Calorimetry and calculating enthalpy change

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Measuring the **temperature change** of a known mass of water (or solution) to find the heat transferred by a reaction.

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The energy needed to raise **1 g** of a substance by **1 K** (1 °C). For water, **c = 4.18 J g⁻¹ K⁻¹**.

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$Q = mc\Delta T$ — heat (J) = mass (g) × specific heat capacity × temperature change.

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ΔT = **T_{final} − T_{initial}**. A change of 1 °C equals a change of 1 K, so the number is the same.

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$\Delta H = -\dfrac{Q}{n}$ — divide Q (in kJ) by the amount that reacted, and add the sign.

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**Exothermic** — heat released to the water — so **ΔH is negative**.

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**Endothermic** — heat absorbed from the water — so **ΔH is positive**.

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The mass of **water** (the substance heated), **not** the mass of fuel or reactant.

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Q from $mc\Delta T$ is in **joules**; enthalpy changes are quoted in **kJ mol⁻¹**, so divide by 1000.

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**Heat loss** to the surroundings and apparatus — so the measured ΔH is **less exothermic** than the true value.

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All the heat goes to the **water**, and the **specific heat capacity** (and density) of the solution equals that of water.

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ΔT → **Q = mcΔT** → ÷1000 for kJ → **÷ n** for per mole → add the **sign**.