a. Calculate the pH of a buffered solution that is 0.100 M in C6H5CO2H (benzoic acid, Ka

In Section 14.5 we found that the equilibrium concentration of H1 in a 1.0-M HF solution is 2.7 3 1022 M, and the percent dissociation of HF is 2.7%. Calculate [H1] and the percent dissociation of HF in a solution containing 1.0 M HF (Ka 5 7.2 3 1024 ) and 1.0 M NaF.

A buffered solution contains 0.50 M acetic acid (HC2H3O2, Ka 5 1.8 3 1025 ) and 0.50 M sodium acetate (NaC2H3O2). Calculate the pH of this solution.

Calculate the change in pH that occurs when 0.010 mole of solid NaOH is added to 1.0 L of the buffered solution described in Example 15.2. Compare this pH change with that which occurs when 0.010 mole of solid NaOH is added to 1.0 L water.

Calculate the pH of a solution containing 0.75 M lactic acid (Ka 5 1.4 3 1024 ) and 0.25 M sodium lactate. Lactic acid (HC3H5O3) is a common constituent of biologic systems. For example, it is found in milk and is present in human muscle tissue during exertion.

A buffered solution contains 0.25 M NH3 (Kb 5 1.8 3 1025 ) and 0.40 M NH4Cl. Calculate the pH of this solution.

Calculate the pH of the solution that results when 0.10 mole of gaseous HCl is added to 1.0 L of the buffered solution from Example 15.5.

Calculate the change in pH that occurs when 0.010 mole of gaseous HCl is added to 1.0 L of each of the following solutions: Solution A: 5.00 M HC2H3O2 and 5.00 M NaC2H3O2 Solution B: 0.050 M HC2H3O2 and 0.050 M NaC2H3O2 For acetic acid, Ka 5 1.8 3 1025 .

The text states that the pKa for a weak acid to be used in the buffer should be as close as possible to the desired pH. What is the problem with choosing a weak acid whose pKa is not close to the desired pH when making a buffer?

A chemist needs a solution buffered at pH 4.30 and can choose from the following acids (and their sodium salts): a. chloroacetic acid (Ka 5 1.35 3 1023 ) b. propanoic acid (Ka 5 1.3 3 1025 ) c. benzoic acid (Ka 5 6.4 3 1025 ) d. hypochlorous acid (Ka 5 3.5 3 1028 ) Calculate the ratio [HA]y[A2] required for each system to yield a pH of 4.30. Which system will work best?

Hydrogen cyanide gas (HCN), a powerful respiratory inhibitor, is highly toxic. It is a very weak acid (Ka 5 6.2 3 10210) when dissolved in water. If a 50.0-mL sample of 0.100 M HCN is titrated with 0.100 M NaOH, calculate the pH of the solution a. after 8.00 mL of 0.100 M NaOH has been added. b. at the halfway point of the titration. c. at the equivalence point of the titration.

A chemist has synthesized a monoprotic weak acid and wants to determine its Ka value. To do so, the chemist dissolves 2.00 mmol of the solid acid in 100.0 mL water and titrates the resulting solution with 0.0500 M NaOH. After 20.0 mL NaOH has been added, the pH is 6.00. What is the Ka value for the acid?

You have read about titrations of strong acids with strong bases, weak acids with strong bases, and weak bases with strong acids. What if you titrated a weak acid with a weak base? Sketch a pH curve and defend its shape. Label the equivalence point and discuss the possibilities for the pH value at the equivalence point.

Bromthymol blue, an indicator with a Ka value of 1.0 3 1027 , is yellow in its HIn form and blue in its In2 form. Suppose we put a few drops of this indicator in a strongly acidic solution. If the solution is then titrated with NaOH, at what pH will the indicator color change first be visible?

What is meant by the presence of a common ion? How does the presence of a common ion affect an equilibrium such as HNO2 1aq2mH1 1aq2 1 NO2 2 1aq2 What is an acidbase solution called that contains a common ion?

Define a buffer solution. What makes up a buffer solution? How do buffers absorb added H1 or OH2 with little pH change? Is it necessary that the concentrations of the weak acid and the weak base in a buffered solution be equal? Explain. What is the pH of a buffer when the weak acid and conjugate base concentrations are equal? A buffer generally contains a weak acid and its weak conjugate base, or a weak base and its weak conjugate acid, in water. You can solve for the pH by setting up the equilibrium problem using the Ka reaction of the weak acid or the Kb reaction of the conjugate base. Both reactions give the same answer for the pH of the solution. Explain. A third method that can be used to solve for the pH of a buffer solution is the HendersonHasselbalch equation. What is the HendersonHasselbalch equation? What assumptions are made when using this equation?

One of the most challenging parts of solving acidbase problems is writing out the correct equation. When a strong acid or a strong base is added to solutions, they are great at what they do and we always react them first. If a strong acid is added to a buffer, what reacts with the H1 from the strong acid and what are the products? If a strong base is added to a buffer, what reacts with the OH2 from the strong base and what are the products? Problems involving the reaction of a strong acid or strong base are assumed to be stoichiometry problems and not equilibrium problems. What is assumed when a strong acid or strong base reacts to make it a stoichiometry problem?

A good buffer generally contains relatively equal concentrations of weak acid and conjugate base. If you wanted to buffer a solution at pH 5 4.00 or pH 5 10.00, how would you decide which weak acid conjugate base or weak baseconjugate acid pair to use? The second characteristic of a good buffer is good buffering capacity. What is the capacity of a buffer? How do the following buffers differ in capacity? How do they differ in pH? 0.01 M acetic acid/0.01 M sodium acetate 0.1 M acetic acid/0.1 M sodium acetate 1.0 M acetic acid/1.0 M sodium acetate

Draw the general titration curve for a strong acid titrated by a strong base. At the various points in the titration, list the major species present before any reaction takes place and the major species present after any reaction takes place. What reaction takes place in a strong acidstrong base titration? How do you calculate the pH at the various points along the curve? What is the pH at the equivalence point for a strong acidstrong base titration? Why?

Instead of the titration of a strong acid by a strong base considered in Question 5, consider the titration of a strong base by a strong acid. Compare and contrast a strong acidstrong base titration with a strong base strong acid titration.

Sketch the titration curve for a weak acid titrated by a strong base. When performing calculations concerning weak acidstrong base titrations, the general two-step procedure is to solve a stoichiometry problem first, then to solve an equilibrium problem to determine the pH. What reaction takes place in the stoichiometry part of the problem? What is assumed about this reaction? At the various points in your titration curve, list the major species present after the strong base (NaOH, for example) reacts to completion with the weak acid, HA. What equilibrium problem would you solve at the various points in your titration curve to calculate the pH? Why is pH . 7.0 at the equivalence point of a weak acidstrong base titration? Does the pH at the halfway point to equivalence have to be less than 7.0? What does the pH at the halfway point equal? Compare and contrast the titration curves for a strong acidstrong base titration and a weak acidstrong base titration.

Sketch the titration curve for a weak base titrated by a strong acid. Weak basestrong acid titration problems also follow a two-step procedure. What reaction takes place in the stoichiometry part of the problem? What is assumed about this reaction? At the various points in your titration curve, list the major species present after the strong acid (HNO3, for example) reacts to completion with the weak base, B. What equilibrium problem would you solve at the various points in your titration curve to calculate the pH? Why is pH , 7.0 at the equivalence point of a weak basestrong acid titration? If pH 5 6.0 at the halfway point to equivalence, what is the Kb value for the weak base titrated? Compare and contrast the titration curves for a strong basestrong acid titration and a weak basestrong acid titration.

What is an acidbase indicator? Define the equivalence (stoichiometric) point and the end point of a titration. Why should you choose an indicator so that the two points coincide? Do the pH values of the two points have to be within 60.01 pH unit of each other? Explain

Why does an indicator change from its acid color to its base color over a range of pH values? In general, when do color changes start to occur for indicators? Can the indicator thymol blue contain only a single OCO2H group and no other acidic or basic functional group? Explain.

What are the major species in solution after NaHSO4 is dissolved in water? What happens to the pH of the solution as more NaHSO4 is added? Why? Would the results vary if baking soda (NaHCO3) were used instead?

A friend asks the following: Consider a buffered solution made up of the weak acid HA and its salt NaA. If a strong base like NaOH is added, the HA reacts with the OH2 to form A2. Thus the amount of acid (HA) is decreased, and the amount of base (A2) is increased. Analogously, adding HCl to the buffered solution forms more of the acid (HA) by reacting with the base (A2). Thus how can we claim that a buffered solution resists changes in the pH of the solution? How would you explain buffering to this friend?

Mixing together solutions of acetic acid and sodium hydroxide can make a buffered solution. Explain. How does the amount of each solution added change the effectiveness of the buffer?

Could a buffered solution be made by mixing aqueous solutions of HCl and NaOH? Explain. Why isnt a mixture of a strong acid and its conjugate base considered a buffered solution?

Sketch two pH curves, one for the titration of a weak acid with a strong base and one for a strong acid with a strong base. How are they similar? How are they different? Account for the similarities and the differences.

Sketch a pH curve for the titration of a weak acid (HA) with a strong base (NaOH). List the major species, and explain how you would go about calculating the pH of the solution at various points, including the halfway point and the equivalence point.

You have a solution of the weak acid HA and add some HCl to it. What are the major species in the solution? What do you need to know to calculate the pH of the solution, and how would you use this information? How does the pH of the solution of just the HA compare with that of the final mixture? Explain.

You have a solution of the weak acid HA and add some of the salt NaA to it. What are the major species in the solution? What do you need to know to calculate the pH of the solution, and how would you use this information? How does the pH of the solution of just the HA compare with that of the final mixture? Explain.

The common ion effect for weak acids is to significantly decrease the dissociation of the acid in water. Explain the common ion effect

Consider a buffer solution where [weak acid] . [conjugate base]. How is the pH of the solution related to the pKa value of the weak acid? If [conjugate base] . [weak acid], how is pH related to pKa?

A best buffer has about equal quantities of weak acid and conjugate base present as well as having a large concentration of each species present. Explain

Consider the following pH curves for 100.0 mL of two different acids with the same initial concentration each titrated by 0.10 M NaOH.a. Which plot represents a pH curve of a weak acid, and which plot is for a strong acid? How can you tell? Cite three differences between the plots that help you decide. b. In both cases the pH is relatively constant before the pH changes greatly. Does this mean that at some point in each titration each solution was a buffered solution? c. True or false? The equivalence point volume for each titration is the same. Explain your answer. d. True or false? The pH at the equivalence point for each titration is the same. Explain your answer

An acid is titrated with NaOH. The following beakers are illustrations of the contents of the beaker at various times during the titration. These are presented out of order. Note: Counter-ions and water molecules have been omitted from the illustrations for clarity. (a) (b) (c) (d) (e) a. Is the acid a weak or strong acid? How can you tell? b. Arrange the beakers in order of what the contents would look like as the titration progresses. c. For which beaker would pH 5 pKa? Explain your answer. d. Which beaker represents the equivalence point of the titration? Explain your answer. e. For which beaker would the Ka value for the acid not be necessary to determine the pH? Explain your answer.

Consider the following four titrations. i. 100.0 mL of 0.10 M HCl titrated by 0.10 M NaOH ii. 100.0 mL of 0.10 M NaOH titrated by 0.10 M HCl iii. 100.0 mL of 0.10 M CH3NH2 titrated by 0.10 M HCl iv. 100.0 mL of 0.10 M HF titrated by 0.10 M NaOH Rank the titrations in order of: a. increasing volume of titrant added to reach the equivalence point. b. increasing pH initially before any titrant has been added. c. increasing pH at the halfway point in equivalence. d. increasing pH at the equivalence point. How would the rankings change if C5H5N replaced CH3NH2 and if HOC6H5 replaced HF?

Figure 15.4 shows the pH curves for the titrations of six different acids by NaOH. Make a similar plot for the titration of three different bases by 0.10 M HCl. Assume 50.0 mL of 0.20 M of the bases and assume the three bases are a strong base (KOH), a weak base with Kb 5 1 3 1025 , and another weak base with Kb 5 1 3 10210.

Acidbase indicators mark the end point of titrations by magically turning a different color. Explain the magic behind acidbase indicators.

How many of the following are buffered solutions? Explain your answer. Note: Counter-ions and water molecules have been omitted from the illustrations for clarity.

Which of the following can be classified as buffer solutions? a. 0.25 M HBr 1 0.25 M HOBr b. 0.15 M HClO4 1 0.20 M RbOH c. 0.50 M HOCl 1 0.35 M KOCl d. 0.70 M KOH 1 0.70 M HONH2 e. 0.85 M H2NNH2 1 0.60 M H2NNH3NO3

A certain buffer is made by dissolving NaHCO3 and Na2CO3 in some water. Write equations to show how this buffer neutralizes added H1 and OH2

A buffer is prepared by dissolving HONH2 and HONH3NO3 in some water. Write equations to show how this buffer neutralizes added H1 and OH2.

Calculate the pH of each of the following solutions. a. 0.100 M propanoic acid (HC3H5O2, Ka 5 1.3 3 1025 ) b. 0.100 M sodium propanoate (NaC3H5O2) c. pure H2O d. a mixture containing 0.100 M HC3H5O2 and 0.100 M NaC3H5O2

Calculate the pH of each of the following solutions. a. 0.100 M HONH2 (Kb 5 1.1 3 1028 ) b. 0.100 M HONH3Cl c. pure H2O d. a mixture containing 0.100 M HONH2 and 0.100 M HONH3Cl

Compare the percent dissociation of the acid in Exercise 21a with the percent dissociation of the acid in Exercise 21d. Explain the large difference in percent dissociation of the acid

Compare the percent ionization of the base in Exercise 22a with the percent ionization of the base in Exercise 22d. Explain any differences.

Calculate the pH after 0.020 mole of HCl is added to 1.00 L of each of the four solutions in Exercise 21

Calculate the pH after 0.020 mole of HCl is added to 1.00 L of each of the four solutions in Exercise 22.

Calculate the pH after 0.020 mole of NaOH is added to 1.00 L of each of the four solutions in Exercise 21.

Calculate the pH after 0.020 mole of NaOH is added to 1.00 L of each of the solutions in Exercise 22.

Which of the solutions in Exercise 21 shows the least change in pH upon the addition of acid or base? Explain

Which of the solutions in Exercise 22 is a buffered solution?

Calculate the pH of a solution that is 1.00 M HNO2 and 1.00 M NaNO2.

Calculate the pH of a solution that is 0.60 M HF and 1.00 M KF.

Calculate the pH after 0.10 mole of NaOH is added to 1.00 L of the solution in Exercise 31, and calculate the pH after 0.20 mole of HCl is added to 1.00 L of the solution in Exercise 31.

Calculate the pH after 0.10 mole of NaOH is added to 1.00 L of the solution in Exercise 32, and calculate the pH after 0.20 mole of HCl is added to 1.00 L of the solution in Exercise 32.

Calculate the pH of each of the following buffered solutions. a. 0.10 M acetic acid/0.25 M sodium acetate b. 0.25 M acetic acid/0.10 M sodium acetate c. 0.080 M acetic acid/0.20 M sodium acetate d. 0.20 M acetic acid/0.080 M sodium acetate

Calculate the pH of each of the following buffered solutions. a. 0.50 M C2H5NH2/0.25 M C2H5NH3Cl b. 0.25 M C2H5NH2/0.50 M C2H5NH3Cl c. 0.50 M C2H5NH2/0.50 M C2H5NH3Cl

Calculate the pH of a buffered solution prepared by dissolving 21.5 g benzoic acid (HC7H5O2) and 37.7 g sodium benzoate in 200.0 mL of solution.

A buffered solution is made by adding 50.0 g NH4Cl to 1.00 L of a 0.75-M solution of NH3. Calculate the pH of the final solution. (Assume no volume change.)

Calculate the pH after 0.010 mole of gaseous HCl is added to 250.0 mL of each of the following buffered solutions. a. 0.050 M NH3/0.15 M NH4Cl b. 0.50 M NH3/1.50 M NH4Cl Do the two original buffered solutions differ in their pH or their capacity? What advantage is there in having a buffer with a greater capacity?

An aqueous solution contains dissolved C6H5NH3Cl and C6H5NH2. The concentration of C6H5NH2 is 0.50 M and pH is 4.20. a. Calculate the concentration of C6H5NH3 1 in this buffer solution. b. Calculate the pH after 4.0 g NaOH(s) is added to 1.0 L of this solution. (Neglect any volume change.)

Calculate the mass of sodium acetate that must be added to 500.0 mL of 0.200 M acetic acid to form a pH 5 5.00 buffer solution

What volumes of 0.50 M HNO2 and 0.50 M NaNO2 must be mixed to prepare 1.00 L of a solution buffered at pH 5 3.55?

Consider a solution that contains both C5H5N and C5H5NHNO3. Calculate the ratio [C5H5N]y[C5H5NH1] if the solution has the following pH values: a. pH 5 4.50 c. pH 5 5.23 b. pH 5 5.00 d. pH 5 5.50

Carbonate buffers are important in regulating the pH of blood at 7.40. If the carbonic acid concentration in a sample of blood is 0.0012 M, determine the bicarbonate ion concentration required to buffer the pH of blood at pH 5 7.40. H2CO3 1aq2mHCO3 2 1aq2 1 H1 1aq2 Ka 5 4.3 3 1027

When a person exercises, muscle contractions produce lactic acid. Moderate increases in lactic acid can be handled by the blood buffers without decreasing the pH of blood. However, excessive amounts of lactic acid can overload the blood buffer system, resulting in a lowering of the blood pH. A condition called acidosis is diagnosed if the blood pH falls to 7.35 or lower. Assume the primary blood buffer system is the carbonate buffer system described in Exercise 45. Calculate what happens to the [H2CO3]/[HCO3 2] ratio in blood when the pH decreases from 7.40 to 7.35.

Consider the acids in Table 14.2. Which acid would be the best choice for preparing a pH 5 7.00 buffer? Explain how to make 1.0 L of this buffer.

Consider the bases in Table 14.3. Which base would be the best choice for preparing a pH 5 5.00 buffer? Explain how to make 1.0 L of this buffer.

Calculate the pH of a solution that is 0.40 M H2NNH2 and 0.80 M H2NNH3NO3. In order for this buffer to have pH 5 pKa, would you add HCl or NaOH? What quantity (moles) of which reagent would you add to 1.0 L of the original buffer so that the resulting solution has pH 5 pKa?

Calculate the pH of a solution that is 0.20 M HOCl and 0.90 M KOCl. In order for this buffer to have pH 5 pKa, would you add HCl or NaOH? What quantity (moles) of which reagent would you add to 1.0 L of the original buffer so that the resulting solution has pH 5 pKa?

Which of the following mixtures would result in buffered solutions when 1.0 L of each of the two solutions are mixed? a. 0.1 M KOH and 0.1 M CH3NH3Cl b. 0.1 M KOH and 0.2 M CH3NH2 c. 0.2 M KOH and 0.1 M CH3NH3Cl d. 0.1 M KOH and 0.2 M CH3NH3Cl

Which of the following mixtures would result in a buffered solution when 1.0 L of each of the two solutions are mixed? a. 0.2 M HNO3 and 0.4 M NaNO3 b. 0.2 M HNO3 and 0.4 M HF c. 0.2 M HNO3 and 0.4 M NaF d. 0.2 M HNO3 and 0.4 M NaOH

What quantity (moles) of NaOH must be added to 1.0 L of 2.0 M HC2H3O2 to produce a solution buffered at each pH? a. pH 5 pKa b. pH 5 4.00 c. pH 5 5.00

Calculate the number of moles of HCl(g) that must be added to 1.0 L of 1.0 M NaC2H3O2 to produce a solution buffered at each pH. a. pH 5 pKa b. pH 5 4.20 c. pH 5 5.00

Consider the titration of a generic weak acid HA with a strong base that gives the following titration curve: 5 pH 10 15 20 25 mL of base On the curve, indicate the points that correspond to the following: a. the stoichiometric (equivalence) point b. the region with maximum buffering c. pH 5 pKa d. pH depends only on [HA] e. pH depends only on [A2] f. pH depends only on the amount of excess strong base added

Sketch the titration curve for the titration of a generic weak base B with a strong acid. The titration reaction is B 1 H1mBH1 On this curve, indicate the points that correspond to the following: a. the stoichiometric (equivalence) point b. the region with maximum buffering c. pH 5 pKa d. pH depends only on [B] e. pH depends only on [BH1] f. pH depends only on the amount of excess strong acid added

Consider the titration of 40.0 mL of 0.200 M HClO4 by 0.100 M KOH. Calculate the pH of the resulting solution after the following volumes of KOH have been added. a. 0.0 mL d. 80.0 mL b. 10.0 mL e. 100.0 mL c. 40.0 mL

Consider the titration of 80.0 mL of 0.100 M Ba(OH)2 by 0.400 M HCl. Calculate the pH of the resulting solution after the following volumes of HCl have been added. a. 0.0 mL d. 40.0 mL b. 20.0 mL e. 80.0 mL c. 30.0 mL

Consider the titration of 100.0 mL of 0.200 M acetic acid (Ka 5 1.8 3 1025 ) by 0.100 M KOH. Calculate the pH of the resulting solution after the following volumes of KOH have been added. a. 0.0 mL d. 150.0 mL b. 50.0 mL e. 200.0 mL c. 100.0 mL f. 250.0 mL

Consider the titration of 100.0 mL of 0.100 M H2NNH2 (Kb 5 3.0 3 1026) by 0.200 M HNO3. Calculate the pH of the resulting solution after the following volumes of HNO3 have been added. a. 0.0 mL d. 40.0 mL b. 20.0 mL e. 50.0 mL c. 25.0 mL f. 100.0 mL

Lactic acid is a common by-product of cellular respiration and is often said to cause the burn associated with strenuous activity. A 25.0-mL sample of 0.100 M lactic acid (HC3H5O3, pKa 5 3.86) is titrated with 0.100 M NaOH solution. Calculate the pH after the addition of 0.0 mL, 4.0 mL, 8.0 mL, 12.5 mL, 20.0 mL, 24.0 mL, 24.5 mL, 24.9 mL, 25.0 mL, 25.1 mL, 26.0 mL, 28.0 mL, and 30.0 mL of the NaOH. Plot the results of your calculations as pH versus milliliters of NaOH added.

Repeat the procedure in Exercise 61, but for the titration of 25.0 mL of 0.100 M propanoic acid (HC3H5O2, Ka 5 1.3 3 1025 ) with 0.100 M NaOH.

Repeat the procedure in Exercise 61, but for the titration of 25.0 mL of 0.100 M NH3 (Kb 5 1.8 3 1025 ) with 0.100 M HCl.

Repeat the procedure in Exercise 61, but for the titration of 25.0 mL of 0.100 M pyridine with 0.100 M hydrochloric acid (Kb for pyridine is 1.7 3 1029). Do not calculate the points at 24.9 and 25.1 mL.

Calculate the pH at the halfway point and at the equivalence point for each of the following titrations. a. 100.0 mL of 0.10 M HC7H5O2 (Ka 5 6.4 3 1025 ) titrated by 0.10 M NaOH b. 100.0 mL of 0.10 M C2H5NH2 (Kb 5 5.6 3 1024 ) titrated by 0.20 M HNO3 c. 100.0 mL of 0.50 M HCl titrated by 0.25 M NaOH

In the titration of 50.0 mL of 1.0 M methylamine, CH3NH2 (Kb 5 4.4 3 1024), with 0.50 M HCl, calculate the pH under the following conditions. a. after 50.0 mL of 0.50 M HCl has been added b. at the stoichiometric point

You have 75.0 mL of 0.10 M HA. After adding 30.0 mL of 0.10 M NaOH, the pH is 5.50. What is the Ka value of HA?

A student dissolves 0.0100 mol of an unknown weak base in 100.0 mL water and titrates the solution with 0.100 M HNO3. After 40.0 mL of 0.100 M HNO3 was added, the pH of the resulting solution was 8.00. Calculate the Kb value for the weak base

Two drops of indicator HIn (Ka 5 1.0 3 1029 ), where HIn is yellow and In2 is blue, are placed in 100.0 mL of 0.10 M HCl. a. What color is the solution initially? b. The solution is titrated with 0.10 M NaOH. At what pH will the color change (yellow to greenish yellow) occur? c. What color will the solution be after 200.0 mL NaOH has been added?

Methyl red has the following structure: Ka = 5.0 106 CO2H N N N(CH3)2 It undergoes a color change from red to yellow as a solution gets more basic. Calculate an approximate pH range for which methyl red is useful. What is the color change and the pH at the color change when a weak acid is titrated with a strong base using methyl red as an indicator? What is the color change and the pH at the color change when a weak base is titrated with a strong acid using methyl red as an indicator? For which of these two types of titrations is methyl red a possible indicator?

Potassium hydrogen phthalate, known as KHP (molar mass 5 204.22 g/mol), can be obtained in high purity and is used to determine the concentration of solutions of strong bases by the reaction HP2 1aq2 1 OH2 1aq2 h H2O1l2 1 P22 1aq2 If a typical titration experiment begins with approximately 0.5 g KHP and has a final volume of about 100 mL, what is an appropriate indicator to use? The pKa for HP2 is 5.51.

A certain indicator HIn has a pKa of 3.00 and a color change becomes visible when 7.00% of the indicator has been converted to In2. At what pH is this color change visible?

Which of the indicators in Fig. 15.8 could be used for the titrations in Exercises 57 and 59?

Which of the indicators in Fig. 15.8 could be used for the titrations in Exercises 58 and 60?

Which of the indicators in Fig. 15.8 could be used for the titrations in Exercises 61 and 63?

Which of the indicators in Fig. 15.8 could be used for the titrations in Exercises 62 and 64?

Estimate the pH of a solution in which bromcresol green is blue and thymol blue is yellow. (See Fig. 15.8.)

Estimate the pH of a solution in which crystal violet is yellow and methyl orange is red. (See Fig. 15.8.)

A solution has a pH of 7.0. What would be the color of the solution if each of the following indicators were added? (See Fig. 15.8.) a. thymol blue c. methyl red b. bromthymol blue d. crystal violet

A solution has a pH of 4.5. What would be the color of the solution if each of the following indicators were added? (See Fig. 15.8.) a. methyl orange c. bromcresol green b. alizarin d. phenolphthalein

Derive an equation analogous to the HendersonHasselbalch equation but relating pOH and pKb of a buffered solution composed of a weak base and its conjugate acid, such as NH3 and NH4 1.

a. Calculate the pH of a buffered solution that is 0.100 M in C6H5CO2H (benzoic acid, Ka 5 6.4 3 1025 ) and 0.100 M in C6H5CO2Na. b. Calculate the pH after 20.0% (by moles) of the benzoic acid is converted to benzoate anion by addition of a strong base. Use the dissociation equilibrium C6H5CO2H1aq2mC6H5CO2 2 1aq2 1 H1 1aq2 to calculate the pH. c. Do the same as in part b, but use the following equilibrium to calculate the pH: C6H5CO2 2 1aq2 1 H2O1l2mC6H5CO2H1aq2 1 OH2 1aq2 d. Do your answers in parts b and c agree? Explain

Tris(hydroxymethyl)aminomethane, commonly called TRIS or Trizma, is often used as a buffer in biochemical studies. Its buffering range is pH 7 to 9, and Kb is 1.19 3 1026 for the aqueous reaction 1HOCH22 3CNH2 1 H2Om1HOCH22 3CNH3 1 1 OH2 TRIS TRISH1 a. What is the optimal pH for TRIS buffers? b. Calculate the ratio [TRIS]y[TRISH1] at pH 5 7.00 and at pH 5 9.00. c. A buffer is prepared by diluting 50.0 g TRIS base and 65.0 g TRIS hydrochloride (written as TRISHCl) to a total volume of 2.0 L. What is the pH of this buffer? What is the pH after 0.50 mL of 12 M HCl is added to a 200.0-mL portion of the buffer?

You make 1.00 L of a buffered solution (pH 5 4.00) by mixing acetic acid and sodium acetate. You have 1.00 M solutions of each component of the buffered solution. What volume of each solution do you mix to make such a buffered solution?

You have the following reagents on hand: Solids (pKa of Acid Form Is Given) Solutions Benzoic acid (4.19) 5.0 M HCl Sodium acetate (4.74) 1.0 M acetic acid (4.74) Potassium fluoride (3.14) 2.6 M NaOH Ammonium chloride (9.26) 1.0 M HOCl (7.46) What combinations of reagents would you use to prepare buffers at the following pH values? a. 3.0 b. 4.0 c. 5.0 d. 7.0 e. 9.0

Amino acids are the building blocks for all proteins in our bodies. A structure for the amino acid alanine is Amino group Carboxylic acid group C CH3 H C O H2N OH All amino acids have at least two functional groups with acidic or basic properties. In alanine, the carboxylic acid group has Ka 5 4.5 3 1023 and the amino group has Kb 5 7.4 3 1025. Because of the two groups with acidic or basic properties, three different charged ions of alanine are possible when alanine is dissolved in water. Which of these ions would predominate in a solution with [H1] 5 1.0 M? In a solution with [OH2] 5 1.0 M?

Phosphate buffers are important in regulating the pH of intracellular fluids at pH values generally between 7.1 and 7.2. a. What is the concentration ratio of H2PO4 2 to HPO4 22 in intracellular fluid at pH 5 7.15? H2PO4 2 1aq2mHPO4 22 1aq2 1 H1 1aq2 Ka 5 6.2 3 1028 b. Why is a buffer composed of H3PO4 and H2PO4 2 ineffective in buffering the pH of intracellular fluid? H3PO4 1aq2mH2PO4 2 1aq2 1 H1 1aq2 Ka 5 7.5 3 1023

What quantity (moles) of HCl(g) must be added to 1.0 L of 2.0 M NaOH to achieve a pH of 0.00? (Neglect any volume changes.)

Calculate the value of the equilibrium constant for each of the following reactions in aqueous solution. a. HC2H3O2 1 OH2mC2H3O2 2 1 H2O b. C2H3O2 2 1 H1mHC2H3O2 c. HCl 1 NaOHmNaCl 1 H2O

The following plot shows the pH curves for the titrations of various acids by 0.10 M NaOH (all of the acids were 50.0-mL samples of 0.10 M concentration). Vol 0.10 M NaOH added (mL) 10 20 30 40 50 60 2.0 4.0 6.0 8.0 10.0 12.0 0 pH a c d e f b a. Which pH curve corresponds to the weakest acid? b. Which pH curve corresponds to the strongest acid? Which point on the pH curve would you examine to see if this acid is a strong acid or a weak acid (assuming you did not know the initial concentration of the acid)? c. Which pH curve corresponds to an acid with Ka < 1 3 1026 ?

Calculate the volume of 1.50 3 1022 M NaOH that must be added to 500.0 mL of 0.200 M HCl to give a solution that has pH 5 2.15.

Repeat the procedure in Exercise 61, but for the titration of 25.0 mL of 0.100 M HNO3 with 0.100 M NaOH

A certain acetic acid solution has pH 5 2.68. Calculate the volume of 0.0975 M KOH required to reach the equivalence point in the titration of 25.0 mL of the acetic acid solution

A 0.210-g sample of an acid (molar mass 5 192 g/mol) is titrated with 30.5 mL of 0.108 M NaOH to a phenolphthalein end point. Is the acid monoprotic, diprotic, or triprotic?

The active ingredient in aspirin is acetylsalicylic acid. A 2.51-g sample of acetylsalicylic acid required 27.36 mL of 0.5106 M NaOH for complete reaction. Addition of 13.68 mL of 0.5106 M HCl to the flask containing the aspirin and the sodium hydroxide produced a mixture with pH 5 3.48. Determine the molar mass of acetylsalicylic acid and its Ka value. State any assumptions you must make to reach your answer

One method for determining the purity of aspirin (C9H8O4) is to hydrolyze it with NaOH solution and then to titrate the remaining NaOH. The reaction of aspirin with NaOH is as follows: C9H8O4 1s2 1 2OH2 1aq2 Aspirin 888n C7H5O3 2 1aq2 1 C2H3O2 2 1aq2 1 H2O1l2 Salicylate ion Acetate ion A sample of aspirin with a mass of 1.427 g was boiled in 50.00 mL of 0.500 M NaOH. After the solution was cooled, it took 31.92 mL of 0.289 M HCl to titrate the excess NaOH. Calculate the purity of the aspirin. What indicator should be used for this titration? Why?

A student intends to titrate a solution of a weak monoprotic acid with a sodium hydroxide solution but reverses the two solutions and places the weak acid solution in the buret. After 23.75 mL of the weak acid solution has been added to 50.0 mL of the 0.100 M NaOH solution, the pH of the resulting solution is 10.50. Calculate the original concentration of the solution of weak acid.

A student titrates an unknown weak acid, HA, to a pale pink phenolphthalein end point with 25.0 mL of 0.100 M NaOH. The student then adds 13.0 mL of 0.100 M HCl. The pH of the resulting solution is 4.70. How is the value of pKa for the unknown acid related to 4.70?

A sample of a certain monoprotic weak acid was dissolved in water and titrated with 0.125 M NaOH, requiring 16.00 mL to reach the equivalence point. During the titration, the pH after adding 2.00 mL NaOH was 6.912. Calculate Ka for the weak acid.

Consider 1.0 L of a solution that is 0.85 M HOC6H5 and 0.80 M NaOC6H5. (Ka for HOC6H5 5 1.6 3 10210.) a. Calculate the pH of this solution. b. Calculate the pH after 0.10 mole of HCl has been added to the original solution. Assume no volume change on addition of HCl. c. Calculate the pH after 0.20 mole of NaOH has been added to the original buffer solution. Assume no volume change on addition of NaOH.

What concentration of NH4Cl is necessary to buffer a 0.52-M NH3 solution at pH 5 9.00? (Kb for NH3 5 1.8 3 1025 .)

Consider the following acids and bases: HCO2H Ka 5 1.8 3 1024 HOBr Ka 5 2.0 3 1029 (C2H5)2NH Kb 5 1.3 3 1023 HONH2 Kb 5 1.1 3 1028 Choose substances from the following list that would be the best choice to prepare a pH 5 9.0 buffer solution. a. HCO2H e. (C2H5)2NH b. HOBr f. (C2H5)2NH2Cl c. KHCO2 g. HONH2 d. HONH3NO3 h. NaOBr

Consider a buffered solution containing CH3NH3Cl and CH3NH2. Which of the following statements concerning this solution is(are) true? (Ka for CH3NH3 1 5 2.3 3 10211.) a. A solution consisting of 0.10 M CH3NH3Cl and 0.10 M CH3NH2 would have a higher buffering capacity than one containing 1.0 M CH3NH3Cl and 1.0 M CH3NH2. b. If [CH3NH2] . [CH3NH3 1], then the pH is larger than the pKa value. c. Adding more [CH3NH3Cl] to the initial buffer solution will decrease the pH.d. If [CH3NH2] , [CH3NH3 1], then pH , 3.36. e. If [CH3NH2] 5 [CH3NH3 1], then pH 5 10.64.

Consider the titration of 150.0 mL of 0.100 M HI by 0.250 M NaOH. a. Calculate the pH after 20.0 mL of NaOH has been added. b. What volume of NaOH must be added so that the pH 5 7.00?

Consider the titration of 100.0 mL of 0.100 M HCN by 0.100 M KOH at 258C. (Ka for HCN 5 6.2 3 10210.) a. Calculate the pH after 0.0 mL of KOH has been added. b. Calculate the pH after 50.0 mL of KOH has been added. c. Calculate the pH after 75.0 mL of KOH has been added. d. Calculate the pH at the equivalence point.

Consider the titration of 100.0 mL of 0.100 M HCN by 0.100 M KOH at 258C. (Ka for HCN 5 6.2 3 10210.) a. Calculate the pH after 0.0 mL of KOH has been added. b. Calculate the pH after 50.0 mL of KOH has been added. c. Calculate the pH after 75.0 mL of KOH has been added. d. Calculate the pH at the equivalence point. e. Calculate the pH after 125 mL of KOH has been added

Consider the titration of 100.0 mL of 0.200 M HONH2 by 0.100 M HCl. (Kb for HONH2 5 1.1 3 1028 .) a. Calculate the pH after 0.0 mL of HCl has been added. b. Calculate the pH after 25.0 mL of HCl has been added. c. Calculate the pH after 70.0 mL of HCl has been added. d. Calculate the pH at the equivalence point. e. Calculate the pH after 300.0 mL of HCl has been added. f. At what volume of HCl added does the pH 5 6.04?

Consider the following four titrations (iiv): i. 150 mL of 0.2 M NH3 (Kb 5 1.8 3 1025 ) by 0.2 M HCl ii. 150 mL of 0.2 M HCl by 0.2 M NaOH iii. 150 mL of 0.2 M HOCl (Ka 5 3.5 3 1028 ) by 0.2 M NaOH iv. 150 mL of 0.2 M HF (Ka 5 7.2 3 1024 ) by 0.2 M NaOH a. Rank the four titrations in order of increasing pH at the halfway point to equivalence (lowest to highest pH). b. Rank the four titrations in order of increasing pH at the equivalence point. c. Which titration requires the largest volume of titrant (HCl or NaOH) to reach the equivalence point?

Another way to treat data from a pH titration is to graph the absolute value of the change in pH per change in milliliters added versus milliliters added (DpHyDmL versus mL added). Make this graph using your results from Exercise 61. What advantage might this method have over the traditional method for treating titration data?

A buffer is made using 45.0 mL of 0.750 M HC3H5O2 (Ka 5 1.3 3 1025 ) and 55.0 mL of 0.700 M NaC3H5O2. What volume of 0.10 M NaOH must be added to change the pH of the original buffer solution by 2.5%?

A 0.400-M solution of ammonia was titrated with hydrochloric acid to the equivalence point, where the total volume was 1.50 times the original volume. At what pH does the equivalence point occur?

What volume of 0.0100 M NaOH must be added to 1.00 L of 0.0500 M HOCl to achieve a pH of 8.00?25.0 mL of 0.100 M Ba(OH)2, and 10.0 mL of 0.150 M KOH. Calculate the pH of this solution

Consider a solution formed by mixing 50.0 mL of 0.100 M H2SO4, 30.0 mL of 0.100 M HOCl, 25.0 mL of 0.200 M NaOH,

When a diprotic acid, H2A, is titrated with NaOH, the protons on the diprotic acid are generally removed one at a time, resulting in a pH curve that has the following generic shape: pH Vol NaOH added a. Notice that the plot has essentially two titration curves. If the first equivalence point occurs at 100.0 mL NaOH added, what volume of NaOH added corresponds to the second equivalence point? b. For the following volumes of NaOH added, list the major species present after the OH2 reacts completely. i. 0 mL NaOH added ii. between 0 and 100.0 mL NaOH added iii. 100.0 mL NaOH added iv. between 100.0 and 200.0 mL NaOH added v. 200.0 mL NaOH added vi. after 200.0 mL NaOH added c. If the pH at 50.0 mL NaOH added is 4.0 and the pH at 150.0 mL NaOH added is 8.0, determine the values Ka1 and Ka2 for the diprotic acid.

Consider the following two acids: Salicylic acid CO2H OH pKa1 5 2.98; pKa2 5 13.40 HO2CCH2CH2CH2CH2CO2H Adipic acid pKa1 5 4.41; pKa2 5 5.28 In two separate experiments the pH was measured during the titration of 5.00 mmol of each acid with 0.200 M NaOH. Each experiment showed only one stoichiometric point when the data were plotted. In one experiment the stoichiometric point was at 25.00 mL added NaOH, and in the other experiment the stoichiometric point was at 50.00 mL NaOH. Explain these results. (See Exercise 113.)

The titration of Na2CO3 with HCl has the following qualitative profile: mL HCl pH A B C D E F V a. Identify the major species in solution at points AF. b. Calculate the pH at the halfway points to equivalence, B and D. (Hint: Refer to Exercise 113.)

Consider the titration curve in Exercise 115 for the titration of Na2CO3 with HCl. a. If a mixture of NaHCO3 and Na2CO3 was titrated, what would be the relative sizes of V1 and V2? b. If a mixture of NaOH and Na2CO3 was titrated, what would be the relative sizes of V1 and V2? c. A sample contains a mixture of NaHCO3 and Na2CO3. When this sample was titrated with 0.100 M HCl, it took 18.9 mL to reach the first stoichiometric point and an additional 36.7 mL to reach the second stoichiometric point. What is the composition in mass percent of the sample?

A few drops of each of the indicators shown in the accompanying table were placed in separate portions of a 1.0-M solution of a weak acid, HX. The results are shown in the last column of the table. What is the approximate pH of the solution containing HX? Calculate the approximate value of Ka for HX. Indicator Color of HIn Color of In2 pKa of HIn Color of 1.0 M HX Bromphenol blue Yellow Blue 4.0 Blue Bromcresol purple Yellow Purple 6.0 Yellow Bromcresol green Yellow Blue 4.8 Green Alizarin Yellow Red 6.5 Yellow

Malonic acid (HO2CCH2CO2H) is a diprotic acid. In the titration of malonic acid with NaOH, stoichiometric points occur at pH 5 3.9 and 8.8. A 25.00-mL sample of malonic acid of unknown concentration is titrated with 0.0984 M NaOH, requiring 31.50 mL of the NaOH solution to reach the phenolphthalein end point. Calculate the concentration of the initial malonic acid solution. (See Exercise 113.)

A buffer solution is prepared by mixing 75.0 mL of 0.275 M fluorobenzoic acid (C7H5O2F) with 55.0 mL of 0.472 M sodium fluorobenzoate. The pKa of this weak acid is 2.90. What is the pH of the buffer solution?

A 10.00-g sample of the ionic compound NaA, where A2 is the anion of a weak acid, was dissolved in enough water to make 100.0 mL of solution and was then titrated with 0.100 M HCl. After 500.0 mL HCl was added, the pH was 5.00. The experimenter found that 1.00 L of 0.100 M HCl was required to reach the stoichiometric point of the titration. a. What is the molar mass of NaA? b. Calculate the pH of the solution at the stoichiometric point of the titration.

Calculate the pH of a solution prepared by mixing 250. mL of 0.174 m aqueous HF (density 5 1.10 g/mL) with 38.7 g of an aqueous solution that is 1.50% NaOH by mass (density 5 1.02 g/mL). (Ka for HF 5 7.2 3 1024 .)

Consider a solution prepared by mixing the following: 50.0 mL of 0.100 M Na3PO4 100.0 mL of 0.0500 M KOH 200.0 mL of 0.0750 M HCl 50.0 mL of 0.150 M NaCN Determine the volume of 0.100 M HNO3 that must be added to this mixture to achieve a final pH value of 7.21.