Problem Set 1.1

1. A home uses 4000 m3 of natural gas (methane, CH4) annually for heating purposes. Assumethat methane behaves as an ideal gas under the conditions of temperature and pressure which arerelevant, i.e., 1 atm and 25o C. What is the mass of the methane gas used per year ?

2. 10 moles of C2H6 gas are confined to a volume of 4.86 L at 27o C. Calculate the pressureexerted by the ethane based on (i) the ideal gas law, and (ii) the Van der Waals equation of state,given that: a = 5.507 L2 atm mol-2 and b = 0.0651 L mol-1 for ethane.

3. The critical volume and critical pressure of a newly discovered gas are 160 cm3 mol-1 and 40atm, respectively. Estimate the critical temperature by assuming that the gas obeys the Berthelotequation of state. Estimate the radii of the individual gas molecules, assuming they are spherical.

4. Consider a spherical balloon, of radius 10 m, filled with hydrogen (H2) to an internal pressureof 1 atm. Assume that (i) hydrogen obeys the ideal gas law, and (ii) the ambient temperature atwhich the balloon is filled is 25o C. The density of air at this temperature is approximately 1.22kg m-3 (at sea level). Calculate (i) the mass of hydrogen filling the balloon, and (ii) the mass theballoon can lift under these conditions (“payload”). (Hint: This problem requires using theprinciple of Archimedes).

5. What is the mean speed of N2 molecules in air at a temperature of 25° C ?

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Problem Set 1

(These are simply EXAMPLE problems: they are NOT due to be handed in as arequirement for credit, but are purely meant as illustrative, sample, “worked examples”)

1. To what temperature must a 1.0 Liter sample of a perfect gas be cooled from 25o C in order toreduce its volume to 100 cm3 (when the pressure is held fixed) ?

2. A car tire was inflated to a pressure of 24 lb in -2 (1 atm = 14.7 lb in-2) on a winter’s day whenthe temperature was –5o C. What pressure will be found, assuming (i) there have been no leaks,and (ii) the volume has remained constant, on a summer day when the temperature is 37o C ?

3. Calculate the pressure exerted by 1 mole of C2H6 as predicted by (i) the perfect gas model, and(ii) the Van der Waals equation of state, when it is confined to a volume of 100 cm3 and is at atemperature of 1000 K. The Van der Waals “a” and “b” constants for C2H6 are: a = 5.507 L2 atmmol-2, and b = 0.0651 L mol-1.

4. 92.4 kg of nitrogen (N2) is heated to a temperature of 500 K at a constant volume of 1 m 3.Calculate the approximate pressure of the gas under these conditions from the Van der Waalsequation of state, assuming that: a = 1.352 L2 atm mol-2, and b = 0.0387 L mol-1 for nitrogen.

5. The critical constants of methane (CH4) are: pc = 45.6 atm, Vc = 98.7 cm3 mol-1, and Tc = 190.6K. Calculate from these data the Van der Waals parameters for methane, and use these toestimate the approximate radius of an individual CH4 molecule.

6. Determine the ratios of: (i) the mean speeds, and (ii) the mean kinetic energies, of H2

molecules and Hg atoms at 20o C.

7. A spherical weather balloon had a radius of 1.0 m when released at sea level at a temperatureof 20o C, and expanded to a radius of 3.0 m when it had risen to maximum altitude where thetemperature was –30o C. What is the pressure within the balloon at that altitude ?

8. Consider the following proposed equation of a state for a newly-discovered substance:

P = R T / Vm - B / Vm2 + C / Vm

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Show that this proposed equation does indeed predict (or, in other words, is capable ofdescribing) critical behavior/condensation.

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Problem Set 2.1

1. A sample of 1 mole of H2O(g) is condensed isothermally and reversibly to liquid water at 100o

C and 1 atm. The standard enthalpy of vaporization of water at 100o C is 40.66 kJ mol-1. Find w,q, ∆U, and ∆H for this process (condensation).

2. A sample of 65 g of xenon is initially confined in a container at 2.5 atm and 298 K, and thenallowed to expand adiabatically: (i) reversibly to a final pressure of 1 atm, and: (ii) against afixed external pressure of 1 atm. Calculate the final temperature in each case, assuming thatxenon is an ideal gas with Cv,m = 3 R/2. (In case (ii), assume that the expansion ceases when thepressure of the gas becomes equal to the applied external pressure of 1 atm).

3. Table sugar consists of the disaccharide sucrose, molecular formula C12H22O11. (i) Calculatethe energy released as heat when a typical sugar cube of mass 1 g is burned in air. (ii) To whatheight could you climb on the energy released by complete combustion of such a sugar cube,assuming that 20 % of the energy released is available for muscular work ? (The standardenthalpy for combustion for sucrose is –5645 kJ mol-1).

4. Consider the following reaction:

6 CH4(g) (methane) → C6H6(g) (benzene) + 9 H2(g).

(The benzene molecule has 6 carbon atoms in a planar ring, and each carbon atom isbonded to 2 other carbon atoms and 1 hydrogen atom).

(a)- Given that the heats of combustion, ∆Hc(0), for methane, benzene, and hydrogen gas

are –890, -3300, and –286 kJ/mol, respectively, estimate the standard enthalpy change forthe above reaction. (In this context, “combustion” = complete reaction with oxygen toyield carbon dioxide and water as the sole products).

(b)- From the standard reaction enthalpy you found in Part (a), estimate the standardenergy change ∆U(0) for this reaction. (Standard conditions are P = 105 Pascals and T =298 K). Assume ideal behavior for all gases.

(c)- Assume that the energy of this reaction, ∆U(0), is purely due to the making andbreaking of chemical bonds. Given that the energies to BREAK any of the C-H bonds inthe problem is 418 kJ/mol, and that to break H-H bonds is 436 kJ/mol, estimate thestrength of the C-C bond in benzene. (Treat ALL the 6 C-C bonds in benzene as beingequivalent).

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Problem Set 2

(Solutions to this problem set need NOT be turned in for credit; these are intended asexamples/sample problems )

1. A sample of 5.5 g of methane initially occupies 14 L at 310 K. (i) Calculate the work donewhen the gas expands isothermally against a constant external pressure of 0.25 atm until itsvolume has increased by 3.3 L. (ii) Calculate the work that would have been done had the sameexpansion been performed reversibly (and isothermally).

2. Consider a system consisting of 3 moles of oxygen (O2) at 25o C confined within a cylinder ofcross sectional area 25 cm2 at 800 kPa. The gas expands adiabatically and irreversibly against aconstant external pressure of 110 kPa. Assume oxygen to be an ideal gas. Calculate w, q, ∆ U, ∆H, and ∆ T when the piston has moved through a distance of 20 cm.

3. From the following data, calculate the standard enthalpy of formation of diborane, B2H6(g), at298 K:

B2H6(g) + 3 O2(g) → B2O3(s) + 3 H2O(g) ∆ Hrxn = -1941 kJ mol-1

2 B(s) + 3/2 O2(g) → B2O3(s) ∆ Hrxn = -2368 kJ mol-1

H2(g) + ½ O2(g) → H2O(g) ∆ Hrxn = -242 kJ mol-1

4. Consider the formation of diamond from amorphous carbon (graphite) at a pressure of 500kbar (1 bar = 105 Pascal). The densities of diamond and graphite are, respectively, 3.52 g cm-3 and2.27 g cm-3. By how much does ∆ U differ from ∆ H for the transition between graphite →diamond under these conditions ?

5. 1 mole of an ideal gas expands adiabatically from a volume of V = 10 L and a temperature of400 K to a volume of 20 L. Assume that Cv,m = 5 R/2 and Cp,m = 7 R/2. (a)- Find the final temperature if the expansion is reversible. (b)- Find the final temperature if the expansion is performed at constant external pressure of 1

atmosphere. (c)- Find the final temperature if the gas expands into a vacuum. (d)- Find ∆U and δw for each of the processes in (a) to (c) above.

6. Evaluate ∆H if 1 mole of supercooled liquid water at -100 C freezes irreversibly at a constantpressure of 1 atm to ice at -100 C. The molar heat capacities of liquid water and of ice are 76 J K-1

mol-1 and 37 J K-1 mol-1 , and the latent heat of fusion for ice at 00 C is 333 J gm-1.

7. Evaluate ∆H for the formation of methane (CH4) from its elements, given that the heats ofcombustion of C(graphite), H2(g), and CH4(g) are -393.5, -285.8, and -890.4 kJ mol-1,respectively.

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Problem Set 3.1

1. The enthalpy of vaporization of chloroform (CHCl3) is 29.4 kJ mol-1 at its normal boiling pointof 334.9 K. Calculate (i) the molar entropy of vaporization of chloroform at this temperature, and(ii) the entropy change of the surroundings.

2. One mole of a monatomic ideal gas initially at 298 K and occupying 10 L is caused to expand,with the surroundings maintained at 298 K, to a final volume of 20 L, in three ways: (i)isothermally and reversibly, (ii) isothermally against a fixed external pressure of 0.5 atm, and (iii)adiabatically against an external pressure which is held fixed during the process of expansion at0.5 atm. Calculate ∆S, ∆Senv, ∆H, ∆T, ∆A, and ∆G, for each path, where the data permit.

3. Calculate the standard enthalpies and entropies at 298 K and 398 K for the reaction: CO2(g) +H2(g) → CO(g) + H2O(g). Assume that the heat capacities are independent of temperature. (Thestandard molar enthalpies of formation for H2O, CO, and CO2, are: -241.8 kJ mol-1, -110.5 kJmol-1, and –393.5 kJ mol-1, respectively. The standard molar entropies for H2O, CO, H2, and CO2

are: 188.8 J/K, 197.7 J/K, 130.7 J/K, and 213.7 J/K, respectively, and their molar heat capacitiesat constant pressure are: 33.6 J K-1 mol-1, 29.1 J K-1 mol-1, 28.8 J K-1 mol-1, and 37.1 J K-1 mol-1,respectively).

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Problem Set 4

1. Calculate the change in Gibbs free energy of 35 g of ethanol (density 0.789 g cm -3) when thepressure is increased isothermally from 1 atm to 30,000 atm.

2. Calculate the change in molar Gibbs free energy of an ideal gas when its pressure is increasedisothermally from 1.5 atm to 37.3 atm at 50o C.

3. The vapor pressure of a liquid between -100oC and 0o C fits the expression: ln (p (in torr)) =16.26 – 2502/T, where T is in Kelvin. Calculate the enthalpy of vaporization of the liquid. (1 torr= 133.32 Pascal).

4. An open vessel containing mercury stands in a sealed room measuring 5m x 5m x 6m at 25o C.What mass of mercury will be found in the air ? The vapor pressure of mercury at thistemperature is 1.7 x 10-3 torr. (1 torr = 133.32 Pascal).

5. Calculate the melting point of ice under a pressure of 100 atm. Assume that the densities of iceand liquid water are 0.92 g cm-3 and 1.0 g cm-3, respectively, and that both phases areincompressible. The standard enthalpy of fusion for ice equals 6.01 kJ mol-1.

6. The temperature dependence of the vapor pressure of solid sulphur dioxide is given by: log (p(in torr)) = 10.59 – 1871 / T, and that of liquid sulphur dioxide by: log(p (in torr)) = 8.32 – 1426 /T, where T is in Kelvin. Estimate the temperature and pressure of the triple point of sulphurdioxide.

7. The enthalpy of vaporization of freon-12 at its normal boiling point of –29o C is 20.3 kJ mol-1.Estimate the pressure that a spray can using freon-12 would have had to withstand at atemperature of 40o C. (Assume that at least some of the freon-12 in the container must remain inthe liquid state).

8. A certain liquid has an enthalpy of vaporization of 19.7 kJ mol -1 at its normal boiling point of192 K. The molar volumes of the liquid and vapor phases at the normal boiling point are 115 cm3

mol-1 and 15.1 L mol-1, respectively. Estimate the slope dP/dT for the vapor-liquid co-existencecurve in the vicinity of the normal boiling point.

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Problem Set 5

1. The vapor pressure of pure benzene is 400 Torr at 60.6o C, and it fell to 386 Torr when 19 g ofan involatile organic compound was dissolved in 500 g of benzene. Calculate the molecularweight of the compound.

2. The addition of 5 g of a compound to 250 g of naphthalene (C10H8) lowered the freezing pointof the solvent by 0.78 K. Calculate the molecular weight of the compound. The freezing pointand standard enthalpy of fusion for pure C10H8 are 354 K and 18.8 kJ mol-1, respectively.

3. The osmotic pressure of an aqueous solution is 120 kPa at 288 K. Estimate the boiling andfreezing points of the solution (assume the solute to be non-volatile). The molar enthalpies forvaporization and fusion for pure H2O, measured at the transition temperatures, are 40.66 kJ mol-1

and 6.01 kJ mol-1, respectively.

4. An aqueous solution of Th(NO3)4 (molal concentration: 9.6 x 10-3 mol kg-1) exhibited afreezing point depression of 0.0703 K. What is the apparent/effective number of ions present insolution per formula unit of Th(NO3)4 ?

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Problem Set 6

1. At 90o C, the vapor pressure of methylbenzene is 400 Torr and that of 1,2 dimethylbenzene is150 Torr. What are: (i) the composition of a liquid mixture of these two compounds that boils at90o C when the pressure is 0.5 atm, and: (ii) the composition of the vapor produced ? (1atmosphere = 760 Torr).

2. At 300 K, the vapor pressures of the pure liquids A and B are 575 and 390 Torr, respectively.Assume that this pair of compounds forms ideal mixtures in both the liquid and vapor phases.For a two-phase mixture in which the mole fraction of A in the vapor is 0.35, calculate (i) thetotal pressure of the vapor, and (ii) the composition of the liquid phase.

3. At a temperature of 273 K, a mixture of hexane and benzene separates into 2 liquidlayers/phases. The mole fraction of benzene in the hexane-rich layer is 0.09, and the molefraction of hexane in the benzene-rich layer is 0.05. Calculate the amounts (by mass) of bothhexane and benzene present in each layer formed upon mixing 50 g of hexane with 100 g ofbenzene at 273 K.

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Problem Set 7

1. The standard enthalpy of a certain reaction is approximately constant at + 125 kJ mol -1 from800 K up to 1500 K. The standard reaction Gibbs free energy is + 22 kJ mol -1 at 1100 K.Estimate the temperature at which the equilibrium constant equals 1.

2. The equilibrium constant of a certain reaction is found to be described by the equation:

ln (Keq) = -1.04 – 1088 / T + 1.51x105 / T2 ,

where T is in Kelvin. Calculate the standard reaction enthalpy and entropy at 400 K.

3. What is the standard enthalpy change for a reaction for which the equilibrium constant is (i)doubled, (ii) halved, when the temperature is increased from 298 K to 308 K ?

4. Consider the decomposition of methane (CH4(g)) into the elements H2(g) and C(s) (graphite). Thestandard enthalpy and entropy for the formation of CH4(g) equal –74.9 kJ mol-1 and –80.7 J K-1

mol-1, at 298 K, respectively. Calculate (i) the equilibrium constant at 298 K, (ii) the equilibriumconstant at 100o C, and (iii) the degree of dissociation of methane at 298 K and a total pressure of0.01 atm.

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Problem Set 8: Electrochemistry

(Notation: “m” = Solute concentration in molal units, moles of solute/kg of solvent)

1. Calculate the ionic strength of a solution that is 0.1 mol kg-1 in KCl (aq) and 0.2 molkg-1 in CuSO4(aq). Assume complete dissociation of the solutes.

2. Calculate the ionic activity coefficients and ionic activities (using the Debye-Huckel limiting law) at 25o C for each of the ionic species in an aqueous solutionthat is 0.01 mol kg-1 in CaCl2(aq) and 0.03 mol kg-1 in NaF(aq). Assume completedissociation of the solutes.

3. Devise cells for which the following are the cell reactions, and calculate thestandard EMFs of the resulting cells:

(a) Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)

(b) 2 AgCl(s) + H2(g) → 2 HCl(aq) + 2 Ag(s)

(c) 2 H2(g) + O2(g) → 2 H2O(l)

(Standard reduction potentials: Cu+2(aq) + 2 e- → Cu(s) : E(0) = +0.34 V; Zn+2

(aq) +2 e- → Zn(s) : E(0) = -0.76 V; AgCl(s) + e- → Ag(s) + Cl-

(aq) : E(0) = +0.22 V; O2(g)

+ 4 H+(aq) + 4 e- → 2 H2O(l) : E(0) = +1.23 V)

4. Using data from Problem (3) above, calculate the EMF (at 25o C) for thefollowing cell: Zn(s) | ZnSO4(aq) (m=0.002) || CuSO4(aq) (m=0.001) | Cu(s) . Assumethat the salt bridge (denoted by the “||”) makes the liquid junction potentialnegligibly small.

5. Calculate the equilibrium constants of the following reactions at 25o C from thestandard potential data that are provided:

(a) Sn(s) + Sn+4(aq) ↔ 2 Sn+2

(aq)

(b) Sn(s) + 2 AgCl(s) ↔ SnCl2(aq) + 2 Ag(s)

(Standard reduction potentials: Sn+4(aq) + 2 e- → Sn+2

(aq) : E(0) = +0.15 V; Sn+2(aq)

+ 2 e- → Sn(s) : E(0) = -0.14 V; AgCl(s) + e- → Ag(s) + Cl-(aq) : E(0) = +0.22 V)

6. The solubility (in water) of AgBr is 2.6 x 10-6 mol kg-1 at 25o C. What is the EMFof the cell: Ag(s) | AgBr(aq) | AgBr(s) | Ag(s) at that temperature ?

7. The standard EMF (E(0)cell) of the cell: Ag(s) | AgI(s) | AgI(aq) | Ag(s) is +0.95 V at 25o

C. Calculate the solubility constant of AgI at this temperature. The solubility

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constant is denoted: Ksp, and is defined as the product: Ksp = −+ IAgaa , where the

activity coefficients refer to the Ag+ and I- ions for a saturated aqueous solution ofAgI. From the solubility constant, estimate the solubility of AgI in water at thistemperature (in moles of AgI per kg of water).

8. Show that the solubility (denoted S) of a sparingly soluble 1:1 salt (e.g., AgCl), inwater at 25o C, is related to the solubility constant for the salt (denoted Ksp) by:

.172.1 Ssp eKS =

9. Given the standard reduction potentials: AgCl(s) + e- → Ag(s) + Cl-(aq) : E(0) = +0.22

V; Zn+2(aq) + 2 e- → Zn(s) : E(0) = -0.76 V, use the Debye-Huckel limiting law to

estimate the EMF of the cell: Zn(s) | ZnCl2(aq) (m = 0.0102), AgCl(s) | Ag(s) at 25o C.

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FCH 360: Fall, 2012EXAM I (1 hour)

USEFUL DATA: Ideal Gas Constant, R = 0.08206 L atm K-1 mol-1 = 8.314 J K-1 mol-1;Boltzmann constant, kB = 1.3807x10-23 J K-1; Avogadro’s Number, NA = 6.022x1023 mol-1; 1atmosphere = 1.013x105 Pascals; 1 m3 = 103 L; 1 L = 103 cm3

Atomic weights: O = 16 gm mol-1, C = 12 gm mol-1

You may assume that for ideal gases: Cv,m = 3 R/2, and Cp,m = 5 R/2

1. Assume that oxygen (O2) is described accurately by the following equation of state:

( ) ( ) 3mm VT

a

bV

TRP −

−= ,

where “a” and “b” are substance-specific parameters, and R is the ideal gas constant.

A sample of one mole of oxygen occupies a volume of 5000 cm3 at a temperature of 200 K undera pressure of 1 atm. Assume that b = 3.19 x 10-2 L mol-1. Use these data to (i) calculate the valueof the parameter a, and (ii) to estimate the pressure exerted by a sample of two moles of oxygenconfined to a volume of 20 L at a temperature of 300 K. (35)

2. Calculate the average speed of molecules of CO2(g) at a temperature of 400 K. (15)

3. Calculate the average kinetic energy per mole for an ideal gas at a temperature of 298 K. (15)

4. A sample of two moles of an ideal gas is initially at a temperature of 300 K and occupies avolume of 20 Liters. The gas undergoes a reversible and isothermal expansion such that the finalpressure equals one-half of the initial pressure that the gas was confined under. Calculate (i) thework done by the gas during this expansion, (ii) the final pressure of the gas, and (iii) the heattransfer (Q) during the process. (35)

FCH 360: FALL 2012: EXAM II (1 hour)

USEFUL DATA: Ideal Gas Constant, R = 0.08206 L atm K-1 mol-1 = 8.314 J K-1 mol-1;Avogadro’s Number, NA = 6.022x1023 mol-1; 1 atm = 1.013x105 Pascals; Atomic weights: H: 1gram mol-1; O: 16 gram mol-1; C: 12 gram mol-1; S: 32 gram mol-1

You may assume that, for an ideal gas: Cv,m = 3 R / 2, and: Cp,m = 5 R / 2

1. Consider the reaction: 2 SO2(g) + O2(g) → 2 SO3(g). The standard enthalpies of formation forSO2(g) and SO3(g) are: -296.8 kJ mol-1 and -395.7 kJ mol-1, respectively. The absolute entropies permole under standard conditions for SO2(g), SO3(g), and O2(g) are: 248.2 J K-1 mol-1, 256.8 J K-1 mol-

1, and: 205.1 J K-1 mol-1, respectively. Based on this information, calculate each of: ∆ G, ∆ H, and∆ S for the above reaction, under standard conditions. (25)

2. Three moles of an ideal gas are initially at a temperature of 300 K and confined under apressure of 2 atmospheres. The gas undergoes an expansion such that the final temperature andpressure are 200 K and 0.5 atmospheres, respectively. Calculate ∆ H and ∆ S for the gas for thisprocess. (25)

3. Consider the equilibrium between carbon dioxide (CO2) in the solid phase (CO2(s), “dry ice”)and carbon dioxide in the gas phase (CO2(g)). The vapor pressure of solid CO2(s) equals 1atmosphere at a temperature of -78o C, and equals 5.1 atmospheres at a temperature of -56o C.Based on these data, estimate the molar enthalpy of sublimation for CO2. (You may assume thatCO2(g) may be treated as an ideal gas, and that the enthalpy of sublimation is independent of thetemperature and pressure). (30)

4. Under a pressure of 1 atmosphere, naphthalene (C10H8(s)) melts at a temperature of 80.3o C. Thelatent heat of fusion (enthalpy of fusion) at this temperature equals 18.8 kJ mol -1. Calculate thechange in entropy (of the naphthalene) when 10 grams of naphthalene are melted from the solidto the liquid state at a pressure of 1 atmosphere and at a constant temperature of 80.3o C. (20)

FCH 360: Fall, 2012

FINAL EXAMINATION

TIME: 2 hours

Attempt ANY FIVE of the problems

USEFUL DATA: Ideal gas constant, R = 8.314 J/mole-K = 0.08206 L atm/mole-K;Avogadro’s Number, NA = 6.022 x 1023 /mole; Boltzmann constant, kB = 1.381 x 10-23

J/K; 1 atm = 1.013 x 105 Pascals; Faraday Constant, F = 96500 Coulomb mol-1

Abbreviations: gm = gram, mol = mole, L = Liter

Atomic weights: He = 4 gm/mol, H = 1 gm/mol, N = 14 gm/mol, O = 16 gm/mol, I =127 gm/mol

Assume that all vapors and gases obey the ideal gas equation of state unless stated otherwise

1. The escape velocity at the surface of the earth equals 11.2 x 103 m s-1. Estimate thetemperatures at which (i) atoms of helium (He), and (ii) molecules of oxygen (O2), have amean speed equal to the escape velocity from the gravitational field of the earth. (20)

2. Calculate the changes in entropy when:(i) 100 grams of liquid water (H2O) are heated from a temperature of 25o C to a

temperature of 50o C at a fixed pressure of 1 atm (for water, Cp,m = 75.4 J K-1 mol-1,independent of temperature);

(ii) 1 mole of oxygen (initially at P = 1 atm and T = 298 K) is mixed with 4 molesof nitrogen (initially at P = 1 atm and T = 298 K), and the final state of the mixture isspecified by T = 298 K and P = 25 atm. (20)

3. The osmotic pressure of a solution of salt in water is found to equal 10 atm at atemperature of 298 K. Estimate the temperature at the boiling point for this solutionunder a pressure of 1 atm. (The density of water may be assumed to be 1 gm/mL; theboiling point for pure water is 373 K under a pressure of 1 atm; and the enthalpy ofvaporization for pure water equals 40.6 kJ mol-1). (20)

4. For the element iodine (I2), the triple point is located at T = 387 K and P = 12.1 x 103

Pascal. The critical point for iodine corresponds to T = 819 K and P = 11.7 x 106 Pascal.Use these data, and the Clapeyron equation, to estimate the molar enthalpy ofvaporization for iodine. Experimentally, it is found that the enthalpy of vaporization foriodine is 41.7 kJ mol-1 at its normal boiling point (T = 457 K and P = 1 atm). Provide atleast two reasons for why your answer differs from the experimentally measured value. (20)

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5. Consider the formation of nitrogen dioxide (NO2(g)) from N2(g) and O2(g).

(i) Write a balanced equation for the formation of NO2(g) from N2(g) and O2(g).

(ii) The standard enthalpy of formation for NO2(g) equals 34 kJ mol-1, and the standardGibbs free energy (at T = 298 K) for formation of NO2(g) equals 51 kJ mol-1. Based uponthese data, calculate the equilibrium constants at both T = 298 K and at T = 1500 K forthe reaction that is your answer to Part (i) above. (You may assume that the enthalpy offormation for NO2(g) is independent of temperature). (20)

6. Use the following standard reduction potentials to estimate the equilibrium constant forthe following reaction at T = 25o C: Mg(s) + Cu+2

(aq) → Mg+2(aq) + Cu(s).

Mg+2(aq) + 2 e → Mg(s) E(0) = -2.36 Volts

Cu+2(aq) + 2 e → Cu(s) E(0) = +0.34 Volts

(20)

7. Use the Debye-Huckel limiting law to calculate the activities and activity coefficientsfor each of the ionic species present in a solution prepared by dissolving 0.001 moles ofMgCl2(s) and 0.002 moles of NaCl(s), in one kilogram of water at T = 25o C. (You mayassume that the MgCl2 and NaCl undergo complete dissociation into Mg+2, Na+, and Cl-

ions). (20)

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