# First Year Physical Chemistry Tutorials by sdfsb346f

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First Year Physical Chemistry Tutorials

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```									                    First Year Physical Chemistry Tutorials

Michaelmas Term Week 3

Thermodynamics I : Entropy, Enthalpy & Internal energy

There are three tasks for this week;
(1) Reading – minimum required reading is the two chapters of Atkins listed
below (supplement this as necessary from the other references, and your
lecture notes so far)
(2) Questions for discussion in tutorial; please think about these questions and
make notes on your answers as necessary - you should come to the tutorial
(3) Numerical problems – please attempt as many of these as possible. Hand in
your answers to my pigeon hole in Merton by 5.00 on Tuesday 28th.

-Physical Chemistry, Chapters 2 - 3, P.W. Atkins and J. De Paula, 8th ed., OUP
-Thermodynamics of Chemical Processes, Chapters 1-3, G. Price, Oxford Chemistry
Primer #56, OUP, 1998.
-Basic Chemical Thermodynamics, Chapters 1-3, E.B. Smith, OUP.

Questions for discussion at tutorial

(1) What is the definition of enthalpy, H, and why is this quantity more commonly
used in chemistry than the internal energy U?

(2) What is meant by the following: (a) adiabatic change (b) isothermal change (c)
reversible change (d) state function?

(3) (a) How is entropy defined (i) thermodynamically, (ii) in statistical terms?
(b) What role does S play in determining the direction of spontaneous change?
(c) On cold nights, water spontaneously freezes to form ice, yet the entropy of ice is lower
than that of water. How is this change consistent with your answer to (b)?
(d) In a certain sense both the following statements (i) and (ii) are true - think about how one
or both of these could be rewritten to clarify their meaning.
(i) “In a reversible process the entropy change is dq/T ”
(ii) “In a reversible process there is no change in entropy.”

(4) Would you expect the entropy changes of the following systems to be positive or negative?
(i) A sample of water is heated from 300K to 325K at constant pressure.
(ii) Nitrogen gas is compressed at constant temperature.
(iii) A sample of argon gas at 1 atm pressure is allowed to mix with a sample of krypton at the
same pressure, such that the total pressure remains at 1 atm.
(iv) One mole of water forms one mole of ice at 273K.
(v) The following reactions occur at constant pressure;
(i) 2AgCl(s) + Br2(l)  2AgBr(s) + Cl2(g)
(ii) N2(g) + 3H2(g)  2NH3(g)
(vi) Crystalline sodium chloride dissolves in water.

(5) How do the entropy and enthalpy of a substance vary with temperature? (You
should be able to derive relationships for these starting with definitions of the heat
capacity and the 2nd law of thermodynamics)

Numerical Problems
1.  (a) Calculate the work done when 1 mole of an ideal gas (initial volume V1 ) expands
isothermally and reversibly to a final volume V2 = 3V1 at 298K.
(b) Calculate the work done when 1 mole of an ideal gas (initial volume V1) expands
isothermally into an evacuated space to a final volume V2 = 3V1 at 298K.
(c) The initial and final states of the gas are the same in parts (a) and (b) so that the change in
internal energy is the same in the two cases. How can the different results for (a) and (b) be
reconciled with the First Law.
(d) One mole of CaCO3 was heated in an open vessel at 1 atm pressure to 700C when it
decomposed into CaO(s) and CO2(g). Calculate the work done during the decomposition
assuming that CO2(g) can be regarded as an ideal gas.

2. A 0.825g sample of   benzoic acid was ignited in a bomb calorimeter in the presence of
excess oxygen. The temperature of the calorimeter rose by 1.940 K from 298 K. In two
separate experiments in the same apparatus, 0.498 g of fumaric acid and 0.509 g of maleic
acid were ignited and gave temperature rises of 0.507 K and 0.528 K respectively. In all
experiments a drop of water was added to the calorimeter.

Using the data given below:
(i) Determine the heat capacity of the calorimeter.

(ii) For both fumaric and maleic acids, calculate

(a) the molar internal energy of combustion;
(b) the molar enthalpy of combustion
(c) the molar enthalpy of formation

Comment on the difference between the enthalpies of formation of the two isomers.

(iii) Why would it have been necessary to add a drop of water to the calorimeter?

The standard enthalpy of formation of water is 285.8 kJ mol1 and of CO2 393.5 kJ mol1.
The internal energy of combustion of benzoic acid is 3251 kJ mol1. The relative molecular
masses of benzoic, fumaric and maleic acids are 122, 116 and 116 respectively.
3.  At 298K, the standard enthalpy of formation (Hf) of NH3(g) is 46.11 kJ mol1.
Assuming that the molar heat capacities can be represented by expressions of the form: Cp,m =
a + bT with the coefficients given below, calculate Hf at 1000K.

N2                   H2               NH3
1     1
a / J K mol              28.58                 27.28             29.75
103 b / J K2 mol1         3.77                 3.26               25.1

(Hint: you need to perform an integration)

4. The following table gives  the molar heat capacity of lead over a range if temperatures.
What is the standard molar Third Law entropy of lead at 25C?

T/K            10    15     20     25     30      50     70     100    150    200    250    298
Cp,m / J K1 mol1    2.8   7.0   10.8   14.1   16.5    21.4   23.3    24.5   25.3   25.8   26.2   26.6

(Hint: You need to estimate this from the area under a suitable graph)

5. (a) Calculate the change in entropy of a system comprising one mole of water at 10C that
is cooled to 0C and then freezes to form ice at 0C.
Cp(H2O(liq)) =75.3 J K1 mol1; Hfus(H2O) = 6.0 kJ mol1.

(b) Comment on the sign of S you obtained in (a) in the light of the microscopic changes
occurring in the system.

6. (a) Calculate the entropy change of 3 moles of CH4 that is heated from 298K to 1098K at 1
atm pressure. Cp(CH4) / J K1 mol1 = 23.64 + 4.79102T 1.93105T2 over the temperature
range 298  2000 K.

(Hint: Again, a mathematical integration is required)

(b) The entropy change of 2 moles of an ideal gas when it was expanded isothermally from V1
to V2 was found to be 5.595 J K1. Calculate the ratio V2/V1.

If this isothermal expansion takes place with the gas doing no work, what is the total entropy
change of the system plus surroundings? Show that your result is consistent with the second law
of thermodynamics.

7. Calculate the entropy changes when:
(a) 0.5 moles of H2O(l) at 0C is added to 0.5 moles of H2O(l) at 100C in an insulated vessel.
(b) the water is then heated to 100C at 1 atm pressure,
(c) and then evaporated at 100C at 1 atm pressure,
(d) the water vapour so formed is compressed isothermally to half its volume.
(e) and then heated at constant volume to twice its absolute temperature.

Cp(H2O(liq)) = 75.48 J K1 mol1; Cv(H2O(g)) = 25.3 J K1 mol1 and Hvap(H2O) = 40.7 kJ mol1
(all of which may be assumed to be independent of temperature).

8. (a) Calculate the difference in molar entropy (i) between water at 5C and ice at 5C
and (ii) between water at 95C and steam at 95C and 1 atm pressure.
(b) Calculate the entropy change of the surroundings and hence (c) the total entropy change of
the universe for the two cases. Discuss the spontaneity of transitions between phases at these
temperatures. The difference in molar heat capacities on melting is 37.3 J K1 mol1 and on
vaporisation is 41.9 J K1 mol1. Further Hmelt(273) = 6.01 kJ mol1 and Hvap(373) = 40.7 kJ
mol1.

9. A sample of ideal gas initially at 1 atm pressure and 273K is expanded to a volume greater
by 36.6% in four different ways:
(a) reversibly and isothermally;