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Unit 6
7–9% of examThermochemistry tracks the energy that flows as heat when substances warm up, change phase or react. You'll learn to tell endothermic from exothermic processes, use calorimetry and q = mcΔT, and calculate enthalpy changes from bond enthalpies, enthalpies of formation and Hess's law.
Longer videos that cover the whole unit. Good for a first pass or a final review.
In an exothermic process the system releases energy to the surroundings, so they warm up; in an endothermic process the system absorbs energy, so the surroundings cool. Dissolving can go either way, depending on whether the attractions formed are stronger or weaker than the ones broken.
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An energy diagram compares the energy of reactants and products: products lower than reactants means exothermic, and products higher means endothermic. The same kind of diagram works for physical changes such as melting or dissolving.
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Particles in a warmer object have a higher average kinetic energy, and their collisions with particles of a cooler object pass energy along as heat. This keeps going until thermal equilibrium, when both have the same average kinetic energy and so the same temperature.
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q = mcΔT gives the heat gained or lost when a mass m of a substance with specific heat capacity c changes temperature by ΔT. In a calorimeter, energy released by a reaction or by dissolving is absorbed by the surrounding solution (and the reverse), so a temperature rise means the process was exothermic.
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Melting and boiling absorb energy, while freezing and condensing release the same amount, and a pure substance's temperature stays constant while it changes phase. You find the heat by multiplying the moles by the molar enthalpy of the change, such as the enthalpy of fusion or of vaporization.
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The enthalpy of reaction, ΔH, is the heat released (negative) or absorbed (positive) by a reaction at constant pressure, for the amounts in the balanced equation. Scale it to the moles that actually react: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) has ΔH° ≈ −890 kJ/mol, so burning 2.0 mol of CH₄ releases about 2.0 × 890 ≈ 1.8 × 10³ kJ.
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Breaking bonds takes energy and forming bonds releases it, so you can estimate ΔH as the energy of the bonds broken minus the energy of the bonds formed. If forming the product bonds releases more energy than breaking the reactant bonds takes, the reaction is exothermic. The answer is an estimate because bond enthalpies are averages.
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The standard enthalpy of formation (ΔH°f) is the enthalpy change for making one mole of a compound from its elements in their standard states; for an element in its standard state it is zero. A reaction's ΔH° equals the sum of the products' ΔH°f values minus the sum of the reactants', each multiplied by its coefficient.
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Because energy is conserved, the ΔH of an overall reaction equals the sum of the ΔH values of any steps that add up to it. If you reverse a step, flip the sign of its ΔH; if you multiply a step by a number, multiply its ΔH by that number too.
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