting system gains energy, while the surroundings lose it. Saying that a reaction 'contains cold' reverses the idea. Cold is not a substance being released. The lower temperature comes from energy leaving the surroundings and entering the process.

An exothermic reaction transfers energy from the reacting system to the surroundings overall. Combustion is a familiar example because it can release substantial heat. An endothermic reaction does the opposite overall, transferring energy into the system. Comparing the two is often easier than memorising each definition separately.

Both kinds of reaction can require activation energy to begin. An endothermic reaction does not simply absorb energy because it has activation energy, and an exothermic reaction can still need an initial spark. To explain what an endothermic reaction is correctly, separate the starting energy barrier from the overall difference between reactants and products. The overall energy change determines whether the reaction is endothermic.

Photosynthesis is the endothermic process that matters most: plants take in light energy to build glucose, which is why the sugar can later release energy when it is broken down. In the laboratory, dissolving ammonium nitrate in water makes the beaker noticeably cold, and that is the reaction inside many instant cold packs. Thermal decomposition is endothermic too, such as heating calcium carbonate to make quicklime in a kiln, where the heat has to be supplied continuously or the reaction stops.

Melting and evaporating are physical changes rather than reactions, but they absorb energy in the same way, which is why sweating cools you. Each example shows the same signature: energy goes in, the surroundings lose it, and the products sit at a higher energy level than what you started with.

It can seem strange that anything would soak up energy and happen on its own, since we usually picture reactions running downhill. The missing piece is entropy, a measure of how spread out energy and matter are. When ammonium nitrate dissolves, an orderly crystal breaks into ions scattered through the water, and that large increase in entropy outweighs the energy cost. Chemists combine the two effects in a quantity called Gibbs free energy. An endothermic change can proceed spontaneously when its entropy gain is large enough, and higher temperatures make that entropy term count for more.

You can measure the energy change with simple calorimetry. Mix known amounts of reactants in an insulated cup, record the starting and lowest temperatures, and use the relationship energy equals mass times specific heat capacity times temperature change. For dilute solutions, chemists usually treat the mixture as water, which needs about 4.18 joules to warm one gram by one degree Celsius. Dividing the energy by the number of moles that reacted gives an enthalpy change per mole. Heat exchanged with the cup and the air makes school results smaller than textbook values, a useful point when evaluating an experiment.

A helpful rule is that if a reaction is endothermic in one direction, the reverse reaction is exothermic by exactly the same amount of energy. Energy is conserved, so whatever the products gained compared with the reactants must be given back if the products turn back into the starting substances. On an energy profile diagram, this simply means reading the same picture from right to left: the products sit higher, so returning to the reactants is a downhill step overall.

Splitting water into hydrogen and oxygen using electricity is a clear example. The electrolysis is strongly endothermic, needing a continuous supply of electrical energy. When hydrogen later burns or reacts in a fuel cell to form water again, it releases that energy, which is why hydrogen is discussed as a way to store energy rather than as a source of it. In practice some energy is always lost as heat along the way, so you never recover all of the energy you put in during the endothermic step.

Many reactions are reversible and settle into an equilibrium, with the forward and backward reactions running at equal rates. Temperature shifts that balance in a predictable way: heating favours the endothermic direction, because that direction absorbs the extra energy, while cooling favours the exothermic one. A classic demonstration uses a sealed tube of nitrogen dioxide, a brown gas, in equilibrium with dinitrogen tetroxide, which is colourless. Splitting dinitrogen tetroxide into nitrogen dioxide is endothermic. Place the tube in hot water and the brown colour deepens, then move it into iced water and the colour fades.

The same principle explains a source of air pollution. Nitrogen and oxygen make up almost all of the air, yet they barely react at everyday temperatures, because forming nitrogen monoxide from them is strongly endothermic, taking in about 180 kilojoules for every two moles of product. In the intense heat of a lightning bolt, or inside a running car engine, the balance shifts and significant amounts of nitrogen monoxide form. When the gas cools quickly it does not have time to change back, and in the air it goes on to form nitrogen dioxide, which contributes to smog and acid rain.

A large share of the world's hydrogen is made from natural gas by an endothermic reaction called steam methane reforming. Methane and steam pass over a nickel catalyst inside tubes heated to roughly 700 to 1,000 degrees Celsius, forming carbon monoxide and hydrogen: CH4 plus H2O gives CO plus 3H2. The reaction absorbs about 206 kilojoules for every mole of methane, so heat must be supplied continuously, usually by burning extra natural gas in the furnace around the tubes. The high temperature does two jobs at once, speeding the reaction up and pushing the equilibrium towards the products.

The carbon monoxide is then usually reacted with more steam in a second stage, the water-gas shift, which forms carbon dioxide and extra hydrogen. That step is mildly exothermic, releasing about 41 kilojoules per mole, so it is run at lower temperatures, the opposite choice to the first stage. Much of the hydrogen goes on to make ammonia for fertilisers through the Haber process, which means an endothermic reaction sits near the start of the chain that feeds a large part of the world's population. It also explains why hydrogen made this way has a sizeable carbon footprint, from both the furnace and the shift.

An endothermic reaction is a reaction that absorbs energy from its surroundings overall, leaving products at a higher energy level than the reactants. Look for a positive energy change, possible cooling of the surroundings, and an energy diagram with products above reactants. Keep activation energy separate from the overall energy change.