LAB-BASED QUESTIONS
Year
Form A
Form B
#2 -- QUANTITATIVE
#5 -- ESSAY
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.
2010
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.
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
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.
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
Distilled water
Balance
100 mL beaker
100
mL
mL
volumetric
100
pipet
flask
(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.