Calorimetry and enthalpy - Chemistry AP Study Notes

Calorimetry is the experimental technique used to measure heat energy changes during chemical reactions or physical processes. Enthalpy (H) represents the total heat content of a system at constant pressure, though we can only measure enthalpy change (ΔH), the heat absorbed or released.

Key Definitions:

• Exothermic reaction: ΔH is negative; heat is released to surroundings (products have less energy than reactants)
• Endothermic reaction: ΔH is positive; heat is absorbed from surroundings (products have more energy than reactants)
• Standard enthalpy change (ΔH°): measured at 298K, 100 kPa pressure, with substances in standard states
• Specific heat capacity (c): energy required to raise 1g of substance by 1K; water = 4.18 J g⁻¹ K⁻¹

Essential Equations:

q = mcΔT Where: q = heat energy (J), m = mass (g), c = specific heat capacity (J g⁻¹ K⁻¹), ΔT = temperature change (K or °C)

ΔH = -q/n Where: ΔH = enthalpy change (kJ mol⁻¹), q = heat measured (convert to kJ), n = moles of limiting reactant

The negative sign converts heat gained by water (positive q) to heat lost by reaction (negative ΔH for exothermic).

Mnemonic - MCAT: Mass × Capacity × ΔT = Heat

Key Principle: In simple calorimetry, we assume all heat from the reaction transfers to the water/solution with negligible heat loss to surroundings (though real experiments have significant errors).

The Coffee Cup Analogy: Imagine your morning coffee in a polystyrene cup (a simple calorimeter). When you add hot coffee to cold milk, the mixture's final temperature tells you how much heat transferred. Similarly, when magnesium burns in acid inside a polystyrene cup, the acid's temperature rise reveals the reaction's enthalpy change.

Real-World Applications:

Food Energy Labels: Food calories are determined using bomb calorimetry. Food burns completely in pure oxygen inside a sealed steel container surrounded by water. The temperature rise of the water reveals the energy content. That "250 calories" on your chocolate bar was literally measured by burning chocolate and watching water heat up!

Hand Warmers: These contain supersaturated sodium acetate. When triggered, crystallization occurs (exothermic ΔH = -19.7 kJ mol⁻¹), releasing heat. The reverse reaction—dissolving the crystals—is endothermic, requiring heat input. This demonstrates enthalpy as a state function: the forward and reverse ΔH values are equal but opposite in sign.

Industrial Process Design: Chemical engineers use calorimetry data to design reactors. Exothermic reactions like ammonia synthesis (Haber process, ΔH = -92 kJ mol⁻¹) require cooling systems to remove heat and prevent dangerous temperature rises. Endothermic reactions like thermal decomposition need continuous heating.

Why Polystyrene Cups? They're excellent insulators (low thermal conductivity), minimizing heat loss to surroundings—a critical assumption in calculations. Professional calorimeters have stirrers, thermometers, and insulated lids, but the principle remains: measure temperature change, calculate heat transfer.

Remember: The surroundings (water/solution) experience the opposite temperature change to the system (reaction). Exothermic reactions warm the water; endothermic reactions cool it.