as level

1. Calorimetry
Measurement of Enthalpy Change

2. Specific heat capacity is the amount of heat needed to
raise the temperature of 1g of substance by 1K
Specific heat capacity of water = 4.18 KJ kg-1
K-1
or 4.18 J g-1
K1
Be careful with the units it could also be quoted as KJ g-1
K-1
Ensure you use the correct units in your calculation!
To measure the heat released in a process we arrange for
the heat to be transferred to a substance (usually water)
then measure the temperature rise.

3. Then :
∆H = mass of water x specific heat capacity x temp rise
∆ H = m x c x ∆T
Note m = mass of water not mass of any solids present

4. Measuring Enthalpy Changes in the Laboratory
Apparatus needed
An insulated container to serve as a calorimeter
A thermometer
A balance
Volumetric appaaratus (e.g burette, pipette, measuring
cylinder)

5. A simple calorimeter

7. Some general steps in the procedure
1) Allow a known mass or volume of reactants to reach
the temperature of the surroundings
2) Thoroughly mix the reactants and record the highest or
lowest temperature reached
3) Determine the temperature change for the reaction
4) Calculate the enthalpy change for the reaction
For a given mass (m kg) of reacting substance the heat
energy released is calculated using the equation
Heat = m x c x ∆T

8. Assumptions and Errors
• For aqueous solutions we assume that 1ml has a mass
of 1g and that for dilute solutions the specific heat
capacity is the same as that of water.
• These assumptions will give minor errors in our
calculations
• The biggest error will be heat lost to the surroundings (i.e
to the thermometer, the surrounding air and the
container) This can be minimised by the use of an
adequately insulated calorimeter

9. Excess powdered zinc was added to 100ml of 0.2 mol/L copper (II)
sulphate solution. A temperature rise of 10o
C was recorded. Find the
enthalpy change for the reaction.
∆H = m x c x ∆T
∴ ∆H = 100g x 4.18 KJ kg-1
K-1
x 10 KJ
1000
= 4.180 KJ
This is for the no of moles of CuSO4 used in the experiment
No of moles of CuSO4 = 0.2 x 100 = 0.02 moles
1000
∴ ∆H = 4.180 = 209 KJ mol-1
0.02
The reaction is exothermic so we need to put in a negative sign
∴ ∆Hr = - 209 KJ mol-1
Note We do not use the standard sign as standard conditions were not
used.

10. Combustion
To find the heat of combustion of a substance a known
mass of the substance is burned, the heat released
transferred to water and the enthalpy change found as
before

11. In an experiment to find the heat of combustion of ethanol
the following results were obtained
Initial mass of lamp + ethanol = 65.20g
Final mass of lamp = 64.28g
Final temperature of water = 47.1o
C
Initial temperature of water = 28.5o
C
Mass of the water = 300g
What are the products of complete combustion of ethanol?
What mass of ethanol was burnt? How many moles is this?
What quantity of heat was transferred to the water?
Find ∆Hc of ethanol
Identify any sources of error
Is ethanol a good fuel?

12. C2H5OH + 3O2 → 2CO2 + 3H2O
∆H = 300 x 4.18 KJ kg-1
K-1
x 18.6K = 23.3KJ
1000
Mass of ethanol used = 0.92g
0.92g = 0.92 = 0.02 mol
46
∴ ∆Hc = - 23.3 = -1160 KJ/mol
0.02

13. Errors
Heat lost to surroundings (air, can thermometer)
Errors in measuring temperature change (unavoidable error
in reading thermometer)
Errors in measuring masses (unavoidable error in reading
balance)
The enthalpy change of combustion is high ∴ethanol is a
good fuel