Vanier College Faculty of Science and General Studies



GENERAL CHEMISTRY II CHEM206 /4 02

Final Examination April 26, 2004 0930-1230

Dr. Cerrie ROGERS

x periodic table and formula sheet provided

x non-programmable calculators allowed

Chem 206 --- GENERAL CHEMISTRY II

LAST NAME: ____________________________ FIRST NAME: ____________________________

STUDENT NUMBER: _____________________ SIGNATURE: ____________________________

Instructions: PLEASE READ THIS PAGE WHILE WAITING TO START!

• Make sure your exam has 16 pages including this cover page.

• Write your student ID number on all pages.

• Write all answers in the space provided (use examination booklets for rough work only).

• The final page contains optional bonus questions (MAXIMUM GRADE = 107/100).

• Read ALL questions carefully BEFORE starting the exam.

• Non-programmable calculators are allowed.

• You may detach the periodic table and “potentially useful information” pages.

Mark breakdown:

Page 2. / 15 Page 9. / 8

Page 3. / 10 Page 10. / 3

Page 4. / 8 Page 11. / 8

Page 5. / 10 Page 12. / 9

Page 6. / 5 Page 13. / 8

Page 7. / 8 Page 14. / 6 BONUS

Page 8. / 9

TOTAL: / 100 (MAXIMUM MARK = 107)

# 1. (__/ 15 marks) Circle the word(s) that correctly completes each of the following statements.

a) When a piece of hot metal is placed into a beaker of cold water, the sign of the heat flow according to the metal is ( POSITIVE / NEGATIVE ).

b) The second law of thermodynamics states that the ( ENTHALPY / ENTROPY) of the universe increases when a spontaneous process occurs.

c) To remove salt from sea water by reverse osmosis, the external pressure applied to the solution must be ( GREATER / LESS ) than the osmotic pressure exerted by the solution.

d) A reaction that displays overall ( FIRST / SECOND ) order kinetics can be conveniently described

using the term “half life” because the ( RATE / HALF-LIFE ) is independent of concentration.

e) A catalyst ( DECREASES / INCREASES ) the rate of a chemical reaction because it decreases the

reaction’s ( GIBBS FREE ENERGY / ACTIVATION ENERGY ).

f) When a system reaches equilibrium, the forward reaction’s rate is ( GREATER THAN / EQUAL TO ) the reverse reaction’s rate, and the Gibbs free energy for the reaction is ( ZERO / SMALL ).

g) Le Châtelier’s principle is a short-cut based on understanding the ( KINETICS / THERMODYNAMICS ) involved in chemical equilibria.

h) The base dissociation constant or Kb for a base is the ( EQUILIBRIUM / RATE ) constant for the reaction of the base with ( ACID / WATER ).

i) The percent dissociation of an acid is ( LARGER / SMALLER ) in dilute solutions than it is in concentrated solutions.

j) The conjugate base of water is ( HYDRONIUM ION / HYDROXIDE ION ).

k) The equilibrium constant for the reaction of OH- with H3O+ is ( 10-14 / 1014 ).

l) A weak acid whose conjugate base is a different colour is useful as a(n) ( BUFFER / INDICATOR ).

m) A buffer is most effective when the members of the conjugate weak acid - base pair are present in

( LOW / HIGH ) concentration, and in a ratio near ( 1:1 / 10:1 ).

n) The ( SOLUBILITY PRODUCT / SOLUBILITY ) of a substance is the concentration of that substance in a solution that is in equilibrium with undissolved solute.

o) Adding a common ion to a salt solution ( INCREASES / DECREASES ) the solubility of the salt.

|SUBSTANCE |ΔHof (kJ/mol) |Sof (J/mol(K) |

|NH4NO3 (s) |-365.6 |151.1 |

|N2O(g) | 82.0 |219.7 |

|H2O(g) |-241.8 |188.7 |

# 2. (__/ 10 marks) Ammonium nitrate is dangerous because it decomposes, sometimes explosively (as occurred recently in North Korea), when heated:

NH4NO3(s) ( N2O(g) + 2H2O(g)

a) Using logic only, explain whether the entropy of the system increases or decreases during this reaction.

b) Using the thermodynamic data provided, show that this reaction is spontaneous at 25°C.

c) Does the driving force for the reaction increase or decrease when the temperature is raised?

d) Calculate the equilibrium constant at 25°C.

[NOTE: Since both products are gases, you will find Kp directly, and won’t have to convert between Kc and Kp.]

# 3. (__/ 8 marks) Assume that a bottle of wine consists of an 11% by weight (i.e., mass %) solution of ethanol in water (ethanol = CH3CH2OH).

a) What types of intermolecular forces exist between the molecules of water and ethanol?

Circle your choice(s): London dispersion forces, hydrogen bonding, dipole-dipole interactions.

Of these three types of forces, which is the strongest? ___________________________________

b) If the bottle of wine is cooled in a freezer down to -20°C, will the wine begin to freeze?

Show detailed calculations to support your answer. [H2O: freezing point = 0°C; Kf = 1.86 °C(kg(mol-1]

# 4. (__/ 10 marks) The balanced chemical reaction for the oxidation of hydrogen bromide by oxygen is:

4HBr(g) + O2(g) ( 2H2O(g) + 2Br2(g)

The following mechanism has been proposed:

STEP 1: HBr + O2 ( HOOBr

STEP 2: HOOBr + HBr ( 2HOBr

STEP 3: HOBr + HBr ( H2O + Br2

a) Would this mechanism yield the correct stoichiometry overall? (hint: step(s) may need to occur more than once)

b) Write the rate law for each elementary step.

STEP 1: RATE =

STEP 2: RATE =

STEP 3: RATE =

c) Identify any intermediates in the mechanism:

d) By experiment, the reaction was found to be first order with respect to both HBr and O2. Neither HOBr nor

HOOBr was detected among the products. Based on this information, which step must be rate-limiting?

e) HBr does not react with oxygen at a measurable rate at room temperature under ordinary conditions. What

can you infer from this about the magnitude of the activation energy for the rate determining step? Explain.

# 5. (__/ 5 marks) We learned about chemical equilibria AFTER studying kinetics for good reasons.

a) Use KINETICS arguments (i.e., a discussion of reaction rates) to explain the dynamic nature of equilibrium.

b) Next, using KINETICS ARGUMENTS, explain how a reaction mixture that is NOT initially at equilibrium will eventually reach a state of equilibrium. [Hint: reaction rates depend on reactant concentrations….]

# 6. (__/ 8 marks) The reaction between the air pollutants ozone and nitric oxide is important in the formation of the brown “smog” seen in cities (nitrogen dioxide is brown):

O3(g) + NO(g) [pic] O2(g) + NO2(g) K = 6.0(1034

a) Determine whether or not the system is at equilibrium when the gas concentrations in the air are:

[O3] = 1.7(10-6 M, [NO] = 2.3(10-5 M, [NO2] = 3.1(10-4 M, [O2] = 8.2(10-3 M. If it is not at equilibrium, in

which direction will the reaction proceed until it does reach equilibrium?

c) During a nice sunny day in late April, the air temperature changes significantly during the day. Will the concentrations of oxygen and nitrogen dioxide be higher or lower in the warm afternoon than in the cool morning? [Assume that nothing is added or removed, & the only chemistry we have to think about is the reaction shown above.]

|SUBSTANCE |ΔHof (kJ/mol) |

|NO(g) |90.29 |

|NO2(g) |33.1 |

|O2(g) |0 |

|O3(g) |142.67 |

You must justify your answer using calculations involving the thermodynamic data at the right, and explain your conclusion using Le Châtelier’s principle.

# 7. (__/ 9 marks) You are asked to prepare a 100.0 mL sample of a solution with pH of 5.50 by dissolving the appropriate amount of a solute in pure water (pH 7.00). Which ONE of the following solutes would you use, and in what quantity? EXPLAIN YOUR CHOICE, AND SHOW ALL RELEVANT CALCULATIONS.

|SUBSTANCE |Ka |

|HCl |very large |

|HC2H3O2 |1.8(10-5 |

|NH4+ |5.6(10-10 |

|H2O |Kw = 1(10-14 |

CIRCLE YOUR CHOICE:

a) 15 M NH3(aq)

b) 12 M HCl(aq)

c) NH4Cl(s)

d) Pure (“glacial”) acetic acid, HC2H3O2

# 8. (__/ 8 Marks) You are doing a titration to quantify the amount of lactic acid (HC3H5O3, Ka = 1.4(10-4) in a 25.00 mL sample. The titrant solution you are using is 8.00(10-2 M NaOH.

a) If you require 31.35 mL of NaOH titrant to reach the equivalence point, what was the original concentration of lactic acid in your sample?

b) At the equivalence point, the sample is essentially a solution of sodium lactate. What is the pH?

# 8. (__/ 3 Marks) CONTINUED FROM PREVIOUS PAGE…

c) Which of the indicators in the table would minimize indicator error for titrations of other samples like the one from parts a & b ? Briefly explain what indicator error is and the reason for your choice of indicator.

|INDICATOR |Colour of conjugate acid |pH range of |Colour of conjugate base |

| | |colour change | |

|Crystal violet |Green |0.0 - 1.8 |Violet |

|Bromthymol blue |Green |6.0 - 8.0 |Blue |

|Phenolphthalein |Colourless |8.0 - 9.0 |Pink |

|Alizarin yellow |Yellow |10.0 - 11.8 |Red |

# 9. (__/ 8 marks) You have prepared a buffer solution by adding 4.95 g of sodium acetate, NaC2H3O2, to 2.50(102 mL of 0.150 M acetic acid, HC2H3O2 (assume volume unchanged). The Ka of acetic acid is 1.8(10-5.

a) What is the pH of this buffer?

b) If you then bubble 35.5 mg of HCl gas into this solution, what will the new pH be?

# 10. (__/ 9 Marks) Imagine that your bathwater becomes saturated with magnesium palmitate, a component of bathtub soap-scum, while you take a nice soapy bath in 50°C water. What mass of magnesium palmitate, Mg(C16H31O2)2, will precipitate from 115 L of your bathwater when it is cooled from 50°C to 25°C?

[For Mg(C16H31O2)2, Ksp = 4.8(10-12 at 50°C and 3.3(10-12 at 25°C.]

# 11. (__/ 8 Marks) The main compound in marble, which has been widely used for statues and ornamental work on buildings, is CaCO3 (Ksp = 2.8(10-9). Unfortunately, marble is readily attacked by acids because of the basicity of the carbonate ion.

Determine the solubility of marble:

a) In normal rainwater, which has pH = 5.60

b) In acid rainwater, which has pH = 4.20

BONUS QUESTIONS

B1. (__/ 1 mark) Even though the carbonic acid-hydrogen carbonate buffer system is crucial to maintaining the “correct” pH of blood, it has no practical use as a laboratory buffer solution. Can you think of a reason for this?

B2. (__/ 5 marks) Phosphorous acid, H3PO3, is a diprotic acid (Ka1 = 3.7(10-2, Ka2 = 2.1(10-7).

a) Propose a Lewis structure for phosphorous acid that is consistent with the fact that it is diprotic.

b) Draw the structure of phosphoric acid, H3PO4 (Ka1 = 7.1(10-3, Ka2 = 6.8(10-8, Ka3 = 4.2(10-13).

c) Can you explain why phosphorous acid is a stronger acid than phosphoric acid?

------------------------------------------------------ HAVE A GOOD SUMMER ! ---------------------------------------------------

POTENTIALLY USEFUL INFORMATION

R = 8.314 J∙mol-1K-1 = 0.08206 L∙atm∙mol-1K-1

C(H2O) = 4.184 J∙g-1K-1

ΔG0 = ΔH0 - TΔS0

ΔG = ΔG0 + RT lnQ

ΔG0 = -RT lnKeq

PV = nRT

C = k P

P =χ P0

ΔT = K m

k = A e(-Ea/RT)

[A] = -k t + [A]0

ln[A] = -k t + ln[A]0

(1/[A]) = k t + (1/[A]0)

x = -b ± √(b2-4ac)

2a

pH = -log[H3O+]

pH = pKa + log [A-] .

[HA]

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