Calorimetry Experiments (DP IB Chemistry): Revision Note
Calorimetry experiments
Measuring enthalpy changes
Calorimetry is a technique used to measure changes in enthalpy of chemical reactions
A calorimeter can be made up of a polystyrene drinking cup, a vacuum flask or metal can
Diagram to show how to set up a simple calorimeter
The energy needed to raise the temperature of 1 g of a substance by 1 K is called the specific heat capacity (c) of the liquid
The specific heat capacity of water is 4.18 J g-1 K-1
The energy transferred as heat can be calculated by:
q = m x c x ΔT
q = the heat transferred, J
m = the mass of water, g
c = the specific heat capacity, J g-1 K-1
ΔT = the temperature change, K
Worked Example
The energy from 0.01 mol of propan-1-ol was used to heat up 250 g of water. The temperature of the water rose from 298K to 310K (the specific heat capacity of water is 4.18 J g-1 K-1.
Calculate the enthalpy of combustion.
Answer:
Step 1: Determine the values
q = m x c x ΔT
m (of water) = 250 g
c (of water) = 4.18 J g-1 K-1
ΔT (of water) = 310 - 298 K = 12 K
Step 2: Substitute into the equation
q = 250 g x 4.18 J g-1 K-1 x 12 K
q = 12 540 J
Step 3: This is the energy released by 0.01 mol of propan-1-ol
Total energy ΔH = 12 540 J
Energy released per mole = q ÷ n = 12 540 J ÷ 0.01 mol = 1 254 000 J mol-1
Step 4: Convert J mol-1 to kJ mol-1 by dividing by 1000
Total energy = - 1254 kJ mol-1
Examiner Tips and Tricks
There's no need to convert the temperature units in calorimetry as the change in temperature in oC is equal to the change in temperature in K
Calorimetry experiments
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
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 cooling section of the graph is extrapolated back to the time when the reaction started to allow for heat loss
The steps to make a temperature correction graph are:
Take a temperature reading before adding the reactants for a few minutes to get a steady value
Add the second reactant and continue recording the temperature and time
Plot the graph and extrapolate the cooling part of the graph until you intersect the time at which the second reactant was added
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
Worked Example
Excess iron powder was added to 100.0 cm3 of 0.200 mol dm-3 copper(II) sulfate solution in a calorimeter. The reaction equation was as follows
Fe (s) + CuSO4 (aq) → FeSO4 (aq) + Cu (s)
The maximum temperature rise was 7.5 oC. Determine the enthalpy of reaction, in kJ.
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 (dm3) x concentration (mol dm-3)
moles = 0.1 x 0.2 = 0.02 mol
Step 3: Calculate ΔH and convert from J to kJ mol-1
ΔH = q ÷ n
ΔH = -3135 J ÷ 0.02 mol = -156 750 J mol-1
ΔH = -156.75 kJ mol-1
ΔH = -160 kJ mol-1 (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
Diagram to show the set up of a typical 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
Worked Example
1.023 g of propan-1-ol (M = 60.11 g mol-1) was burned in a spirit burner and used to heat 200 g of water in a copper calorimeter. The temperature of the water rose by 30 oC.
Calculate the enthalpy of combustion of propan-1-ol using this data.
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
moles = 1.023 g ÷ 60.11 g mol-1
moes = 0.01702 mol
Step 3: Calculate ΔH and convert to kJ mol-1
ΔH = q ÷ n
ΔH = -25 080 J ÷ 0.01702 mol
ΔH = -1 473 560 J mol-1
ΔH = -1 474 kJ mol-1
ΔH = -1.5 x 103 kJ mol-1 (2 sig figs)
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