EQUILIBRIUM QUESTIONS



LAB-BASED QUESTIONS -- ANSWERS

|Year |Form A |Form B |

|2011 |#2 -- QUANTITATIVE |# 5 – ESSAY |

| |Gravimetric Analysis |Titration |

| |(Stoichiometry) |(Acid-base) |

|2010 |#2 -- QUANTITATIVE |#5 -- ESSAY |

| |Enthalpy of Solution (Thermodynamics) |Titration |

| | |(Acid-base) |

|2009 |#2 -- QUANTITATIVE |#2 -- QUANTITATIVE |

| |Molar Mass of an Unknown Gas (Gases) |Rate of Reaction |

| | |(Kinetics) |

|2008 |#2 -- QUANTITATIVE |#5 -- ESSAY |

| |Formula of a Hydrate |Quantitative Analysis |

| |(Stoichiometry) |(Descriptive Chemistry/Reactions) |

|2007 |#5 -- ESSAY |#5 -- ESSAY |

| |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 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 |28.7210 g |

| Mass of filter crucible and precipitate |29.3587 g |

| (first weighing) | |

| Mass of filter crucible and precipitate |29.2599 g |

| (second weighing) | |

| Mass of filter crucible and precipitate |29.2598 g |

| (third weighing) | |

(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

[pic]= 0.003759 mol or 3.759 mmol AgCl = 3.759 mmol Ag

(c) [pic]= 62.51% Ag

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.

[pic]

(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[pic] of urea in kJ mol-1.

(d) Using the information in the table below, calculate the value of the molar entropy of solution, [pic] of urea at 298 K. Include units with your answer.

| |Accepted Value |

|(H[pic] of urea |14.0 kJ mol-1 |

|[pic] of urea |–6.9 kJ mol-1 |

(e) The student repeats the experiment and this time obtains a result for (H[pic]of urea that is 11 percent below the accepted value. Calculate the value of (H[pic] 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[pic]? 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

(d) [pic] = (H[pic] - T[pic]; -6900 J mol-1 = 14000 J mol-1 – (298K)([pic])

[pic] = +70. J mol-1 K-1

(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 ( [pic]= 0.995 g air

(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

(d) n = [pic]0.0343 mol

1.37 g gas / 0.0343 mol = 40.0 g mol-1

(e) [pic]( 100 = -9.04% error

(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

1.630 g MgCl2 [pic]= 0.01712 mol MgCl2 ; 1.848 g water [pic] = 0.1026 mol water

[pic] = 6; therefore, the formula is MgCl2•6H2O

(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 [pic]= 0.0191 MgCl2

(ii) 0.0191 mol MgCl2[pic] = 1.82 g MgCl2 ; [pic]( 100 = 61.9% MgCl2

2007 B

5 Fe2+(aq) + MnO4–(aq) + 8 H+(aq) ( 5 Fe3+(aq) + Mn2+(aq) + 4 H2O(l)

5. 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) [pic]( 100

(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.

[pic]

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.

| |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 |

(i) 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) [pic]

(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.

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