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Unformatted text preview: 1. a. b. 2. a. b. An unknown compound is found to contain only carbon, hydrogen and oxygen.
Complete combustion of a 10.38 mg sample of the compound yielded 16.01 mg of
CO2 and 4.37 mg of H2O. The molar mass of the compound is 176.1 g/mol.
How many mg of each of C and H were in the original sample?
4.37 mg C, 0.49 mg H
What is the % w/w of each of C, H and O in the compound?
42.10% C, 4.72% H, 53.18% O
What are the empirical and molecular formula of the compound?
C2H3O2 empirical formula, C6H9O6 molecular formula
The mass spectrum of an unknown ketone showed the molecular ion peak at m/z
= 86. The relative abundances of the M to the M+1 peak were in the ratio of
about 100 to 5. The other main peaks appeared at m/z = 71 and 43. Suggest a
possible structure for this ketone and briefly explain your reasoning.
A sample of a mixture of KClO3 and KBr weighing 7.0950 g is heated in the
presence of a small amount of MnO2 which acts as a catalyst of the decomposition
2 KClO3(s) 3O2(g) + 2KCl(s) .
The O2(g) is collected in a container over water. The total volume and pressure of
the gas in the container was 865.5 mL and 741.5 mm Hg respectively at 27.2oC.
The vapour pressure of water at that temperature is 27.1 mm Hg. Assume that all
of the KClO3(s) decomposed. What was the %w/w of KClO3(s) in the original
Briefly describe the role of a catalyst in a reaction. Give two other specific
examples of reactions and their catalysts. 3. In a laboratory experiment, 0.03404 g of Mg was completely reacted with dilute, aqueous
HCl generating H2(g). The H2(g) was collected in a gas burette over the aqueous solution at
26.4oC. The final level of aqueous solution in the burette was 5.00 cm higher than the
level in the beaker. The vapour pressure of the aqueous solution at 26.4oC was 25.8 mm
Hg and the barometric pressure that day was 764.8 mm Hg. Given that the density of the
aqueous solution is 1.00 g/mL and the density of Hg is 13.6 g/mL, calculate the volume
of the H2(g) that was collected. (atomic mass of Mg = 24.30 g/mol)
V = 0.0356 L 4. Calculate the empirical formula of hydrated ferrous ammonium sulfate,
Fea(NH4)b(SO4)c(H2O)d, from the following data.
0.7840 g of the salt gives 0.1600 g Fe2O3(s) when heated strongly in air to constant
0.7840 g of the salt dissolved in water gives 0.9336 g BaSO4(s) when excess
BaCl2(aq) is added.
When 0.3920 g of the salt is dissolved in water and boiled with excess NaOH(aq),
NH3(g) is liberated. When this gas is absorbed in 50.0 mL of 0.10 M HCl(aq), the excess acid remaining after reaction with the NH3(g) requires 30.0 mL of 0.10 M
NaOH for neutralization.
(a=1, b=2, c=2, d=6)
5. M is a metal; X is a non-metal. Upon strong heating of 0.5000g of MXO3(s), it
decomposes forming MO(s) and XO2(g). Then, all of the MO(s) was dissolved in water and
excess acid, 100 mL of 0.120M HCl was added. The resulting solution required 40.0 mL
of 0.050 M NaOH for complete neutralization. The rate of effusion of the XO2(g) was
found to be 0.797 times that of N2(g) at the same temperature.
a. Find the number of moles of MO(s) formed.
b. Identify X.
c. Identify M.
Ca 6. A sample of a mixture of CaC2O4 and MgC2O4 was dissolved in about 100 mL of aqueous
acid. One half of the solution was titrated with 15.05 mL of 0.1000 M KMnO4 to the end
MnO4- + C2O4-2 = Mn+2 + CO2 (unbalanced)
The other half of the solution was titrated with 27.40 mL of 0.1000 M EDTA to the faint
pink murexide indicator end point which only titrates the calcium ions.
Calculate the mass percentage of the original sample that was CaC2O4.
(75.35%) 7. Explain the principles upon which mass spectrometry is based. How could one
distinguish between ethanol and dimethylether using mass spectrometry? 8. Organic reactions can be classified as substitutions, eliminations or additions and the
reagents can be classified as electrophiles, nucleophiles or radicals. Reactions from each
classification are involved in the synthesis of 2,3,4-trimethylpentan-3-ol from ethanol (as
the only source of carbon atoms) and inorganic reagents. Write out the steps involved in
the synthesis. For each step,
write an equation,
name the reactants and products, and
classify the type of reaction (e.g. electrophilic substitution). 9. Briefly explain each of the following facts:
Water has a boiling point about 200 degrees greater than would be expected in
comparison with the boiling points of H2S, H2Se and H2Te.
Water has its maximum density at 3.98oC.
Methanol is soluble with water, while hexane is not.
Na2O is readily soluble in water, while MgO is not. e. 10. The freezing point of 0.1 M BaCl2 is lower than that of 0.1M NaI and both are
lower than that of pure water. a. Consider the reaction by which HBr is added to ethene to give bromoethane.
Write the mechanism of the reaction, using curly arrows to show the movement of
electrons. If the reactant was 2-methylpropene, rather than ethene, what would be
For the reaction above, the following thermodynamic data were determined:
Ho = 84.1 kJ/mol ; So = 132 J/K mol ; Go = 44.8 kJ/mol . Note that
these are the absolute values of the thermodynamic parameters. From the
reaction, explain whether the sign of each of these parameters is expected to be
positive or negative.
All are negative.
Then calculate the equilibrium constant for the reaction at 298 K.
7.12 x 107
The rate law for the reaction above is rate = k1[ethene][HBr] and its activation
energy is 140 kJ/mol. Draw an energy profile for the reaction, plotting potential
energy vs progress of reaction. Label this diagram showing reactants,
intermediates (if any) and products and as much of the data given as is possible. b. c. 11. a. b. 12. For a certain first order reaction, the reaction goes to 38.5% completion in 480 s.
Calculate the rate constant for the reaction. Show the derivation of the equation
that you use for the calculation.
1.01 x 10-3 s-1
At 25oC the first-order rate constant for a certain reaction is 2.0 x 103 s1. The
activation energy is 15.0 kJ/mol. What is the value of the rate constant at 75oC?
Start with the Arrhenius equation, k=PZeEa/RT , and show the derivation of the
equation that you use for the calculation.
4.75 x 103 s-1 F or the r eaction
CH3OH + Br
two mechanisms have been pr oposed.
(1) CH3OH + OH CH3O + H2O CH3O + Br [Br CH3O 2] [Br CH3O 2]
O 2 + H2O
CH3Br + O 2
slow CH3Br + OH the following (2)
CH3OH + H3O+ CH3OH2+ + Br fast [Br CH3OH2] fast 2H2O CH3OH2+ + H2O fast
[Br CH3OH2] CH3Br + H2O H3O+ + OH slow
d. 13. a. b. 14. What would you expect the rate law to be for each of these mechanisms? Show
how you figur e this out.
Br iefly explain how one could experimentally distinguish between the two
Briefly explain which mechanism you think is most likely to occur.
Consider mechanism (2). If the obser ved r ate constant is
kobs = 1. 6 x 10 4sec 1M 2 and the r eaction is car r ied out with
[CH3OH]o = 0. 5M, [H3O+ ]o = 0. 5M, and [Br ]o = 4 x 10 5M, how long will
it take to r educe the [Br ] to 4 x 10 6M?
t = 14391 s Der ive the equations that show that while the half-life for a fir st or der r eaction
is independent of the or iginal r eactant concentr ation, the half-life for both zer oor der and second-or der r eactions depends on the or iginal r eactant concentr ation.
The r adioactive isotope 32P decays accor ding to fir st-or der kinetics and has a
half-life of 14. 3 days. How long does it take for 95. 0% of a sample of 32P to
61. 8 days a. At what temper atur es will each of the pr ocesses with the following values of
enthalpy changes and entr opy changes be spontaneous? Br iefly explain any
H = 25 kJ, S = 5. 0 J K 1
H = + 25 kJ, S = + 5. 0 J K 1
H = + 25 kJ, S = 5. 0 J K 1
not spontaneous at any T
H = 25 kJ, S = + 5. 0 J K 1
spontaneous at all T b. Wr ite the balanced equation for the dissolving of Mg(s) in HCl(aq). How much
wor k is done when 50. 00 g of Mg(s) is dissolved in excess HCl(aq) in a container
of flexible volume at a constant temper atur e of 23 " C? Assume ideal gas
w = -5. 06 kJ/ mol 15. Given the following standar d r eduction potentials, calculate the KSP for AgCl(s).
AgCl(s) + e Ag(s) + Cl(aq)
+ 0. 22
Ag (aq) + e
+ 0. 80
Br iefly descr ibe another way by which the KSP of AgCl could be measur ed.
KSP = 1. 56 x 10 10 b. 16. a. Balance the following r edox r eaction that occur s in aqueous acid:
MnO4 (aq) + F e(s) F e+ 2(aq) + Mn+ 2(aq) .
Dr aw, and label the par ts of, a galvanic cell in which F e is oxidized and MnO4
is r educed. a. Answer this question using the simultaneous equations method. Note any
simplifying assumptions made in the derivation.
What is the pH of a buffer solution that was made by mixing 45.0 mL of 0.750M
propanoic acid (KA = 1.3 x 105 ) with 55.0 mL of 0.700 M sodium propanoate?
pH = 4.94
What volume of 0.100 M NaOH would have to be added to the above buffer
solution in order to increase the pH by 5.0 % ?
97mL b. 17. Given an ample supply of a 0.200 M aqueous solution of the weak, monoprotic acid,
formic acid (KA = 1.78 x 104) and a 0.200 M aqueous solution of NaOH, describe how
you would prepare 500 mL of a buffer solution whose pH = 4.0. Use the simultaneous
equations method for the calculations. Explain any simplifying assumptions made.
Mix 0.305 L of acid with 0.195 L of base.
Describe one other way by which a pH 4.0 formic acid buffer solution could be prepared. 18. a.
b. 19. Given that pKB for NH3 = 4.75 and pKW for H2O = 14.00, calculate the equilibrium constant for the reaction NH3 + H3O+ NH4+ + H2O.
A 25.00 mL sample of a weak base, 0.1063M B, was titrated with a strong acid,
0.1495M HCl. When 10.00 mL of acid were added, the pH was 7.72.
Calculate the KB of the base.
Calculate the pH at the equivalence point of the titration.
Use the simultaneous equations method for the calculations and be sure to explain
any simplifying assumptions made in your calculations. A 25.00 mL sample of 0.1002 M lactic acid, a monoprotic acid, was titrated with 0.1997 M NaOH. When 10.00 mL of base were added, the pH was 4.46.
Calculate the KA of lactic acid.
1.365 x 104
Calculate the pH at the equivalence point of the titration.
Use the simultaneous equations method for your calculations and be sure to explain any
simplifying assumptions made in your calculations.
When about 6.25 mL of base were added, a buffer solution resulted. Explain.
20. In a laboratory practical exam students were asked to prepare an aqueous buffer solution
having a pH close to 9. Which of these students prepared the desired buffer? Note that
more than one used a correct procedure. Briefly explain you conclusion regarding each
student's work. Detailed pH calculations are not required.
KB of NH3 = 1.8 x 10-5M ; KA of HA = 1.8 x 10-5M.
Amanda mixed 50mL 0.10M NH4Cl + 50 mL 0.10M NaOH.
Balal mixed 25 mL 0.10M NaOH + 50 mL 0.10M NH4Cl
Chandra mixed 50 mL 0.10M NH4Cl + 50 mL 0.10M NH3.
Daniella mixed 50 mL NaA + 50 mL 0.10M HA.
Erum mixed 25 mL 0.10M HCl + 50 mL 0.10M NH3.
cor r ect 21. A 2.51 g sample of the monoprotic acid, HA, required 27.36 mL of 0.5106 M NaOH for
complete reaction. Addition of 15.44 mL of 0.4524 M HCl to the flask containing the
HA and the NaOH produced a mixture with pH = 3.48.
Find the molar mass of HA.
Using the simultaneous equations method for the calculations, find the KA value of HA.
3.3 x 10-4 22. a. b. 23. What volume of a 0.10 M sodium acetate (NaA) should be added to 500 mL of
0.10 M acetic acid (HA), whose KA is 1.8 x 10-5, in order to produce a buffer
solution of pH 5.5? Use the simultaneous equations method to solve this problem
and show any simplifying assumptions made.
Describe another way by which an identical buffer solution could be prepared. To 10.00 mL of a weak acid, HA, was added 9.00 mL of 0.085 M NaOH. This
corresponded to the 75% point in the titration. The pH was 6.0. Find the original acid
concentration, Ca, (0.102 M)
and the acid dissociation constant, Ka,
(3.01 x 10-6)
at 25 C. Use the simultaneous equations method for the calculations and explain any
simplifying assumptions made.
24. Describe the range of intermolecular forces which hold together groups of molecules or
ions in the solid, liquid and non-ideal gaseous state (giving specific examples). Compare
the strength of these forces and the distances over which they operate.
Based upon the above, explain the following observations.
The boiling point of dimethyl ether is less than that of ethanol.
At high densities, gases do not obey the ideal gas law.
KF is much more soluble in water than CaO. 25. a.
b. c. What parameters are used to define each of the following electronegativity scales:
Pauling, Mulliken, Allred-Rochow? What do they all have in common?
Using the following data, calculate the relative electronegativities of H, F, Cl and
Bond Dissociation Energy (kJ/mol)
Briefly comment on the observed trends in electronegativities with reference to
atomic orbital energy levels. 26. Using the Bohr theory, derive a general equation for the frequencies of the lines of the
spectrum of atomic hydrogen. Identify the expression for the Rydberg constant in your
equation. 27. The van’t Hoff equation which relates the equilibrium constant of a reaction to the
reaction temperature has the form
ln K = X + Y
Show the derivation of this equation:
from the equations defining the standard free energy change of the reaction and
from the Arrhenius equation for the rate constant of the reaction.
Note any assumptions made in the derivations. Comment on the physical meaning of the
terms Y which result in each derivation and the relationship between them. 28. Beginning with H=E+PV and G=H-TS, derive an equation that relates the standard free
energy change for a reaction to the equilibrium constant for that reaction. Note any
simplifications or assumptions made in the derivation. 29. a. b. 30. Beginning with the Arrhenius equation, k = PZeEa/RT , derive an equation from
which the rate constant for a reaction at one temperature could be calculated if the
rate constant for that reaction at another temperature and the activation energy
were known. Note any simplifications or assumptions made in the derivation.
Explain the difference between the effect of increasing the temperature on the rate
constant for a reaction, k, and the effect of increasing the temperature on the
equilibrium constant for a reaction, K. a. A 1.00 g sample of enriched water, a mixture of H2O and D2O, reacted completely
with Cl2 to give a mixture of HCl and DCl. The HCl and DCl were then dissolved
in pure H2O to make a 1.00 L solution. A 25.00 mL sample of the 1.00 L solution
was reacted with excess AgNO3 and 0.3800 g of an AgCl precipitate formed.
What was the mass % of D2O in the original sample of enriched water?
Atomic masses (g/mol): H = 1.0, D = 2.0, O = 16.0, Cl = 35.5, Ag = 107.9 b. The mass % natural abundance of the isotopes H and D are 99.985 and 0.015
respectively. The mass % natural abundance of the isotopes 16O, 17O and 18O are
99.759, 0.037 and 0.204 respectively. In a mass spectrometry experiment on
naturally occurring water, at what m/z values would you expect to observe the
three most abundant peaks for the molecular ions and what would you expect to
be their relative abundances.
m/z = 18
relative abundance = 100.
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This note was uploaded on 12/11/2011 for the course CHEM 140 taught by Professor P during the Spring '10 term at University of Toronto- Toronto.
- Spring '10