Calorimetry Measurement of Enthalpy Change

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CalorimetryMeasurement of Enthalpy Change
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• 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.

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• 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

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• 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)

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A simple calorimeter
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• Some general steps in the procedure
• 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
• Determine the temperature change for the reaction
• 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

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• 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

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• Excess powdered zinc was added to 100ml of 0.2
  mol/L copper (II) sulphate solution. A
  temperature rise of 10oC 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.

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• 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

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• 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 47.1oC
• Initial temperature of water
  28.5oC
• 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?

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• 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

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• 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