Exothermic and endothermic reactions
On a cold morning at a GAA match, a physio snaps a cold pack and it turns icy within seconds. Down the road, a hiker cracks open a hand warmer and it heats up. Nothing was plugged in, so where does the cold or the heat come from? By the end of this page you will be able to explain both.
Energy in and energy out
Every chemical reaction involves energy. Bonds in the reactants are broken and new bonds form in the products. Overall, energy is either given out to the surroundings or taken in from them.
The "surroundings" are the water, the air, the beaker or your hand: whatever is around the reacting chemicals. That is why you feel a cold pack get cold: it is taking heat from your skin.
| Exothermic | Endothermic | |
|---|---|---|
| Energy | Given out | Taken in |
| Surroundings | Get warmer | Get cooler |
| Temperature change | Rise (+) | Fall (−) |
| Examples | Burning natural gas, respiration, neutralisation, hand warmers | Photosynthesis, cold packs, thermal decomposition of limestone |
To find the temperature change, subtract: $\Delta T = T_{\text{final}} - T_{\text{initial}}$. A positive answer means exothermic. A negative answer means endothermic.
Define an exothermic reaction and an endothermic reaction. State what happens to the temperature of the surroundings in each case.
Define an exothermic reaction and an endothermic reaction. State what happens to the temperature of the surroundings in each case.
Classify each of the following as exothermic or endothermic: a) burning natural gas in a boiler b) photosynthesis in a leaf c) a self-heating coffee can d) a cold pack…
Classify each of the following as exothermic or endothermic:
- a) burning natural gas in a boiler
- b) photosynthesis in a leaf
- c) a self-heating coffee can
- d) a cold pack used for sports injuries
When ammonium nitrate dissolves in water in a beaker, the temperature falls from $21\ °\text{C}$ to $14\ °\text{C}$. a) Calculate the temperature change. b) Is the…
When ammonium nitrate dissolves in water in a beaker, the temperature falls from $21\ °\text{C}$ to $14\ °\text{C}$.
- a) Calculate the temperature change.
- b) Is the process exothermic or endothermic? Explain.
Measuring temperature changes
In a practical, you can tell whether a change is exothermic or endothermic with a thermometer. Measure the starting temperature, carry out the reaction or dissolve the solid, then record the highest or lowest temperature reached.
Use a polystyrene cup with a lid. Polystyrene is a good insulator, so little heat escapes to the air or is absorbed by the cup. A metal can would conduct heat away, and your readings would be too small.
| Solid | Initial (°C) | Final (°C) | $\Delta T$ (°C) | Type |
|---|---|---|---|---|
| X | 20.0 | 28.0 | +8.0 | Exothermic |
| Y | 20.0 | 12.0 | −8.0 | Endothermic |
| Z | 20.0 | 20.0 | 0.0 | No measurable change |
A $\Delta T$ of zero means that dissolving Z gives out or takes in no detectable energy. It is neither clearly exothermic nor endothermic.
How much energy?
You can calculate the energy transferred to the water using
$$Q = mc\Delta T$$where $m$ is the mass of the liquid in g, $c$ is its specific heat capacity (for water, $4.2\ \text{J/g °C}$) and $\Delta T$ is the temperature change in °C. $Q$ is in joules.
A student measures the temperature change when different solids dissolve in $100\ \text{cm}^3$ of water. | Solid | Initial temp (°C) | Final temp (°C) | |---|---|---| |…
A student measures the temperature change when different solids dissolve in $100\ \text{cm}^3$ of water.
| Solid | Initial temp (°C) | Final temp (°C) |
|---|---|---|
| A | 20.0 | 26.5 |
| B | 20.0 | 15.5 |
| C | 20.0 | 20.0 |
- a) Calculate the temperature change for each solid.
- b) Identify which dissolving is exothermic and which is endothermic.
- c) Explain what the result for C shows.
A student mixes $4.0\ \text{g}$ of citric acid solution with sodium hydrogencarbonate in $50\ \text{g}$ of water in a polystyrene cup. The temperature falls from…
A student mixes $4.0\ \text{g}$ of citric acid solution with sodium hydrogencarbonate in $50\ \text{g}$ of water in a polystyrene cup. The temperature falls from $22.0\ °\text{C}$ to $16.0\ °\text{C}$. Assume the specific heat capacity of the mixture is $4.2\ \text{J/g°C}$ and the total mass is $54\ \text{g}$.
- a) Calculate the energy absorbed using $Q = mc\Delta T$.
- b) State whether the reaction is endothermic and justify.
- c) Explain why a polystyrene cup is used instead of a metal can.
- d) Suggest one reason why the calculated value may differ from the true value.
Energy profile diagrams
An energy profile diagram shows how the energy of the chemicals changes during a reaction. The vertical axis is energy. The horizontal axis is the progress of the reaction.
Even exothermic reactions need a push to get going. A match will not light on its own: you must strike it first. The minimum energy needed to start a reaction is the activation energy.
-
Activation energy = peak − reactants.
-
Energy change $\Delta H$ = products − reactants.
For an exothermic reaction, the products are lower than the reactants, so $\Delta H$ is negative. For an endothermic reaction, the products are higher, so $\Delta H$ is positive.
Why does an endothermic reaction need energy supplied continuously? Its products hold more energy than the reactants, so energy must keep flowing in. A candle, once lit, gives out enough energy to supply the activation energy for the next bit of wax, so it keeps burning by itself.
An energy profile diagram for a reaction shows reactants at $50\ \text{kJ}$, a peak at $130\ \text{kJ}$ and products at $20\ \text{kJ}$. a) Calculate the activation…
An energy profile diagram for a reaction shows reactants at $50\ \text{kJ}$, a peak at $130\ \text{kJ}$ and products at $20\ \text{kJ}$.
- a) Calculate the activation energy.
- b) Calculate the energy change of the reaction ($\Delta H$ = products − reactants) and state if it is exothermic or endothermic.
- c) Describe where the products appear on the diagram compared with the reactants.
Another reaction has reactants at $30\ \text{kJ}$, a peak at $110\ \text{kJ}$ and products at $70\ \text{kJ}$. a) Calculate the activation energy. b) Calculate…
Another reaction has reactants at $30\ \text{kJ}$, a peak at $110\ \text{kJ}$ and products at $70\ \text{kJ}$.
- a) Calculate the activation energy.
- b) Calculate $\Delta H$ and state whether the reaction is exothermic or endothermic.
- c) Explain why this reaction needs a continuous supply of energy to keep going, while the burning of a candle does not once started.
Catalysts and energy calculations
A catalyst speeds up a reaction without being used up. It does this by providing an alternative pathway with a lower activation energy. More particles have enough energy to react, so the reaction is faster.
A catalyst does not change the energy of the reactants or products. So $\Delta H$ stays exactly the same.
Energy and moles
For some reactions, $\Delta H$ is given per mole. You can scale it up:
$$\text{energy} = \text{number of moles} \times \Delta H$$Thermal decomposition of calcium carbonate in a lime kiln is endothermic: $\text{CaCO}3 \rightarrow \text{CaO} + \text{CO}2$, with $\Delta H = +178\ \text{kJ}$ per mole…
Thermal decomposition of calcium carbonate in a lime kiln is endothermic: $\text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2$, with $\Delta H = +178\ \text{kJ}$ per mole of $\text{CaCO}_3$.
- a) Describe the shape of the energy profile for this reaction, stating the relative levels of reactants and products.
- b) Explain why the kiln must be kept hot continuously.
- c) Calculate the energy needed to decompose $2.5$ moles of $\text{CaCO}_3$.
- d) Calculate the mass of $\text{CaCO}_3$ ($M_r = 100$) that requires $890\ \text{kJ}$.
On an energy profile diagram, the uncatalysed reaction has an activation energy of $90\ \text{kJ}$. With a catalyst, the activation energy is $55\ \text{kJ}$. The…
On an energy profile diagram, the uncatalysed reaction has an activation energy of $90\ \text{kJ}$. With a catalyst, the activation energy is $55\ \text{kJ}$. The reactants are at $20\ \text{kJ}$ and the products are at $-30\ \text{kJ}$.
- a) Calculate $\Delta H$ for the uncatalysed and the catalysed reaction.
- b) Explain why $\Delta H$ is the same in both cases.
- c) State the heights of the peaks with and without the catalyst.
- d) Explain why the catalysed reaction is faster.
Back to the start: a cold pack contains chemicals that dissolve in a way that is endothermic, so it takes energy from your skin and feels cold. A hand warmer reaction is exothermic, so it gives out energy and feels hot.