Energy, heat, and enthalpy
Every chemical reaction absorbs or releases energy. Thermochemistry measures that energy, explaining why fuels burn hot and cold packs get cold.
Energy and the first law
Energy is the capacity to do work or transfer heat. The first law of thermodynamics says energy is conserved: it changes form but is never created or destroyed. In chemistry we track energy flowing between a system (the reacting substances) and its surroundings.
Exothermic and endothermic
Exothermic reactions release heat to the surroundings, which get warmer: combustion, respiration, and the setting of concrete. Endothermic reactions absorb heat, so the surroundings cool: photosynthesis, melting ice, and dissolving ammonium nitrate in an instant cold pack.
Enthalpy
At constant pressure, the heat transferred equals the change in enthalpy, \(\Delta H\). Exothermic reactions have \(\Delta H < 0\); endothermic ones have \(\Delta H > 0\). For burning methane, \(\Delta H \approx -890\) kJ per mole of methane. Enthalpy changes scale with amount: burning 2 mol releases twice as much heat.
q_p = \Delta HCalorimetry
Heat is measured by calorimetry: \(q = mc\Delta T\), where \(c\) is specific heat. Water's specific heat is 4.18 J/(g \(^\circ\)C), unusually high, which is why oceans moderate climates. If a reaction warms 100 g of water by 5.0 \(^\circ\)C, it released \(100 \times 4.18 \times 5.0 = 2090\) J.
q = mc\Delta THess's law
Because enthalpy depends only on the start and end states, the \(\Delta H\) of a reaction equals the sum of \(\Delta H\) values for any series of steps that adds up to it. This lets chemists compute enthalpies of reactions that are hard to measure directly, by combining known reactions like puzzle pieces.
Standard enthalpies of formation
\(\Delta H_f^\circ\) is the enthalpy change to form one mole of a compound from its elements in their standard states; for elements themselves it is zero. Then \(\Delta H^\circ_{\text{rxn}} = \sum \Delta H_f^\circ(\text{products}) - \sum \Delta H_f^\circ(\text{reactants})\). Tables of formation enthalpies let you compute the heat of thousands of reactions.
\Delta H^\circ_{rxn} = \sum \Delta H_f^\circ(\text{prod}) - \sum \Delta H_f^\circ(\text{react})Bond enthalpies
Breaking bonds absorbs energy; forming bonds releases it. Estimate \(\Delta H\) as bonds broken minus bonds formed. A reaction is exothermic when the bonds formed are stronger than those broken, as in combustion, where strong C=O and O–H bonds form.
Worked example
Using formation enthalpies: \(\text{CH}_4\) \(-74.8\), \(\text{CO}_2\) \(-393.5\), \(\text{H}_2\text{O}(l)\) \(-285.8\) kJ/mol. For \(\text{CH}_4 + 2\text{O}_2 \to \text{CO}_2 + 2\text{H}_2\text{O}\): \(\Delta H = [-393.5 + 2(-285.8)] - [-74.8 + 0] = -890.3\) kJ. Methane combustion releases about 890 kJ per mole.
Key takeaways
Energy is conserved. Exothermic reactions have \(\Delta H < 0\) and warm the surroundings; endothermic reactions have \(\Delta H > 0\). Calorimetry uses \(q = mc\Delta T\). Hess's law and formation enthalpies let you calculate \(\Delta H\) for any reaction; bond enthalpies give estimates.
Flashcard
Energy is conserved
Check your understanding
An instant cold pack gets cold when activated. The dissolving process is: