C2.5 Exothermic and endothermic reactions
When chemical reactions occur, energy is transferred to (exothermic) or from (endothermic) the surroundings.
Importantly, it is the particular process of the chemical reaction, that either generates heat (exothermic) to, or absorbs heat (endothermic) from, the surroundings. A common misconception is that students believe that the “reactants” alone release heat or absorb heat-this is incorrect!.
An exothermic chemical reaction is a reaction that results in the generation of heat energy as well as the chemical products of the reaction.
In doing so, REACTANTS gives rise to both HEAT ENERGY and PRODUCTS in an exothermic chemical reaction.
A classic example of this is the combustion of methane gas by a Bunsen burner”
Methane + Oxygen → Carbon Dioxide + Water + Heat energy
CH4 + 2O2 → CO2 + 2H2O + Heat energy
The heat energy that is produced, as a consequence of the chemical reaction, is transferred to the surroundings, hence the air surrounding a Bunsen flame heats up.
Apart from combustion, other examples of exothermic reactions include: many oxidation and acid/base neutralisations. Everyday uses of exothermic reactions include self-heating cans (eg for coffee) and hand warmers.
Because chemical ‘stored’ energy is released as heat energy during the chemical reaction, it follows that:
energy stored by the chemical bonds in the reactants > energy stored by the chemical bonds of the products
Hence, the reactants are situated at a higher chemical potential energy level than the products.
The overall energy released from the chemical bonds of the reactants during their conversion to the products is shown as as a downward displacement arrow (ΔH=−ve) in an energy level diagram.
Knowing that every chemical reaction must first break the chemical bonds present in the reactants then form new bonds that are present in the products, it follows that in an exothermic reaction,
Energy released by forming new bonds> Energy required to break the bonds
exothermic heat energy is released by the chemical reaction process.
An endothermic chemical reaction is a reaction that occurs by absorbing heat energy from the surroundings, and by doing so, will cool down the surroundings of the chemical reaction. The heat energy that is absorbed by the chemical reaction process essentially ‘steals’ this from the energy of the particles that surrounds the reactant particles.
In doing so, HEAT ENERGY + REACTANTS gives rise to the PRODUCTS of an endothermic chemical reaction.
Classic examples of endothermic reactions include dissolving ammonium nitrate crystals in water and the thermal decomposition of calcium carbonate. Some sports injury packs are based upon endothermic reactions.
Heat energy + Ammonium nitrate crystals + Water → Ammonium nitrate solution
Heat energy + NH4NO3 (s) → NH4+(aq) + NO3−(aq)
Heat energy + Calcium carbonate → Calcium oxide + Carbon Dioxide
Heat energy + CaCO3 → CaO + CO2
Because heat energy is added to the chemical ‘stored’ energy during the chemical reaction, it follows that:
energy stored by the chemical bonds in the products > energy stored by the chemical bonds of the reactants
Hence, the products are situated at a higher chemical potential energy level than the reactants.
The overall energy absorbed by the chemical bonds of the reactants during their conversion to the products is shown as as an upward displacement arrow (ΔH=+ve) in an energy level diagram.
Knowing that every chemical reaction must first break the chemical bonds present in the reactants then form the new bonds that are present in the products, it follows that in an endothermic reaction,
Energy released by forming new bonds< Energy required to break the bonds
endothermic heat energy is absorbed by the chemical reaction process
Exothermic or endothermic experiment
Boardworks simulation (exothermic or endothermic reaction)
BBC Bitesize (exothermic or endothermic reaction)
Practice questions (four)
Assessment and practical opportunities
■investigating temperature changes of neutralisations and displacement reactions, eg zinc and copper sulfate
■ investigating temperature changes when dissolving ammonium nitrate, or reacting citric acid and sodium hydrogencarbonate
■ adding ammonium nitrate to barium hydroxide
■ demonstration of the addition of concentrated sulfuric acid to sugar
■ demonstration of the reaction between iodine and aluminium after activation by a drop of water
■ demonstration of the screaming jelly baby
■ demonstration of the thermite reaction, ie aluminium mixed with iron(III) oxide
■ investigation of hand warmers, self-warming cans, sports injury packs.
There are opportunities here for measurements using temperature sensors to investigate energy transfer.