Chemistry / Unit 8: Stoichiometry
8.5 Energy in Reactions
Notes
Reactions either release or absorb energy as bonds break and form.
- Exothermic: releases heat to the surroundings, which feel warmer. $\Delta H$ is negative. Example: combustion.
- Endothermic: absorbs heat, so the surroundings feel colder. $\Delta H$ is positive. Example: a cold pack, photosynthesis.
- Breaking bonds absorbs energy; forming bonds releases energy.
- Thermochemical stoichiometry: $\Delta H$ is for the moles in the equation. $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$, $\Delta H = -890$ kJ, so 2.00 mol $\text{CH}_4$ releases 1780 kJ.
- Heat absorbed by water: $q = mc\Delta T$, with $c = 4.18$ J/(g·°C).
Practice
A reaction has $\Delta H = +45$ kJ. Which is true?
A positive $\Delta H$ means heat is absorbed from the surroundings.
$\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$, $\Delta H = -890$ kJ. How much heat is released when 32.1 g of $\text{CH}_4$ (16.05 g/mol) burns?
$\displaystyle \frac{32.1}{16.05} = 2.00$ mol, and $2.00 \times 890 = 1780$ kJ.
How much heat is needed to warm 50.0 g of water from 20.0 °C to 30.0 °C? ($c = 4.18$ J/(g·°C))
$q = mc\Delta T = 50.0 \times 4.18 \times 10.0 = 2090$ J.
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