Calorimetry

1
Calorimetry

• The enthalpy change associated with a chemical
  reaction or process and the specific heat of a
  substance can be measured experimentally using
  calorimetry.
• The experimental measurement of the amount of
  heat gained or lost
• Determined by measuring the temperature change
  that occurs.
• Calorimeter
• An instrument used to measure the heat gained or
  lost

2
Calorimetry

• Two types of calorimetry are commonly used
• Constant-pressure Calorimetry
• Specific heat
• DHrxn or DHsoln
• Bomb Calorimetry (Constant Volume Calorimetry)
• DHcombustion

3
Calorimetry

• In constant-pressure calorimetry,
• Measurements are made in an open container.
• Reactions occur at constant pressure
• Calorimeter is insulated
• Assume that the amount of heat gained from or
  lost to the surroundings is negligible
• any heat gained or lost during a chemical
  reaction or process comes from or goes into the
  solution being studied.

4
Calorimetry

• In order to measure the specific heat of a
  substance, we need to know three values
• qsubstance
• mass
• DT
• In order to measure a molar enthalpy change for a
  reaction, we need to know two values
• qsubstance or qreactant
• moles of substance or reactant

5
Calorimetry

• In both cases, the unknown that must be
  determined experimentally is the amount of heat
  gained or lost by the substance being studied
  (i.e. qsubstance).
• In constant-pressure calorimetry, qsubstance is
  determined indirectly by measuring the heat
  gained or lost by the liquid present in the
  calorimetry (whose specific heat is known).

6
Calorimetry

• We then apply the First Law of Thermodynamics
• If heat is lost by the chemicals (or object)
  during a reaction or process, then it must be
  gained by the solution (liquid) and vice versa.
• The heat gained or lost by the reactants (or
  object) and the solution (liquid) are equal in
  magnitude but opposite in sign.
• qobject - qsoln or qrxn -qsoln

7
Calorimetry

• Calorimetry Simulation
• Measuring the specific heat of a metal
• How is the experiment performed?
• What kind of data is collected?
• How are the data used to determine specific heat?

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rojectfolder/animationsindex.htm
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Calorimetry

• Determining specific heat or molar heat capacity
  experimentally
• Heat an object with a known mass to 95-100oC.
• Measure initial temperature of object.
• Place known mass of water into calorimeter and
  measure its temperature.
• Add hot object to the water
• Measure the equilibrium temperature (i.e. final
  temperature) of the water and object.

9
Calorimetry

• Types of data obtained from or used in a
  calorimetry experiment often include
• Sample Data
• Mass of sample (object or compound)
• Tinitial (sample)
• Tfinal (sample)
• Solution Data
• Mass of solution
• Tinitial (solution)
• Tfinal (solution)
• Csoln
• Often use specific heat of water if water is the
  solvent

10
Calorimetry

• Using the data to find specific heat
• Calculate the amount of heat gained by the water
  
• qwater Cwater x mwater x DTwater
• Calculate the amount of heat lost by the object
• qobj - qwater
• Calculate specific heat of object
• Cs qobj
• mobj x DTobj

11
Calorimetry

• Example A 55.0 g piece of aluminum metal at
  100.0oC was added to 51.3 g of water at 20.0oC.
  The equilibrium temperature of the system was
  35.0oC. If the specific heat of water is 4.18
  J/g.K, what is the specific heat of aluminum?

12
Calorimetry
13
Calorimetry

• This is the same procedure you will use to
  calculate the specific heat of the metal sample
  you use in the lab.
• Follow this procedure to do the calculations for
  Expt. 2.
• You should expect a similar problem on your exam!

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Calorimetry

• Calorimetry Simulation
• Measuring the molar heat of solution (DHsoln)
• How is the experiment performed?
• What kind of data is collected?
• How are the data used to determine molar heat of
  solution?

15
Calorimetry

• Determining molar heat of solution (DHsoln per
  mole of solute) from calorimetry data.
• Place known mass of water into calorimeter.
• Measure the initial temperature of the water.
• Add known mass of solute to water.
• Record equilibrium temperature of the resulting
  solution.

16
Calorimetry

• Types of data collected when measuring molar heat
  of solution
• Mass of water
• Initial temperature of water
• Mass of solute
• Final temperature of solution

17
Calorimetry

• Determining DHsoln per mole of solute from
  calorimetry data.
• Assume Csoln Cwater 4.18 J/g-K unless told
  otherwise
• Calculate the heat gained or lost by the entire
  solution
• qsoln Csoln x mass of solution x DTsoln
• Mass of solution mass of H2O mass of solute

18
Calorimetry

• Calculate heat gained or lost by the solute
  (salt)
• qsolute - qsoln
• Calculate molar heat of solution
• DHsoln per mole qsolute
• mole solute

19
Calorimetry

• Example When a 2.125 g sample of solid ammonium
  nitrate dissolves in 30.000 g of water in a
  constant pressure calorimeter, the temperature
  drops from 22.0oC to 16.9oC. Calculate the molar
  DHsoln (in kJ/mole) for the dissolution process
• NH4NO3 (s) ? NH4 (aq) NO3- (aq)
• Assume that the specific heat of the solution is
  the same as pure water (4.18 J/g.K).

20
Calorimetry
21
Calorimetry

• You should expect a calorimetry problem similar
  to this on your exam.
• What would you actually observe when the
  dissolution process is exothermic?
• What would you actually observe when the
  dissolution process is endothermic?

22
Calorimetry

• For an exothermic process, the heat produced
  causes the temperature of the solution to
  increase. (DTsoln gt0)

Chemical particle
q
q
q
q
Heat released by chemical particle
q
q
q
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Calorimetry

• For an endothermic process, the heat gained comes
  from the reaction mixture and causes the
  temperature of the solution to decrease (DTsoln lt
  0)

q
Chemical particle
q
q
q
Heat gained by chemical particle from the rxn
mixture
q
q
24
Calorimetry

• Bomb Calorimetry
• Constant Volume Calorimetry
• Used to study combustion reactions
• Measure DHcombustion

25
Calorimetry

• As combustion occurs,
• Heat is released
• Heat is absorbed by the calorimeter and its
  contents
• Temperature of calorimeter contents increases.
• The change in temperature of the calorimeter and
  its contents can be used to determine the heat of
  combustion.

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Calorimetry

• To calculate the molar heat of combustion from a
  calorimetry experiment
• Calculate the heat absorbed by the calorimeter
  and its contents using the heat capacity of the
  calorimeter
• qcal Ccal x DT
• (Notice that you do not need a mass term because
  heat capacity has units of J/K)
• Calculate the heat lost by the reactants
• qrxn -qcal

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Calorimetry

• Calculate molar heat of combustion (DHcomb per
  mole reactant)
• DHcomb qrxn
• mole reactant

28
Calorimetry

• Example When 2.00 g of methylhydrazine (CH6N2)
  is burned in a bomb calorimeter, the temperature
  of the calorimeter increases from 25.00oC to
  32.25oC. If the heat capacity of the calorimeter
  is 7.794 kJ/oC, what is the molar heat of
  combustion for CH6N2?

2 CH6N2 (l) 5 O2 (g) ? 2 N2 (g) 2 CO2(g)
6 H2O (g)
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Calorimetry
30
Hesss Law

• The heats of reaction (DHrxn) have been measured
  and tabulated for many chemical reactions.
• There are two approaches to determining the heat
  of reaction for a particular chemical reaction
• Calorimetry
• Use tabulated DHrxn to calculate the heat of
  reaction for another reaction of interest

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Hesss Law

• The enthalpy change for a reaction or process is
  a state function
• Depends only on the amount of reactants and
  products used/formed and on their physical state
• Does not depend on how the reaction was done.
• one step vs. multiple steps

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Hesss Law

• Hesss Law
• If a reaction is carried out in a series of
  steps, DH for the overall (one-step) reaction is
  equal to the sum of the DHs for the individual
  steps.

33
Hesss Law

• IMPORTANT
• When applying Hesss Law, if you need to multiply
  or divide the coefficients of an equation by a
  number then the DH must also be multiplied by the
  same number
• DH depends on amount of material
• H2O (g) ? H2O (l) DH - 44 kJ
• 2 H2O (g) ? 2 H2O (l) DH 2 (- 44 kJ) -
  88 kJ

34
Hesss Law

• IMPORTANT
• If a reaction has to be reversed, the magnitude
  of DH stays the same but the sign must be
  reversed.
• H2O (g) ? H2O (l) DH - 44 kJ
• H2O (l) ? H2O (g) DH 44 kJ

35
Hesss Law

• Example Calculate the DHrxn for the incomplete
  combustion of C forming CO
• C (s) ½ O2 (g) ? CO (g) DH ???
• given the following reactions
• C (s) O2 (g) ? CO2 (g) DH - 393.5 kJ
• 2 CO (g) O2 (g) ? 2 CO2 (g) DH - 566.0 kJ

36
Hesss Law
37
Hesss Law

• Example Calculate DHrxn for
• NO (g) O (g) ? NO2 (g) DH ???
• using the following thermochemical equations.
• NO (g) O3 (g) ? NO2 (g) O2 (g) DH -198.9
  kJ
• O3 (g) ? 3/2 O2 (g) DH -142.3 kJ
• O2 (g) ? 2 O (g) DH 495.0 kJ

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Hesss Law