Endothermic and Exothermic Reactions
Ximena pulls an instant cold pack out of her backpack after getting hit during soccer practice. She squeezes it, the little bag gets cold in seconds, and it soothes the bruise. On the stove at her house, by contrast, LP gas burns and heats the pot. Both are chemical reactions or processes, but one cools and the other heats. Where does that energy come from, or where does it go? We answer at the end of the page.
System and Surroundings
To talk about energy in a reaction, you have to decide what you are studying. The system is the part we observe (the reactants and products). The surroundings are everything else: the air, the container, your hand.
Energy is conserved: if the system loses energy, the surroundings gain it, and vice versa. That is why what you feel (heat or cold) is what happens to the surroundings, and it is the opposite of what happens to the system.
Classify each process as endothermic or exothermic and justify your answer according to whether the system absorbs or releases energy: a) the combustion of LP gas on a…
Classify each process as endothermic or exothermic and justify your answer according to whether the system absorbs or releases energy: a) the combustion of LP gas on a stove; b) an instant cold pack (the kind used for bruises), which turns cold when activated; c) the oxidation of iron (the formation of rust); d) the evaporation of sweat from the skin, which cools you down.
The Sign of ΔH
Chemists measure the energy exchanged as heat at constant pressure and call it the enthalpy change, $\Delta H$:
$$\Delta H = H_{\text{products}} - H_{\text{reactants}}$$It is measured in kJ (or kJ/mol). The sign is taken from the system's point of view:
| Type | Energy of the system | Sign of $\Delta H$ | Example |
|---|---|---|---|
| Exothermic | The products have less energy than the reactants | $\Delta H < 0$ | Combustion of LP gas |
| Endothermic | The products have more energy than the reactants | $\Delta H > 0$ | Evaporation of sweat |
In a thermochemical equation, the heat can be written next to the equation. For example, the combustion of methane:
$$\text{CH}_4 + 2\,\text{O}_2 \rightarrow \text{CO}_2 + 2\,\text{H}_2\text{O} \qquad \Delta H = -890\ \text{kJ}$$It reads: for every mole of methane that burns, 890 kJ are released. If 2 mol burn, twice as much is released.
Classify each process as endothermic or exothermic and justify your answer according to whether the system absorbs or releases energy: a) the combustion of LP gas on a…
Classify each process as endothermic or exothermic and justify your answer according to whether the system absorbs or releases energy: a) the combustion of LP gas on a stove; b) an instant cold pack (the kind used for bruises), which turns cold when activated; c) the oxidation of iron (the formation of rust); d) the evaporation of sweat from the skin, which cools you down.
Why Is Energy Released or Absorbed?
A reaction regroups atoms: bonds in the reactants are broken and new bonds are formed in the products.
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Breaking a bond costs energy: it has to be supplied (it is always positive).
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Forming a bond releases energy: a stable molecule has less energy than the separated atoms.
The bond energy is the energy needed to break 1 mol of that bond (in kJ/mol); it is also the energy released when the bond forms.
So the balance depends on which side wins:
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If more energy is released in forming bonds than is spent in breaking the reactants' bonds, the reaction is exothermic.
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If more is spent in breaking than is released in forming, it is endothermic.
Calculating ΔH with Bond Energies
To estimate $\Delta H$: (1) write the balanced equation, (2) count the bonds broken and the bonds formed, multiplying by the coefficients, (3) add up each group, (4) subtract.
The approximate bond energies are: $\text{H}-\text{H}$: $436\ \text{kJ}/\text{mol}$, $\text{Cl}-\text{Cl}$: $243\ \text{kJ}/\text{mol}$ and $\text{H}-\text{Cl}$:…
The approximate bond energies are: $\text{H}-\text{H}$: $436\ \text{kJ}/\text{mol}$, $\text{Cl}-\text{Cl}$: $243\ \text{kJ}/\text{mol}$ and $\text{H}-\text{Cl}$: $432\ \text{kJ}/\text{mol}$. For the reaction $\text{H}_2 + \text{Cl}_2 \rightarrow 2\ \text{HCl}$: a) calculate the energy needed to break the bonds in the reactants; b) calculate the energy released when the bonds of the product are formed; c) find the $\Delta H$ of the reaction and state whether it is endothermic or exothermic.
Everyday Examples
| Process | Type | What you feel |
|---|---|---|
| Combustion of LP gas on the stove | Exothermic | Heat and flame |
| Rusting of iron | Exothermic (very slow) | Almost nothing, it is slow |
| Instant cold pack | Endothermic | It gets cold |
| Evaporation of sweat | Endothermic | Your skin cools off |
| Melting of ice | Endothermic | The soda gets cold |
| Photosynthesis | Endothermic | It needs sunlight |
When ice melts there is no reaction, but there is an energy change: physical processes are classified the same way. And a process being exothermic does not mean it is fast: iron releases energy when it rusts, but it is so slow that you do not notice it.
Back to the beginning: Ximena's cold pack contains a salt that, when it dissolves in water, absorbs energy from the surroundings (it is endothermic), and that is why it feels cold. On the stove, when LP gas burns, bonds that are stronger (in CO₂ and H₂O) than the ones that were broken are formed, and the leftover energy comes out as heat and light: it is exothermic.
Classify each process as endothermic or exothermic and justify your answer according to whether the system absorbs or releases energy: a) the combustion of LP gas on a…
Classify each process as endothermic or exothermic and justify your answer according to whether the system absorbs or releases energy: a) the combustion of LP gas on a stove; b) an instant cold pack (the kind used for bruises), which turns cold when activated; c) the oxidation of iron (the formation of rust); d) the evaporation of sweat from the skin, which cools you down.
The approximate bond energies are: $\text{H}-\text{H}$: $436\ \text{kJ}/\text{mol}$, $\text{Cl}-\text{Cl}$: $243\ \text{kJ}/\text{mol}$ and $\text{H}-\text{Cl}$:…
The approximate bond energies are: $\text{H}-\text{H}$: $436\ \text{kJ}/\text{mol}$, $\text{Cl}-\text{Cl}$: $243\ \text{kJ}/\text{mol}$ and $\text{H}-\text{Cl}$: $432\ \text{kJ}/\text{mol}$. For the reaction $\text{H}_2 + \text{Cl}_2 \rightarrow 2\ \text{HCl}$: a) calculate the energy needed to break the bonds in the reactants; b) calculate the energy released when the bonds of the product are formed; c) find the $\Delta H$ of the reaction and state whether it is endothermic or exothermic.