Calorimetry Experiments (DP IB Chemistry: HL)

• There are two types of calorimetry experiments you need to know for IB Chemistry:
  • Enthalpy changes of reactions in solution
  • Enthalpy changes of combustion
• In both cases you should be able to give an outline of the experiment and be able to process experimental data using calculations or graphical methods

Enthalpy changes for reactions in solution

• The principle of these calorimetry experiments is to carry out the reaction with an excess of one reagent and measure the temperature change over the course of a few minutes
• The apparatus needed to carry out an enthalpy of reaction in solution calorimetry experiment is shown in the previous section 5.1.3 Calorimetry
• For the purposes of the calculations, some assumptions are made about the experiment:
  • That the specific heat capacity of the solution is the same as pure water, i.e. 4.18 J g^-1 K^-1
  • That the density of the solution is the same as pure water, i.e. 1 g cm^-3
  • The specific heat capacity of the container is ignored
  • The reaction is complete
  • There are negligible heat losses

Temperature correction graphs

• For reactions which are not instantaneous there may be a delay before the maximum temperature is reached
• During that delay the substances themselves may be losing heat to the surroundings, so that the true maximum temperature is never actually reached
• To overcome this problem we can use graphical analysis to determine the maximum enthalpy change

A temperature correction graph for a metal displacement reaction between zinc and copper sulfate solution. The zinc is added after 4 minutes

The steps to make a temperature correction graph are:

• An assumption made here is that the rate of cooling is constant
• The analysis can also be used for endothermic reactions, but this time there will be a ‘warming’ section as the substances return to room temperature

Answer:

Step 1: Calculate q

q = m x c x ΔT

q = 100 g x 4.18 J g^-1 K^-1 x 7.5 K = - 3135 J

Step 2: Calculate the amount of CuSO₄ (aq)

moles = volume in dm³ x concentration = 0.1 x 0.2 = 0.02 mol

Step 3: Calculate ΔH

ΔH = q ÷ n =  -3135 J ÷ 0.02 mol = - 156 750 J = -156.75 kJ = -160 kJ (2 sig figs)

Enthalpy of Combustion Experiments

• The principle here is to use the heat released by a combustion reaction to increase the heat content of water
• A typical simple calorimeter is used to measure the temperature changes to the water

A simple combustion calorimeter

• Not all the heat produced by the combustion reaction is transferred to the water
  • Some heat is lost to the surroundings
  • Some heat is absorbed by the calorimeter
• To minimise the heat losses the copper calorimeter should not be placed too far above the flame and a lid placed over the calorimeter
• Shielding can be used to reduce draughts
• In this experiment the main sources of error are
  • Heat losses
  • Incomplete combustion

Answer:

Step 1: Calculate q

q = m x c x ΔT

q = 200 g x 4.18 J g^-1 K^-1 x 30 K = - 25 080 J

Step 2: Calculate the amount of propan-1-ol burned

moles = mass ÷ molar mass = 1.023 g ÷  60.11 g mol^-1 = 0.01702 mol

Step 3: Calculate ΔH

ΔH = q ÷ n =  -25 080 J ÷ 0.01702 mol = - 1 473 560 J = -1 474 kJ = -1.5 x 10³ kJ