LAB-BASED QUESTIONS -- ANSWERS
Year
Form A
Form B
#2 -- QUANTITATIVE
# 5 – ESSAY
Gravimetric Analysis
Titration
(Stoichiometry)
(Acid-base)
#2 -- QUANTITATIVE
#5 -- ESSAY
2010
Enthalpy of Solution
Titration
(Thermodynamics)
(Acid-base)
#2 -- QUANTITATIVE
#2 -- QUANTITATIVE
2009
Molar Mass of an Unknown Gas
Rate of Reaction
(Gases)
(Kinetics)
#2 -- QUANTITATIVE
#5 -- ESSAY
2008
Formula of a Hydrate
Quantitative Analysis
(Stoichiometry)
(Descriptive Chemistry/Reactions)
#5
-ESSAY
#5 -- ESSAY
2007
Redox Titration
Acid-Base/Buffers
(Electrochemistry)
(Acid-Base)
***NOTE: Prior to 2007, the test consisted of 8 problems/essays. The lab questions
were spread among the selections.
2011
2011 A
2.
A student is assigned the task of determining the mass percent of silver in an alloy of copper
and silver by dissolving a sample of the alloy in excess nitric acid and then precipitating the
silver as AgCl.
First the student prepares 50. mL of 6 M HNO3.
(a) The student is provided with a stock solution of 16 M HNO3, two 100 mL graduated
cylinders that can be read to ±1 mL, a 100 mL beaker that can be read to ±10 mL, safety
goggles, rubber gloves, a glass stirring rod, a dropper, and distilled H2O.
(i) Calculate the volume, in mL, of 16 M HNO3 that the student should use for preparing
50. mL of 6 M HNO3.
(ii) Briefly list the steps of an appropriate and safe procedure for preparing the 50. mL of 6
M HNO3. Only materials selected from those provided to the student (listed above) may
be used.
(iii) Explain why it is not necessary to use a volumetric flask (calibrated to 50.00 mL ±0.05
mL) to perform the dilution.
(iv) During the preparation of the solution, the student accidentally spills about 1 mL of 16
M HNO3 on the bench top. The student finds three bottles containing liquids sitting near
the spill: a bottle of distilled water, a bottle of 5 percent NaHCO3(aq), and a bottle of
saturated NaCl(aq). Which of the liquids is best to use in cleaning up the spill? Justify
your choice.
Then the student pours 25 mL of the 6 M HNO3 into a beaker and adds a 0.6489 g sample of the
alloy. After the sample completely reacts with the acid, some saturated NaCl(aq) is added to the
beaker, resulting in the formation of an AgCl precipitate. Additional NaCl(aq) is added until no
more precipitate is observed to form. The precipitate is filtered, washed, dried, and weighed to
constant mass in a filter crucible. The data are shown in the table below.
Mass of sample of copper-silver alloy
0.6489 g
Mass of dry filter crucible
Mass of filter crucible and precipitate
(first weighing)
Mass of filter crucible and precipitate
(second weighing)
Mass of filter crucible and precipitate
(third weighing)
28.7210 g
29.3587 g
29.2599 g
29.2598 g
(b) Calculate the number of moles of AgCl precipitate collected.
(c) Calculate the mass percent of silver in the alloy of copper and silver.
Answer:
(a) (i) since nitric acid is monoprotic, (MA)(VA) = (MB)(VB); (16 M)(VA) = (6 M)(50 mL); VA =
19 mL
(ii) (1) Don safety goggles and rubber gloves, (2) measure about 25 mL of distilled water
into one grad. cyl. and 19 mL of nitric acid into the other, (3) pour the water into the beaker
and slowly add the acid, while mixing with stirring rod, (4) pour back into a graduated
cylinder, (5) use dropper to add water to the solution to get to the 50 mL mark.
(iii) the final solution is not that precise, it is 6 M, not 6.00 M and an excess will be used in
this experiment.
(iv) use the bottle of 5% baking soda, NaHCO3, as this will neutralize the acid.
(b) crucible + ppt
29.2598 g
– crucible
28.7210 g
ppt
0.5388 g AgCl
0.5388 g
= 0.003759 mol or 3.759 mmol AgCl = 3.759 mmol Ag
(107.87  35.45)g/mol
(c)
0.003759 mol Ag  107.87g/mol
 100 = 62.51% Ag
0.6489 g sample
2010 A
2. A student performs an experiment to determine the molar enthalpy of solution of urea,
H2NCONH2. The student places 91.95 g of water at 25°C into a coffee-cup calorimeter and
immerses a thermometer in the water.
After 50 s, the student adds 5.13 g of solid urea, also at 25°C, to the water and measures the
temperature of the solution as the urea dissolves. A plot of the temperature data is shown in
the graph below.
(a) Determine the change in temperature of the solution that results from the dissolution of the
urea.
(b) According to the data, is the dissolution of urea in water an endothermic process or an
exothermic process? Justify your answer.
(c) Assume that the specific heat capacity of the calorimeter is negligible and that the specific
heat capacity of the solution of urea and water is 4.2 J g-1 °C-1 throughout the experiment.
(i) Calculate the heat of dissolution of the urea in joules.
(ii) Calculate the molar enthalpy of solution, H osoln of urea in kJ mol-1.
(d) Using the information in the table below, calculate the value of the molar entropy of
o
solution, Ssoln
of urea at 298 K. Include units with your answer.
Accepted Value
H osoln of urea
14.0 kJ mol-1
o
of urea
Gsoln
–6.9 kJ mol-1
(e) The student repeats the experiment and this time obtains a result for H osoln of urea that is 11
percent below the accepted value. Calculate the value of H osoln that the student obtained in
this second trial.
(f) The student performs a third trial of the experiment but this time adds urea that has been
taken directly from a refrigerator at 5°C. What effect, if any, would using the cold urea
instead of urea at 25°C have on the experimentally obtained value of H osoln ? Justify your
answer.
Answer:
(a) 21.8°C - 25.0°C = -3.2°C
(b) endothermic, it took the heat from the surrounding water and the water’s temperature
decreased.
(c) (i) q = mcT = (91.95 g + 5.13 g))(4.2 Jg-1°C-1)(3.2°C) = 1304.7552 J = 1300 J
1304.7552 J / 5.13 g = 254.3382456 = 250 Jg-1
(ii) (250 Jg-1)(60. g mol-1) = 15260.29474 = 15 kJ mol-1
o
o
o
(d) Gsoln
= H osoln - T Ssoln
; -6900 J mol-1 = 14000 J mol-1 – (298K)( Ssoln
)
o
= +70. J mol-1 K-1
Ssoln
(e) (14.0 kJ mol-1)(0.89) = 12.5 kJ mol-1
(f) larger; a colder starting temp of the solid will give a larger T, then a larger q, which results
in a larger H
2009 A
2. A student was assigned the task of determining the molar mass of an unknown gas. The student
measured the mass of a sealed 843 mL rigid flask that contained dry air. The student then flushed the
flask with the unknown gas, resealed it, and measured the mass again. Both the air and the unknown
gas were at 23˚C and 750. torr. The data for the experiment are shown below.
Volume of sealed flask
843 mL
Mass of sealed flask and air
157.70 g
Mass of sealed flask and unknown gas
158.08 g
(a) Calculate the mass, in grams, of the dry air that was in the sealed flask. (The density of dry air is
1.18 g L-1 at 23.0˚C and 750. torr.)
(b) Calculate the mass, in grams, of the sealed flask itself (i.e., if it had no air in it).
(c) Calculate the mass, in grams, of the unknown gas that was added to the sealed flask.
(d) Using the information above, calculate the value of the molar mass of the unknown gas.
After the experiment was completed, the instructor informed the student that the unknown gas was carbon
dioxide (44.0 g mol-1).
(e) Calculate the percent error in the value of the molar mass calculated in part (d).
(f) For each of the following two possible occurrences, indicate whether it by itself could have been
responsible for the error in the student’s experimental results. You need not include any calculations
with your answer. For each of the possible occurrences, justify your answer.
Occurrence 1: The flask was incompletely flushed with CO2(g), resulting in some dry air remaining
in the flask.
Occurrence 2: The temperature of the air was 23.0˚C, but the temperature of the CO2(g) was lower
than the reported 23.0˚C.
(g) Describe the steps of a laboratory method that the student could use to verify that the volume of the
rigid flask is 843 mL at 23.0˚C. You need not include any calculations with your answer.
Answer:
(a) 843 mL 
1.18 g
= 0.995 g air
1000 mL
(b) 157.70 g flask + air
- 0.995 g air
156.71 g flask
(c)
158.08 g flask + gas
-156.71 g flask
1.37 g gas
PV
(d) n =

RT
750torr 760 (0.843 mL) 0.0343 mol
torr
atm
Lgatm
0.0821 molgK
(295.0K)
1.37 g gas / 0.0343 mol = 40.0 g mol-1
(e)
40.0  44.0   100 = -9.04% error
44.0
(f) occurrence 1: yes; dry air, with molar mass of about 28.8 g mol-1, would, if mixed with the
higher molar mass CO2 would give lower results than expected.
occurrence 2: no; if the temperature is less than expected then more of the sample would be
in the flask (as T decreases, n increases) and give too large a calculated result.
(g) measure mass of empty flask
fill with water
measure mass of flask + water
calculate mass of water, look up density of water at 23˚C
calculate volume of water
2008 A
2. Answer the following questions relating to gravimetric analysis.
In the first of two experiments, a student is assigned the task of determining the number of moles of water
in one mole of MgCl2•n H2O. The student collects the data shown in the following table.
Mass of empty container
22.347 g
Initial mass of sample and container
25.825 g
Mass of sample and container after first heating
23.982 g
Mass of sample and container after second heating
23.976 g
Mass of sample and container after third heating
23.977 g
(a) Explain why the student can correctly conclude that the hydrate was heated a sufficient number of
times in the experiment.
(b) Use the data above to
(i) calculate the total number of moles of water lost when the sample was heated, and
(ii) determine the formula of the hydrated compound.
(c) A different student heats the hydrate in an uncovered crucible, and some of the solid spatters out of
the crucible. This spattering will have what effect on the calculated mass of the water lost by the
hydrate? Justify your answer.
In the second experiment, a student is given 2.94 g of a mixture containing anhydrous MgCl2 and KNO3.
To determine the percentage by mass of MgCl2 in the mixture, the student uses excess AgNO3(aq) to
precipitate the chloride ion as AgCl(s).
(d) Starting with the 2.94 g sample of the mixture dissolved in water, briefly describe the steps
necessary to quantitatively determine the mass of the AgCl precipitate.
(e) The student determines the mass of the AgCl precipitate to be 5.48 g. On the basis of this
information, calculate each of the following.
(i) The number of moles of MgC12 in the original mixture
(ii) The percent by mass of MgCl2 in the original mixture
Answer:
(a) a negligible change in mass between the second and third heatings indicates that all the
water has been removed
(b) (i)
25.825 g hydrate + container
–23.977 g anhydrate + container
1.848 g water
(ii)
23.977 g anhydrate + container
–22.347 g container
1.630 g MgCl2
1 mol
1 mol
1.630 g MgCl2 
= 0.01712 mol MgCl2 ; 1.848 g water 
= 0.1026 mol
95.206 g
18.01 g
water
0.1026 mol water
= 6; therefore, the formula is MgCl2•6H2O
0.01712 mol anhydrate
(c) splattering will result in a greater loss of mass that is calculated as water, this will produce a
higher water/anhydrate ratio
(d) add silver nitrate solution dropwise with stirring; stop when no more ppt results
mass filter paper and add AgCl ppt
wash with plenty of water
allow filter paper and ppt to thoroughly dry
mass filter paper and AgCl
subtract mass of original filter paper to obtain mass of AgCl
(e) (i) 5.48 g AgCl 
35.45 g Cl
1 mol Cl 1 mol MgCl2
= 0.0191 MgCl2


143.32 g AgCl 35.45 g Cl
2 mol Cl
(ii) 0.0191 mol MgCl2 
95.206 g
1.82 g MgCl2
= 1.82 g MgCl2 ;
 100 = 61.9% MgCl2
1 mol
2.94 g mixture
2007 B
5.
5 Fe2+(aq) + MnO4–(aq) + 8 H+(aq)  5 Fe3+(aq) + Mn2+(aq) + 4 H2O(l)
The mass percent of iron in a soluble iron(II) compound is measured using a titration based on the
balanced equation above.
(a) What is the oxidation number of manganese in the permanganate ion, MnO4–(aq)?
(b) Identify the reducing agent in the reaction represented above.
The mass of a sample of the iron(II) compound is carefully measured before the sample is dissolved in
distilled water. The resulting solution is acidified with H2SO4(aq). The solution is then titrated with
MnO4–(aq) until the end point is reached.
(c) Describe the color change that occurs in the flask when the end point of the titration has been
reached. Explain why the color of the solution changes at the end point.
(d) Let the variables g, M, and V be defined as follows:
g = the mass, in grams, of the sample of the iron(II) compound
M = the molarity of the MnO4–(aq) used as the titrant
V = the volume, in liters, of MnO4–(aq) added to reach the end point
In terms of these variables, the number of moles of MnO4–(aq) added to reach the end point of the
titration is expressed as M x V. Using the variables defined above, the molar mass of iron (55.85 g
mol-1), and the coefficients in the balanced chemical equation, write the expression for each of the
following quantities
(i) The number of moles of iron in the sample
(ii) The mass of iron in the sample, in grams
(iii) The mass percent of iron in the compound
(e) What effect will adding too much titrant have on the experimentally determined value of the mass
percent of iron in the compound? Justify your answer.
Answer:
(a) +7
(b) Fe2+
(c) once the end point is reached, the pink-purple permanganate titrant is no longer being
reduced (you have run out of iron(II) ions) and the solution will remain a faint pink
(d) (i) M  V  5
(ii) M  V  5  55.85
(iii)
M  V  5  55.85
 100
g
(e) too big; V would be too large and the expression in part (d)(iii) would produce a larger than
expected value.
2007 B form B
5. Answer the following questions about laboratory situations involving acids, bases, and buffer
solutions.
(a) Lactic acid, HC3H5O3, reacts with water to produce an acidic solution. Shown below are the
complete Lewis structures of the reactants.
In the space provided above, complete the equation by drawing the complete Lewis structures of the
reaction products.
(b) Choosing from the chemicals and equipment listed below, describe how to prepare 100.00 mL of a
1.00 M aqueous solution of NH4C1 (molar mass 53.5 g mol-1). Include specific amounts and
equipment where appropriate.
NH4Cl(s)
50 mL buret
100 mL graduated cylinder
100
mL
pipet
Distilled water
100 mL beaker
100
mL
volumetric
flask
Balance
(c) Two buffer solutions, each containing acetic acid and sodium acetate, are prepared. A student adds
0.10 mol of HCl to 1.0 L of each of these buffer solutions and to 1.0 L of distilled water. The table
below shows the pH measurements made before and after the 0.10 mol of HCl is added.
(i)
pH Before HCl
Added
pH After HCl
Added
Distilled water
7.0
1.0
Buffer 1
4.7
2.7
Buffer 2
4.7
4.3
Write the balanced net-ionic equation for the reaction that takes place when the HCl is added to
buffer 1 or buffer 2.
(ii) Explain why the pH of buffer 1 is different from the pH of buffer 2 after 0.10 mol of HCl is
added.
(iii) Explain why the pH of buffer 1 is the same as the pH of buffer 2 before 0.10 mol of HCl is
added.
ANSWERS:
(a)
(b) measure 5.35 g of ammonium chloride on the balance
transfer to 100 mL volumetric flask
add about 50 mL of distilled water and swirl until solid is completely dissolved
fill volumetric flask up to the mark on the neck with distilled water, mixing while adding.
(c) (i) CH3COO– + H+  CH3COOH
(ii) buffer 1 has less sodium acetate than buffer 2 and the acetate has been completely
protonated by the acid
(iii) the acetate and acetic acid are in equilibrium with each other in both buffers.