Unit 8B Acid Base Equilibrium Homework AP
Directions: For each question, provide your work and final answer below the question. Without work, no
credit will be given. Check significant figures and units.
1. A sample of 40.0 milliliters of a 0.100 molar HC2H3O2 solution is titrated with a 0.150 molar NaOH solution.
Ka for acetic acid = 1.810–5
(a) What volume of NaOH is used in the titration in order to reach the equivalence point?
(b)What is the molar concentration of C2H3O2– at the equivalence point?
(c) What is the pH of the solution at the equivalence point?
2. A solution of barium hydroxide is titrated with 0.1–M sulfuric acid and the electrical conductivity of the
solution is measured as the titration proceeds. The data obtained are plotted on the graph below.
Conductivity,

10
20
30
40
50
60
70
80
Millilitres of 0.1–M H2SO4
(a) For the reaction that occurs during the titration described above, write a balanced net ionic equation.
(b)Explain why the conductivity decreases, passes through a minimum, and then increases as the volume of
H2SO4 added to the barium hydroxide is increased.
(c) Calculate the number of moles of barium hydroxide originally present in the solution that is titrated.
Unit 8B Acid Base Equilibrium Homework AP
(d)Explain why the conductivity does not fall to zero at the equivalence point of this titration.
3. A 5.00 gram sample of a dry mixture of potassium hydroxide, potassium carbonate, and potassium chloride is
reacted with 0.100 liter of 2.00 molar HCl solution
(a) A 249 milliliter sample of dry CO2 gas, measured at 22ºC and 740 torr, is obtained from this reaction.
What is the percentage of potassium carbonate in the mixture?
(b)The excess HCl is found by titration to be chemically equivalent to 86.6 milliliters of 1.50 molar NaOH.
Calculate the percentages of potassium hydroxide and of potassium chloride in the original mixture.
4.
MgF2(s)  Mg2+(aq) + 2 F-(aq)
In a saturated solution of MgF2 at 18ºC, the concentration of Mg2+ is 1.2110-3 molar. The equilibrium is
represented by the equation above.
(a) Write the expression for the solubility-product constant, Ksp, and calculate its value at 18ºC.
(b) Calculate the equilibrium concentration of Mg2+ in 1.000 liter of saturated MgF2 solution at 18ºC to
which 0.100 mole of solid KF has been added. The KF dissolves completely. Assume the volume
change is negligible.
(c) Predict whether a precipitate of MgF2 will form when 100.0 milliliters of a 3.0010-3-molar Mg(NO3)2
solution is mixed with 200.0 milliliters of a 2.00l0-3-molar NaF solution at 18ºC. Calculations to
support your prediction must be shown.
(d)At 27ºC the concentration of Mg2+ in a saturated solution of MgF2 is 1.1710-3 molar. Is the dissolving of
MgF2 in water an endothermic or an exothermic process? Give an explanation to support your conclusion.
Unit 8B Acid Base Equilibrium Homework AP
Directions: Read the following passage and answer the questions in full sentences.
A Quick Blurb about Titrations
Titration is used in laboratory medicine to determine unknown concentrations of chemicals of interest in blood
and urine. Pharmacists also use titration in the development of new pharmaceuticals. Titration may also be used
to determine the amount of a certain chemical in food. Often, titration is used to determine fat content, water
content, and concentrations of vitamins. Titration is also used to tell if cheeses and wines have aged enough for
distribution to supermarkets and shops. Of course, the main use of titration is in the laboratory. However,
titrations are not limited to the field of chemistry. Scientists of different disciplines use titrations for a variety of
different reasons.
Oxides
Pure water has a pH of 7.0 (neutral); however, natural, unpolluted rainwater actually has a pH of about 5.6
(acidic).[Recall from Experiment 1 that pH is a measure of the hydrogen ion (H+) concentration.] The acidity of
rainwater comes from the natural presence of three substances (CO2, NO, and SO2) found in the troposphere
(the lowest layer of the atmosphere). As is seen in Table I, carbon dioxide (CO2) is present in the greatest
concentration and therefore contributes the most to the natural acidity of rainwater.
Carbon dioxide reacts with water to form carbonic acid (Equation 1). Carbonic acid then dissociates to give the
hydrogen ion (H+) and the hydrogen carbonate ion (HCO3-) (Equation 2). The ability of H2CO3 to deliver H+ is
what classifies this molecule as an acid, thus lowering the pH of a solution.
(1)
(2)
Nitric oxide (NO), which also contributes to the natural acidity of rainwater, is formed during lightning storms
by the reaction of nitrogen and oxygen, two common atmospheric gases (Equation 3). In air, NO is oxidized to
nitrogen dioxide (NO2) (Equation 4), which in turn reacts with water to give nitric acid (HNO3) (Equation 5).
This acid dissociates in water to yield hydrogen ions and nitrate ions (NO3-) in a reaction analagous to the
dissociation of carbonic acid shown in Equation 2, again lowering the pH of the solution.
(3)
(4)
(5)
Unit 8B Acid Base Equilibrium Homework AP
Unfortunately, human industrial activity produces additional acid-forming compounds in far greater quantities
than the natural sources of acidity described above. In some areas of the United States, the pH of rainwater can
be 3.0 or lower, approximately 1000 times more acidic than normal rainwater. In 1982, the pH of a fog on the
West Coast of the United States was measured at 1.8! When rainwater is too acidic, it can cause problems
ranging from killing freshwater fish and damaging crops, to eroding buildings and monuments. About onefourth of the acidity of rain is accounted for by nitric acid (HNO3). In addition to the natural processes that form
small amounts of nitric acid in rainwater, high-temperature air combustion, such as occurs in car engines and
power plants, produces large amounts of NO gas. This gas then forms nitric acid via Equations 4 and 5. Thus, a
process that occurs naturally at levels tolerable by the environment can harm the environment when human
activity causes the process (e.g., formation of nitric acid) to occur to a much greater extent.
What about the other 75% of the acidity of rain? Most is accounted for by the presence of sulfuric acid (H2SO4)
in rainwater. Although sulfuric acid may be produced naturally in small quantities from biological decay and
volcanic activity (Figure 1), it is produced almost entirely by human activity, especially the combustion of
sulfur-containing fossil fuels in power plants. When these fossil fuels are burned, the sulfur contained in them
reacts with oxygen from the air to form sulfur dioxide (SO2). Combustion of fossil fuels accounts for
approximately 80% of the total atmospheric SO2 in the United States. The effects of burning fossil fuels can be
dramatic: in contrast to the unpolluted atmospheric SO2 concentration of 0 to 0.01 ppm, polluted urban air can
contain 0.1 to 2 ppm SO2, or up to 200 times more SO2! Sulfur dioxide, like the oxides of carbon and nitrogen,
reacts with water to form sulfuric acid (Equation 6).
(6)
Sulfuric acid is a strong acid, so it readily dissociates in water, to give an H+ ion and an HSO4- ion (Equation 7).
The HSO4- ion may further dissociate to give H+ and SO42- (Equation 8). Thus, the presence of H2SO4 causes
the concentration of H+ ions to increase dramatically, and so the pH of the rainwater drops to harmful levels.
(7)
(8)
Acid rain triggers a number of inorganic and biochemical reactions with deleterious environmental effects,
making this a growing environmental problem worldwide.
 Many lakes have become so acidic that fish cannot live in them anymore.
 Degradation of many soil minerals produces metal ions that are then washed away in the runoff, causing
several effects:
o The release of toxic ions, such as Al3+, into the water supply.
o The loss of important minerals, such as Ca2+, from the soil, killing trees and damaging crops.
 Atmospheric pollutants are easily moved by wind currents, so acid-rain effects are felt far from where
pollutants are generated.
Marble and limestone have long been preferred materials for constructing durable buildings and monuments.
The Saint Louis Art Museum, the Parthenon in Greece, the Chicago Field Museum, and the United States
Capitol building are all made of these materials. Marble and limestone both consist of calcium carbonate
(CaCO3), and differ only in their crystalline structure. Limestone consists of smaller crystals and is more porous
Unit 8B Acid Base Equilibrium Homework AP
than marble; it is used more extensively in buildings. Marble, with its larger crystals and smaller pores, can
attain a high polish and is thus preferred for monuments and statues. Although these are recognized as highly
durable materials, buildings and outdoor monuments made of marble and limestone are now being gradually
eroded away by acid rain. How does this happen? A chemical reaction (Equation 9) between calcium carbonate
and sulfuric acid (the primary acid component of acid rain) results in the dissolution of CaCO3 to give aqueous
ions, which in turn are washed away in the water flow.
(9)
This process occurs at the surface of the buildings or monuments; thus acid rain can easily destroy the details on
relief work (e.g., the faces on a statue), but generally does not affect the structural integrity of the building. The
degree of damage is determined not only by the acidity of the rainwater, but also by the amount of water flow
that a region of the surface receives. Regions exposed to direct downpour of acid rain are highly susceptible to
erosion, but regions that are more sheltered from water flow (such as under eaves and overhangs of limestone
buildings) are much better preserved. The marble columns of the emperors Marcus Aurelius and Trajan, in
Rome, provide a striking example: large volumes of rainwater flow directly over certain parts of the columns,
which have been badly eroded; other parts are protected by wind effects from this flow, and are in extremely
good condition even after nearly 2000 years! Even those parts of marble and limestone structures that are not
themselves eroded can be damaged by this process (Equation 9). When the water dries, it leaves behind the ions
that were dissolved in it. When a solution containing calcium and sulfate ions dries, the ions crystallize as
CaSO4 2H2O, which is gypsum. Gypsum is soluble in water, so it is washed away from areas that receive a
heavy flow of rain. However, gypsum accumulates in the same sheltered areas that are protected from erosion,
and attracts dust, carbon particles, dry-ash, and other dark pollutants. This results in blackening of the surfaces
where gypsum accumulates. An even more serious situation arises when water containing calcium and sulfate
ions penetrates the stone's pores. When the water dries, the ions form salt crystals within the pore system. These
crystals can disrupt the crystalline arrangement of the atoms in the stone, causing the fundamental structure of
the stone to be disturbed. If the crystalline structure is disrupted sufficiently, the stone may actually crack. Thus,
porosity is an important factor in determining a stone's durability.
Questions:
1. List two or more ways that you could test the acidity of a sample of rainwater.
2. Write a balanced chemical equation for the dissociation of nitric acid in water. (HINT: Draw an analogy
with Equation 2.)
3. The gaseous oxides found in the atmosphere, including CO2 and NO are nonmetal oxides. What would
happen to the pH of rainwater if the atmosphere contained metal oxides instead?
Unit 8B Acid Base Equilibrium Homework AP
4. At sea level and 25oC, one mole of air fills a volume of 24.5 liters, and the density of air is 1.22x106
g/ml. Compute the mole fraction (i.e.,moles of component /total moles) and molarity of SO2 when the
atmospheric concentration of SO2 is 2.0 ppm
5. One strategy for limiting the amount of acid pollution in the atmosphere is scrubbing. In particular,
calcium oxide (CaO) is injected into the combustion chamber of a power plant, where it reacts with the
sulfur dioxide produced, to yield solid calcium sulfite.
a. Write a balanced chemical equation for this reaction.
b. Approximately one ton, or 9.0x102 kg, of calcium sulfite is generated each year for every person
served by a power plant. How much sulfur dioxide (in moles) is prevented from entering the atmosphere
when this much calcium sulfite is generated? Show your calculation.
c. The final stage in the scrubbing process is to treat the combustion gases with a slurry of solid CaO in
water, in order to trap any remaining SO2 and convert it to calcium sulfite. A slurry is a thick suspension
of an insoluble precipitate in water. Using the solubility guidelines predict whether this stage of the
scrubbing process will produce a slurry (i.e., precipitate) or a solution (i.e., no precipitate) of calcium
sulfite.
d. If MgO, rather than CaO, were used for scrubbing, would the product of the final stage be a slurry or
a solution of magnesium sulfite? (Assume that a very large quantity of magnesium sulfite, relative to the
amount of water, is produced.)
6. In water, H2SO4 can dissociate to yield two H+ ions and one SO42- ion. Write the net ionic equation for
the reaction of calcium carbonate and sulfuric acid.
7. Which is a more durable building material, limestone or marble? Briefly, explain your reasoning.