Exothermic and Endothermic Changes
An exothermic change releases energy to the surroundings, usually as heat, so the surroundings get warmer, while an endothermic change absorbs energy from the surroundings, so they get colder. In brief: burning a fuel and freezing water are exothermic, and melting ice and dissolving ammonium nitrate are endothermic.
Energy changes accompany every physical and chemical change. They explain why a campfire warms you, why sweating cools you and why a cold pack feels cold. This article explains the two terms, shows energy level diagrams in words, compares examples in a table, works through a temperature calculation, and discusses uses and mistakes.
What is an exothermic change?
In an exothermic change, the system gives out energy. The word comes from Greek: exo means “out” and thermic relates to heat. The temperature of the surroundings rises. Combustion, respiration, neutralization of an acid with an alkali and the setting of concrete are exothermic. Physical changes that release energy include condensation and freezing.
What is an endothermic change?
In an endothermic change, the system takes in energy. Endo means “in”. The temperature of the surroundings falls. Photosynthesis, thermal decomposition of calcium carbonate and the dissolving of some salts such as ammonium nitrate are endothermic. Physical changes that absorb energy include melting, evaporation and sublimation.
Why do energy changes happen?
Breaking bonds requires energy, and forming bonds releases energy. If the energy released in forming the new bonds in the products is greater than the energy needed to break the bonds in the reactants, the overall change is exothermic. If it is less, the change is endothermic. The energy is conserved: it moves between the chemicals and the surroundings.
On an energy diagram, the products of an exothermic reaction are at a lower energy than the reactants. In an endothermic reaction the products are at a higher energy. Many reactions also need a small input of energy to get started, called the activation energy, even if they are exothermic overall. A match must be struck, for instance, before it starts to burn.
Examples of exothermic and endothermic changes
| Change | Type | Energy flow | Everyday effect |
|---|---|---|---|
| Burning wood | Chemical, exothermic | Out | Heat and light |
| Respiration in cells | Chemical, exothermic | Out | Body warmth |
| Water freezing | Physical, exothermic | Out | Slight warming of surroundings |
| Rain forming in clouds (condensation) | Physical, exothermic | Out | Warms the air and powers storms |
| Photosynthesis | Chemical, endothermic | In | Plants absorb light energy |
| Ice melting | Physical, endothermic | In | Cools a drink |
| Sweat evaporating | Physical, endothermic | In | Cools the skin |
| Cooking an egg | Chemical, endothermic | In | Needs a heat source |
How can you measure the energy change?
- Place a measured volume of a solution in an insulated cup, such as a polystyrene (foam) cup, with a lid.
- Record the starting temperature, stirring gently.
- Add the second reactant, stir and record the highest (or lowest) temperature reached.
- Calculate the temperature change and use Q = m × c × ΔT, where m is the mass of the solution, c is the specific heat capacity (about 4.18 J/(g·°C) for dilute aqueous solutions) and ΔT is the temperature change.
Worked example: heat released in a reaction
In an experiment, 50.0 mL of 1.0 mol/L hydrochloric acid at 21.0 °C is mixed with 50.0 mL of 1.0 mol/L sodium hydroxide solution at 21.0 °C. The highest temperature reached is 27.8 °C.
- Mass of solution. The total volume is 100.0 mL. Assuming a density of 1.00 g/mL, the mass is 100.0 g.
- Temperature change. ΔT = 27.8 − 21.0 = 6.8 °C. The temperature rose, so the reaction is exothermic.
- Heat released. Q = 100.0 × 4.18 × 6.8 = 2,842 J, or about 2.84 kJ.
- Moles reacting. Each solution contains 0.0500 L × 1.0 mol/L = 0.050 mol, so 0.050 mol of water forms.
- Energy per mole. 2.84 kJ ÷ 0.050 mol = 56.8 kJ/mol, which is close to the accepted value of about 57 kJ per mole for the neutralization of a strong acid with a strong base.
The temperature rise showed that energy was released, and the calculation gave the amount. In an endothermic process the temperature would fall, and ΔT would be negative.
Where are these changes used?
- Heating: combustion of fuels and the self-heating packs that use an exothermic reaction.
- Cooling: instant cold packs for injuries use an endothermic dissolving process, and refrigerators use the evaporation of a fluid.
- Food: the cooking of food is endothermic, and the digestion of food releases energy in the body.
- Construction: concrete sets in an exothermic reaction, and large pours must be cooled to prevent cracking.
How did scientists study heat in reactions?
In the 1780s Antoine Lavoisier and Pierre-Simon Laplace built an ice calorimeter to measure the heat released by respiration and combustion. In the 19th century, James Joule showed that heat and mechanical work are equivalent forms of energy, and Germain Hess formulated in 1840 a law about the heat of reactions that is still used in chemistry. These ideas formed the basis of thermochemistry.
Which mistakes are common?
Mistake 1: Saying that endothermic means cold
Endothermic means that energy is absorbed. The surroundings get colder, but the reaction itself may need heating to proceed.
Mistake 2: Thinking that exothermic reactions need no energy to start
Most need some activation energy, such as the heat of a spark.
Mistake 3: Believing that energy is destroyed in endothermic reactions
The energy is stored in the chemical bonds of the products. It is not lost.
Mistake 4: Confusing the system with the surroundings
An exothermic reaction makes the solution warmer because energy is released to the surroundings. Always say which part you describe.