3.4.2Periodic Table and Energy

Calorimetry

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Section 1 of 5The Calorimetry Equation

Calorimetry converts a measured temperature change into the enthalpy change of a chemical reaction.

Heat energy

q = m c ΔT, where m is the mass of liquid heated in grams, c is the specific heat capacity and ΔT is the temperature change.

Specific heat capacity

For water or dilute aqueous solutions, c = 4.18 J g⁻¹ K⁻¹, which is the energy needed to heat 1 g of water by 1 K.

Enthalpy change

ΔH = −q ÷ n, with q in kJ and n in moles, where the negative sign shows that a temperature rise means an exothermic reaction.

Mass of solution

For aqueous reactions, assume that 1 cm³ of solution has a mass of 1 g, so mass = volume in cm³.

Measure a known volume of water or solution into a calorimeter such as a polystyrene cup or a metal can, and record its initial temperature.

For combustion, weigh a spirit burner containing fuel, light it, and heat the water for a time while stirring.

For reactions in solution, add the second reagent, stir constantly and record the highest or lowest temperature reached.

Calculate the temperature change, ΔT, and use q = mcΔT to find the heat transferred to the solution.

Divide by the moles of the limiting reactant, add the correct sign and express ΔH in kJ mol⁻¹.

Errors and Improvements

Heat loss

Heat is lost to the surroundings and to the apparatus,
- so the measured temperature rise is lower and ΔH is less exothermic than the data book value.

Incomplete combustion

Soot formation means less energy is released than for complete combustion, so the result is too low.

Insulation

Using a polystyrene cup with a lid, and a draught shield for combustion, reduces heat loss.

Extrapolating

For solutions, plot temperature against time and extrapolate back to the time of mixing to allow for heat loss.

Question

1.15 g of ethanol (M = 46.0 g mol⁻¹) is burned and heats 150 g of water by 25.0 K.
Calculate the enthalpy change of combustion of ethanol shown by the data.

q = m c ΔT = 150 × 4.18 × 25.0 = 15 675 J = 15.7 kJ.
What the symbols mean: E, Q or q = energy transferred (J) · m = mass (kg; g in chemistry calorimetry) · c = specific heat capacity (J kg⁻¹ °C⁻¹; water 4.18 J g⁻¹ °C⁻¹) · ΔT or Δθ = temperature change (°C or K)
n(ethanol) = 1.15 ÷ 46.0 = 0.0250 mol.
ΔH = −15.675 ÷ 0.0250 = −627 kJ mol⁻¹.

Answer ΔcH ≈ −627 kJ mol⁻¹, much less exothermic than the accepted value because of heat losses.

Specific heat capacity The energy required to raise the temperature of 1 g of a substance by 1 K.
Calorimeter Apparatus used to measure the heat transferred in a chemical reaction.
Enthalpy change of combustion The enthalpy change when 1 mol of a substance burns completely in oxygen under standard conditions.
Exothermic A reaction that releases heat and has a negative ΔH.

Using the mass of the fuel instead of the mass of the water in q = mcΔT.
Leaving q in joules when dividing by moles to give kJ mol⁻¹, so the answer is 1000 times too large.
Forgetting the negative sign for an exothermic reaction.
Using the volume of the total solution but then the moles of the excess reagent instead of the limiting one.