Nobody can measure an enthalpy change directly, but anyone can read a thermometer. Let the reaction heat or cool a known mass of water, and q = mcΔT turns that temperature change into energy.
The heat transferred to or from the surroundings is q = mcΔT, where m is the mass of water or solution heated (not the mass of reactant), c = 4.18 J g−1 K−1 and ΔT is the temperature change. Divide by the amount in moles of the reactant that was used up (the limiting one) and convert J to kJ: ΔH = −q ÷ n. The minus sign means a temperature rise gives a negative (exothermic) ΔH. Heat loss to the surroundings makes experimental values less exothermic than data book values, most of all in combustion experiments.
Run a neutralisation in the calorimetry simulator, then a combustion with a spirit burner. Record the start and maximum temperatures and compare the value you calculate with the data book.
For a fuel, the water in a copper can is heated by the flame; weigh the burner before and after to find the mass of fuel burned. For a metal displacement, the metal is usually in excess, so the solution sets n.
Four quick checks. Each correct answer earns XP and records this skill on your profile.
In q = mcΔT for a fuel heating water, m is:
How much heat is needed to raise the temperature of 150 g of water by 8.0 K? (c = 4.18 J g−1 K−1)
The temperature of a solution rises during a reaction. The sign of ΔH is:
Why do simple combustion experiments give values less exothermic than data book values?
The biggest error in simple calorimetry is heat exchanged with the surroundings. Other assumptions matter too.
| Source of error | Effect | Improvement |
|---|---|---|
| heat lost to surroundings | ΔT too small, ΔH less exothermic | lid, insulation, draught shield; plot a cooling curve and extrapolate |
| incomplete combustion | less heat released | good oxygen supply |
| assuming c and density of solution equal those of water | small systematic error | use measured values |
| heat absorbed by the calorimeter | ΔT too small | include its heat capacity |
In q = mcΔT, m is the mass of the water or solution whose temperature changes, never the mass of fuel or solid added. Remember too that q comes out in joules: divide by 1000 for kJ, and add the minus sign yourself when the temperature rose.
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25.0 cm3 of 1.00 mol dm−3 HCl is mixed with 25.0 cm3 of 1.00 mol dm−3 NaOH; the temperature rises by 6.8 K. What is ΔH of neutralisation? (c = 4.18 J g−1 K−1)
Burning 0.460 g of ethanol (Mr 46.0) raises the temperature of 200 g of water by 12.0 K. What is the experimental ΔHc?
In a displacement reaction, excess zinc is added to 50.0 cm3 of 0.200 mol dm−3 CuSO4(aq). Which amount should be used as n in ΔH = −q/n?
Which change would make a calorimetry result for a reaction in solution more accurate?
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