Consider the titration of 150.0 mL of 0.100 M HI by 0.250

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

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.

What if all you know about two salts is that the value of Ksp for salt A is greater than that of salt B? Why can we not compare relative solubilities of the salts? Use numbers to show how salt A could be more soluble than salt B, and how salt B can be more soluble than salt A.

You and a friend are studying for a chemistry exam. What if your friend tells you that since acids are very reactive, all salts are more soluble in aqueous solutions of acids than in water? How would you explain to your friend that this is not true? Use a specific example to defend your answer.

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? Explain.

A friend asks the following: Consider a buffered solution made up of the weak acid HA and its salt NaA. If a strong base such as NaOH is added, the HA reacts with the OH2 to make 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 your 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? Would a buffered solution made by mixing HCl and NaOH be effective? Explain

Sketch two pH curves, one for the titration of a weak acid with a strong base, and one for the titration of 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 calculate 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

Devise as many ways as you can to experimentally determine the Ksp value of a solid. Explain why each of these would work.

You are browsing through the Handbook of Hypothetical Chemistry when you come across a solid that is reported to have a Ksp value of zero in water at 25C. What does this mean?

A friend tells you: The constant Ksp of a salt is called the solubility product constant and is calculated from the concentrations of ions in the solution. Thus, if salt A dissolves to a greater extent than salt B, salt A must have a higher Ksp than salt B. Do you agree with your friend? Explain.

What happens to the Ksp value of a solid as the temperature of the solution changes? Consider both increasing and decreasing temperatures, and explain your answer.

Which is more likely to dissolve in an acidic solution, silver sulfide or silver chloride? Why?

You have two salts AgX and AgY with very similar Ksp values. You know that the Ka value for HX is much greater than the Ka value for HY. Which salt is more soluble in an acidic solution? Explain

Under what circumstances can the relative solubilities of two salts be compared by directly comparing values of their solubility products?

Define a buffered solution. What makes up a buffered solution? Explain how buffers absorb added H1 or OH2 with little pH change. 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 good buffer generally contains relatively equal concentrations of a weak acid and its 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 acidconjugate 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 acetat

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. H+ A B

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

Consider a buffered 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?

Derive an equation analogous to the Henderson Hasselbalch equation that relates pOH and pKb of a buffered solution composed of a weak base and its conjugate acid, such as NH3 and NH4 1.

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. 0.100 M HC3H5O2 and 0.100 M NaC3H5O2

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 NaOH is added to 1.00 L of each of the four solutions in Exercise 21.

The results of Exercises 2123 illustrate an important property of buffered solutions. Which solution in Exercise 21 is the buffered solution and what important property is illustrated by the results?

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. 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

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

Calculate the pH of a solution that is 0.100 M HONH2 and 0.100 M HONH3Cl.

Calculate the pH after 0.10 mole of NaOH is added to 1.00 L of the solution in Exercise 27, and calculate the pH after 0.20 mole of HCl is added to 1.00 L of the solution in Exercise 27

Calculate the pH after 0.020 mole of NaOH is added to 1.00 L of the solution in Exercise 28, and calculate the pH after 0.020 mole of HCl is added to 1.00 L of the solution in Exercise 28

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 reageant 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?

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

Consider a solution that contains both C5H5N and C5H5NHNO3. Calculate the ratio [C5H5N]/[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

How many 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

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?

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 buffered solution

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 buffered solution. b. Calculate the pH after 4.0 g of NaOH(s) is added to 1.0 L of this solution. (Neglect any volume change.

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?

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(aq) 34 HPO4 22(aq) 1 H1(aq) 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(aq) 34 H2PO4 2(aq) 1 H1(aq) Ka 5 7.5 3 1023

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(aq) 34 HCO3 2(aq) 1 H1(aq) 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 44. Calculate what happens to the [H2CO3]/[HCO3 2] ratio in blood when the pH decreases from 7.40 to 7.35.

Which of the following mixtures would result in a buffered solution when 1.0 L of each of the two solutions are mixed? a. 0.1 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

Calculate the pH of a solution formed by mixing 100.0 mL of 0.100 M NaF and 100.0 mL of 0.025 M HCl.

Consider the acids in Table 7.2. 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 7.3. 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.

A solution contains 1.0 3 1026 M HOCl and an unknown concentration of KOCl. If the pH of the solution is 7.20, calculate the KOCl concentration. (Hint: The contribution of water to the [H1] cannot be ignored.)

In Section 8.3 an equation was derived for the exact treatment of HA/NaA-type buffers. What would be the expression for B/BHCl-type buffers stated in terms of Kb, [OH2], [B], and [BH1]? Would it be necessary to use this exact expression to solve for the pH of a solution containing 1.0 3 1024 M HONH2 and 1.0 3 1025 M HONH3Cl? Explain

Consider a weak acid HA with a Ka value of 1.6 3 1027. Calculate the pH of a solution that is 5.0 3 1027 M HA and 5.0 3 1027 M NaA

Consider the following pH curves for 100.0 mL of two different acids with the same intital concentration each titrated by 0.10 M NaOH: Vol NaOH pH 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 titration of a generic weak acid HA with a strong base that gives the following titration curve: 5 pH 10 15 20 25 mL base On the curve indicate the points that correspond to the following. a. the equivalence point b. the maximum buffering region 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 H1 34 BHOn 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 [B] e. pH depends only on [BH1] f. pH depends only on the amount of excess strong acid added

Draw the general titration curve for a strong acid titrated with 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 acid strong 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? Answer the same questions for a strong base strong acid titration. Compare and contrast a strong acidstrong base titration with a strong basestrong acid titration.

Consider the following four titrations: i. 100.0 mL of 0.10 M HCl titrated with 0.10 M NaOH ii. 100.0 mL of 0.10 M NaOH titrated with 0.10 M HCl iii. 100.0 mL of 0.10 M CH3NH2 titrated with 0.10 M HCl iv. 100.0 mL of 0.10 M HF titrated with 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?

A student titrates an unknown weak acid HA to a palepink phenolphthalein endpoint 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 4.7. How is the value of pKa for the unknown acid related to 4.7? 61. The following plot shows the pH curves for the titrations of various acids with 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.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

The figure in the preceding exercise shows the pH curves for the titrations of six different acids with NaOH. Make a similar plot for the titration of three different bases with 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

Consider the titration of 40.0 mL of 0.200 M HClO4 with 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 with 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) with 0.100 M KOH. Calculate the pH of the resulting solution after each of the following volumes of KOH has 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) with 0.200 M HNO3. Calculate the pH of the resulting solution after each of the following volumes of HNO3 has 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 67 for the titration of 25.0 mL of 0.100 M propanoic acid (HC3H5O2, Ka 5 1.3 3 1025) with 0.100 M KOH.

Repeat the procedure in Exercise 67 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 67 for the titration of 25.0 mL of 0.100 M pyridine (Kb 5 1.7 3 1029) with 0.100 M hydrochloric acid. Do not do the points at 24.9 mL 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 with 0.10 M NaOH b. 100.0 mL of 0.10 M C2H5NH2 (Kb 5 5.6 3 1024) titrated with 0.20 M HNO3 c. 100.0 mL of 0.50 M HCl titrated with 0.25 M NaOH

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 mole 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.

What is an acidbase indicator? Define the equivalence (stoichiometric) point and the endpoint 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 CO2H group and no other acidic or basic functional group? Explain.

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 of NaOH has been added?

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?

Estimate the pH of a solution in which bromcresol green is blue and thymol blue is yellow (see Fig. 8.8)

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. 8.8.) a. thymol blue c. methyl red b. bromthymol blue d. crystal violet

Which of the indicators in Fig. 8.8 could be used for doing the titrations in Exercises 63 and 65?

Which of the indicators in Fig. 8.8 could be used for doing the titrations in Exercises 64 and 66?

Which of the indicators in Fig. 8.8 could be used for doing the titrations in Exercises 67 and 69

Which of the indicators in Fig. 8.8 could be used for doing the titrations in Exercises 68 and 70?

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?

Indicators can be used to estimate the pH values of solutions. To determine the pH of a 0.01 M weak acid (HX) solution, a few drops of three different indicators are added to separate portions of 0.01 M HX. The resulting colors of the HX solution are summarized in the last column of the accompanying table. What is the approximate pH of the 0.01 M HX solution? What is the approximate Ka value for HX? Indicator (HIn) Color of HIn Color of In2 pKa of HIn Color of 0.01 M HX Bromphenol Yellow Blue 4.0 Blue blue Bromcresol Yellow Purple 6.0 Yellow purple Bromcresol Yellow Blue 4.8 Green green

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

A student was given a 0.10 M solution of an unknown diprotic acid H2A and asked to determine the Ka1 and Ka2 values for the diprotic acid. The student titrated 50.0 mL of the 0.10 M H2A with 0.10 M NaOH. After 25.0 mL of NaOH was added, the pH of the resulting solution was 6.70. After 50.0 mL of NaOH was added, the pH of the resulting solution was 8.00. What are the values of Ka1 and Ka2 for the diprotic acid?

Consider the titration of 100.0 mL of a 0.0500 M solution of the hypothetical weak acid H3X (Ka1 5 1.0 3 1023, Ka2 5 1.0 3 1027, Ka3 5 1.0 3 10212) with 0.100 M KOH. Calculate the pH of the solution under the following conditions. a. before any KOH has been added b. after 10.0 mL of 0.100 M KOH has been added c. after 25.0 mL of 0.100 M KOH has been added d. after 50.0 mL of 0.100 M KOH has been added e. after 60.0 mL of 0.100 M KOH has been added f. after 75.0 mL of 0.100 M KOH has been added g. after 100.0 mL of 0.100 M KOH has been added h. after 125.0 mL of 0.100 M KOH has been added i. after 150.0 mL of 0.100 M KOH has been added j. after 200.0 mL of 0.100 M KOH has been added

Consider 100.0 mL of a 0.100 M solution of H3A (Ka1 5 1.5 3 1024, Ka2 5 3.0 3 1028, Ka3 5 5.0 3 10212). a. Calculate the pH of this solution. b. Calculate the pH of the solution after 10.0 mL of 1.00 M NaOH has been added to the original solution. c. Calculate the pH of the solution after 25.0 mL of 1.00 M NaOH has been added to the original solution.

A 0.200-g sample of a triprotic acid (molar mass 5 165.0 g/mol) is dissolved in a 50.00-mL aqueous solution and titrated with 0.0500 M NaOH. After 10.50 mL of the base was added, the pH was observed to be 3.73. The pH at the first stoichiometric point was 5.19 and at the second stoichiometric point was 8.00. a. Calculate the three Ka values for the acid. b. Make a reasonable estimate of the pH after 59.0 mL of 0.0500 M NaOH has been added. Explain your answer. c. Calculate the pH after 59.0 mL of 0.0500 M NaOH has been added.

Consider the titration of 100.0 mL of 0.100 M H3A (Ka1 5 5.0 3 1024, Ka2 5 1.0 3 1028, Ka3 5 1.0 3 10211) with 0.0500 M NaOH. a. Calculate the pH after 100.0 mL of 0.0500 M NaOH has been added. b. What total volume of 0.0500 M NaOH is required to reach a pH of 8.67?

The titration of Na2CO3 with HCl has the following qualitative profile: V1 V2 mL HCl pH A B C D E F a. Identify the major species in solution as points AF. b. For the titration of 25.00 mL of 0.100 M Na2CO3 with 0.100 M HCl, calculate the pH at points AE. (B and D are halfway points to equivalence.)

Consider 100.0 mL of a solution of 0.200 M Na2A, where A22 is a base with corresponding acids H2A (Ka 5 1.0 3 1023) and HA2 (Ka 5 1.0 3 1028). a. What volume of 1.00 M HCl must be added to this solution to reach pH 5 8.00? b. Calculate the pH at the second stoichiometric point of the titration of 0.200 M Na2A, with 1.00 M HCl.

For which of the following is the Ksp value of the ionic compound the largest? The smallest? Explain your answer

Ag2S(s) has a larger molar solubility than CuS even though Ag2S has the smaller Ksp value. Explain how this is possible.

When Na3PO4(aq) is added to a solution containing a metal ion and a precipitate forms, the precipitate generally could be one of two possibilities. What are the two possibilities?

The common ion effect for ionic solids (salts) is to significantly decrease the solubility of the ionic compound in water. Explain the common ion effect.

Calculate the solubility of each of the following compounds in moles per liter and grams per liter. (Ignore any acidbase properties.) a. Ag3PO4, Ksp 5 1.8 3 10218 b. CaCO3, Ksp 5 8.7 3 1029 c. Hg2Cl2, Ksp 5 1.1 3 10218 (Hg2 21 is the cation in solution.)

Calculate the solubility of each of the following compounds in moles per liter. Ignore any acidbase properties. a. PbI2, Ksp 5 1.4 3 1028 b. CdCO3, Ksp 5 5.2 3 10212 c. Sr3(PO4)2, Ksp 5 1 3 10231

Use the following data to calculate the Ksp value for each solid. a. The solubility of CaC2O4 is 6.1 3 1023 g/L. b. The solubility of BiI3 is 1.32 3 1025 mol/L.

The concentration of Pb21 in a solution saturated with PbBr2(s) is 2.14 3 1022 M. Calculate Ksp for PbBr2

The concentration of Ag1 in a solution saturated with Ag2C2O4(s) is 2.2 3 1024 M. Calculate Ksp for Ag2C2O4.

The solubility of the ionic compound M2X3, having a molar mass of 288 g/mol, is 3.60 3 1027 g/L. Calculate the Ksp of the compound

For each of the following pairs of solids, determine which solid has the smallest molar solubility. a. CaF2(s), Ksp 5 4.0 3 10211 or BaF2(s), Ksp 5 2.4 3 1025 b. Ca3(PO4)2(s), Ksp 5 1.3 3 10232 or FePO4(s), Ksp 5 1.0 3 10222

The solubility rules outlined in Chapter 4 say that Ba(OH)2, Sr(OH)2, and Ca(OH)2 are marginally soluble hydroxides. Calculate the pH of a saturated solution of each of these marginally soluble hydroxides

Calculate the molar solubility of Co(OH)3, Ksp 5 2.5 3 10243.

The Ksp for silver sulfate (Ag2SO4) is 1.2 3 1025. Calculate the solubility of silver sulfate in each of the following. a. water b. 0.10 M AgNO3 c. 0.20 M K2SO4

Calculate the solubility (in mol/L) of Fe(OH)3 (Ksp 5 4 3 10238) in each of the following. a. water (assume pH is 7.0 and constant) b. a solution buffered at pH 5 5.0 c. a solution buffered at pH 5 11.0

The Ksp for lead iodide (PbI2) is 1.4 3 1028. Calculate the solubility of lead iodide in each of the following. a. water b. 0.10 M Pb(NO3)2 c. 0.010 M NaI

Calculate the solubility of solid Ca3(PO4)2 (Ksp 5 1.3 3 10232) in a 0.20-M Na3PO4 solution.

The solubility of Ce(IO3)3 in a 0.20 M KIO3 solution is 4.4 3 1028 mol/L. Calculate Ksp for Ce(IO3)3.

What mass of ZnS (Ksp 5 2.5 3 10222) will dissolve in 300.0 mL of 0.050 M Zn(NO3)2? Ignore the basic properties of S

The concentration of Mg21 in seawater is 0.052 M. At what pH will 99% of the Mg21 be precipitated as the hydroxide salt? [Ksp for Mg(OH)2 5 8.9 3 10212.

For the substances in Exercises 97 and 98, which will show increased solubility as the pH of the solution becomes more acidic? Write equations for the reactions that occur to increase the solubility.

Explain the following phenomenon: You have a test tube with an aqueous solution of silver nitrate as shown in test tube 1 below. A few drops of aqueous sodium chromate solution was added with the end result shown in test tube 2. A few drops of aqueous sodium chloride solution was then added with the end result shown in test tube 3. 1 CrO4 2 2 3 Cl Use the Ksp values in the book to support your explanation, and include the balanced equations. Also, list the ions that are present in solution in each test tube

For which salt in each of the following groups will the solubility depend on pH? a. AgF, AgCl, AgBr c. Sr(NO3)2, Sr(NO2)2 b. Pb(OH)2, PbCl2 d. Ni(NO3)2, Ni(CN)2

A solution is prepared by mixing 75.0 mL of 0.020 M BaCl2 and 125 mL of 0.040 M K2SO4. What are the concentrations of barium and sulfate ions in this solution? Assume only SO4 22 ions (no HSO4 2) are present

Calculate the final concentrations of K1(aq), C2O4 22(aq), Ba21(aq), and Br2(aq) in a solution prepared by adding 0.100 L of 6.0 3 1024 M K2C2O4 to 0.150 L of 1.0 3 1024 M BaBr2. (For BaC2O4, Ksp 5 2.3 3 1028.)

A solution is prepared by mixing 50.0 mL of 0.10 M Pb(NO3)2 with 50.0 mL of 1.0 M KCl. Calculate the concentrations of Pb21 and Cl2 at equilibrium. [Ksp for PbCl2(s) 5 1.6 3 1025.]

The Ksp of Al(OH)3 is 2 3 10232. At what pH will a 0.2 M Al31 solution begin to show precipitation of Al(OH)3?

A solution is 1 3 1024 M in NaF, Na2S, and Na3PO4. What would be the order of precipitation as a source of Pb21 is added gradually to the solution? The relevant Ksp values are Ksp(PbF2) 5 4 3 1028, Ksp(PbS) 5 7 3 10229, and Ksp[Pb3(PO4)2] 5 1 3 10254.

A solution contains 1.0 3 1025 M Na3PO4. What is the minimum concentration of AgNO3 that would cause precipitation of solid Ag3PO4 (Ksp 5 1.8 3 10218)?

A solution contains 0.25 M Ni(NO3)2 and 0.25 M Cu(NO3)2. Can the metal ions be separated by slowly adding Na2CO3? Assume that for successful separation, 99% of the metal ion must be precipitated before the other metal ion begins to precipitate, and assume that no volume change occurs on addition of Na2CO3.

Describe how you could separate the ions in each of the following groups by selective precipitation. a. Ag1, Mg21, Cu21 c. Cl2, Br2, I2 b. Pb21, Ca21, Fe21 d. Pb21, Bi31

If a solution contains either Pb21(aq) or Ag1(aq), how can temperature be manipulated to help identify the ion in solution?

Sulfide precipitates are generally grouped as sulfides insoluble in acidic solution and sulfides insoluble in basic solution. Explain why there is a difference between the two groups of sulfide precipitates.

Nanotechnology has become an important field, with applications ranging from high-density data storage to the design of nano machines. One common building block of nanostructured architectures is manganese oxide nanoparticles. The particles can be formed from manganese oxalate nanorods, the formation of which can be described as follows: Mn21(aq) 1 C2O4 22(aq) 34 MnC2O4(aq) K1 5 7.9 3 103 MnC2O4(aq) 1 C2O4 22(aq) 34 Mn(C2O4)2 22(aq) K2 5 7.9 3 101 Calculate the value for the overall formation constant for Mn(C2O4)2 22: K 5 3Mn1C2O42 2 22 4 3Mn21 4 3C2O4 22 4 2

When aqueous KI is added gradually to mercury(II) nitrate, an orange precipitate forms. Continued addition of KI causes the precipitate to dissolve. Write balanced equations to explain these observations. (Hint: Hg21 reacts with I2 to form HgI4 22.)

As a sodium chloride solution is added to a solution of silver nitrate, a white precipitate forms. Ammonia is added to the mixture and the precipitate dissolves. When potassium bromide solution is then added, a pale yellow precipitate appears. When a solution of sodium thiosulfate is added, the yellow precipitate dissolves. Finally, potassium iodide is added to the solution and a yellow precipitate forms. Write reactions for all the changes mentioned above. What conclusions can you draw concerning the sizes of the Ksp values for AgCl, AgBr, and AgI? What can you say about the relative values of the formation constants of Ag(NH3)2 1 and Ag(S2O3)2 32?

The overall formation constant for HgI4 22 is 1.0 3 1030. That is, 1.0 3 1030 5 3HgI4 22 4 3Hg21 4 3I 2 4 4 What is the concentration of Hg21 in 500.0 mL of a solution that was originally 0.010 M Hg21 and had 65 g of KI added to it? The reaction is Hg21(aq) 1 4I2(aq) 34 HgI4 22(aq)

A solution is prepared by adding 0.090 mole of K3[Fe(CN)6] to 0.60 L of 2.0 M NaCN. Assuming no volume change, calculate the concentrations of Fe(CN)6 32 and Fe31 in this solution. The K (overall) for the formation of Fe(CN)6 32 is 1 3 1042

A solution is prepared by mixing 100.0 mL of 1.0 3 1024 M Be(NO3)2 and 100.0 mL of 8.0 M NaF. Be21(aq) 1 F2(aq) 34 BeF1(aq) K1 5 7.9 3 104 BeF1(aq) 1 F2(aq) 34 BeF2(aq) K2 5 5.8 3 103 BeF2(aq) 1 F2(aq) 34 BeF3 2(aq) K3 5 6.1 3 102 BeF3 2(aq) 1 F2(aq) 34 BeF4 22(aq) K4 5 2.7 3 101 Calculate the equilibrum concentrations of F2, Be21, BeF1, BeF2, BeF3 2, and BeF4 22 in this solution.

Kf for the complex ion Ag(NH3)2 1 is 1.7 3 107. Ksp for AgCl is 1.6 3 10210. Calculate the molar solubility of AgCl in 1.0 M NH3.

a. Using the Ksp for Cu(OH)2 (1.6 3 10219) and the overall formation constant for Cu(NH3)4 21 (1.0 3 1013), calculate a value for the equilibrium constant for the reaction Cu(OH)2(s) 1 4NH3(aq) 34 Cu(NH3)4 21(aq) 1 2OH2(aq) b. Use the value of the equilibrium constant you calculated in part a to calculate the solubility (in mol/L) of Cu(OH)2 in 5.0 M NH3. In 5.0 M NH3, the concentration of OH2 is 0.0095 M.

The copper(I) ion forms a chloride salt that has Ksp 5 1.2 3 1026. Copper(I) also forms a complex ion with Cl2: Cu1(aq) 1 2Cl2(aq) 34 CuCl2 2(aq) K 5 8.7 3 104 a. Calculate the solubility of copper(I) chloride in pure water. (Ignore CuCl2 2 formation for part a.) b. Calculate the solubility of copper(I) chloride in 0.10 M NaCl

Solutions of sodium thiosulfate are used to dissolve unexposed AgBr in the developing process for black-andwhite film. What mass of AgBr can dissolve in 1.00 L of 0.500 M Na2S2O3? Assume the overall formation constant for Ag(S2O3)2 32 is 2.9 3 1013 and Ksp for AgBr is 5.0 3 10213.

a. Calculate the molar solubility of AgI in pure water. Ksp for AgI is 1.5 3 10216. b. Calculate the molar solubility of AgI in 3.0 M NH3. The overall formation constant for Ag(NH3)2 1 is 1.7 3 107. c. Compare the calculated solubilities from parts a and b. Explain any differences.

A series of chemicals was added to some AgNO3(aq). NaCl(aq) was added first to the silver nitrate solution, with the end result shown below in test tube 1; NH3(aq) was then added, with the end result shown in test tube 2; and HNO3(aq) was added last, with the end result shown in test tube 3. 1 2 3 Explain the results shown in each test tube. Include a balanced equation for the reaction(s) taking place.

Will a precipitate of Cd(OH)2 form if 1.0 mL of 1.0 M Cd(NO3)2 is added to 1.0 L of 5.0 M NH3? Cd21(aq) 1 4NH3(aq) 34 Cd(NH3)4 21(aq) K 5 1.0 3 107 Cd(OH)2(s) 34 Cd21(aq) 1 2OH2(aq) Ksp 5 5.9 3 10215c. A buffer is prepared by diluting 50.0 g of TRIS base and 65.0 g of 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

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?

The solubility of copper(II) hydroxide in water can be increased by adding either the base NH3 or the acid HNO3. Explain. Would added NH3 or HNO3 have the same effect on the solubility of silver acetate or silver chloride? Explain

The salts in Table 8.5, with the possible exception of the hydroxide salts, have one of the following mathematical relationships between the Ksp value and the molar solubility s. i. Ksp 5 s2 iii. Ksp 5 27s4 ii. Ksp 5 4s3 iv. Ksp 5 108s5 For each mathematical relationship, give an example of a salt in Table 8.5 that exhibits that relationship.

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

Repeat the procedure in Exercise 67, but for the titration of 25.0 mL of 0.100 M HNO3 with 0.100 M NaOH.

One method for determining the purity of aspirin (empirical formula 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(s) 1 2OH2(aq) Aspirin 888888n C7H5O3 2(aq) 1 C2H3O2 2(aq) 1 H2O(l) 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?

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 (DpH/DmL versus mL added). Make this graph using your results from Exercise 67. What advantage might this method have over the traditional method for treating titration data?

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(aq) 1 OH2(aq) 88n H2O(l) 1 P22(aq) If a typical titration experiment begins with approximately 0.5 g of 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 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.

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.

What mass of Ca(NO3)2 must be added to 1.0 L of a 1.0 M HF solution to begin precipitation of CaF2(s)? For CaF2, Ksp 5 4.0 3 10211 and Ka for HF 5 7.2 3 1024. Assume no volume change on addition of Ca(NO3)2(s).

The equilibrium constant for the following reaction is 1.0 3 1023: Ethylenediaminetetraacetate O2C C O H2 N G G G OCH2OCH2ON G O2C C O H2 CH2 CO2 O CH2 CO2 O EDTA4 = Cr3+(aq) CrEDTA(aq) + 2H+ H2EDTA (aq) 2 + (aq)EDTA is used as a complexing agent in chemical analysis. Solutions of EDTA, usually containing the disodium salt Na2H2EDTA, are used to treat heavy metal poisoning. Calculate [Cr31] at equilibrium in a solution originally 0.0010 M in Cr31 and 0.050 M in H2EDTA22 and buffered at pH 5 6.00

Calculate the concentration of Pb21 in each of the following. a. a saturated solution of Pb(OH)2, sp 5 1.2 3 10215 b. a saturated solution of Pb(OH)2 buffered at pH 5 13.00 c. Ethylenediaminetetraacetate (EDTA42) is used as a complexing agent in chemical analysis and has the following structure: Ethylenediaminetetraacetate N CH2 CH2 N O2C CH2 CH2 O2C CO2 CO2 CH 2 CH2 Solutions of EDTA42 are used to treat heavy metal poisoning by removing the heavy metal in the form of a soluble complex ion. The reaction of EDTA42 with Pb21 is Pb21 1aq2 1 EDTA42 1aq2mPbEDTA22 1aq2 K 5 1.1 3 1018 Consider a solution with 0.010 mole of Pb(NO3)2 added to 1.0 L of an aqueous solution buffered at pH 5 13.00 and containing 0.050 M Na4EDTA. Does Pb(OH)2 precipitate from this solution?

Consider saturated solutions of the following compounds. a. Mg(OH)2 b. Cd(OH)2 c. Pb(OH)2 Calculate the pH of each saturated solution

A certain acetic acid solution has pH 5 2.68. Calculate the volume of 0.0975 M KOH required to neutralize 25.0 mL of this solution

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.

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?

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.

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 15.44 mL of 0.4524 M HCl to the flask containing the aspirin and the sodium hydroxide produced a mixture with pH 5 3.48. Find the molar mass of acetylsalicylic acid and its Ka value. Acetylsalicylic acid is a monoprotic acid.

A solution is formed by mixing 50.0 mL of 10.0 M NaX with 50.0 mL of 2.0 3 1023 M CuNO3. Assume that Cu(I) forms complex ions with X2 as follows: Cu1(aq) 1 X2(aq) 34 CuX(aq) K1 5 1.0 3 102 CuX(aq) 1 X2(aq) 34 CuX2 2(aq) K2 5 1.0 3 104 CuX2 2(aq) 1 X2(aq) 34 CuX3 22(aq) K3 5 1.0 3 103 Calculate the following concentrations at equilibrium. a. CuX3 22 b. CuX2 2 c. Cu1

When phosphoric acid is titrated with a NaOH solution, only two stoichiometric points are seen. Why?

Consider the following two acids: pKa1 5 2.98; pKa2 5 13.40 Salicylic acid OH CO2H 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 85.)

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. 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?

Assuming that the solubility of Ca3(PO4)2(s) is 1.6 3 1027 mol/L at 258C, calculate the Ksp for this salt. Ignore any potential reactions of the ions with water

Order the following solids (ad) from least soluble to most soluble. Ignore any potential reactions of the ions with water. a. AgCl Ksp 5 1.6 3 10210 b. Ag2S Ksp 5 1.6 3 10249 c. CaF2 Ksp 5 4.0 3 10211 d. CuS Ksp 5 8.5 3 10245

The Ksp for PbI2(s) is 1.4 3 1028. Calculate the solubility of PbI2(s) in 0.048 M NaI.

The solubility of Pb(IO3)2(s) in a 7.2 3 1022-M KIO3 solution is 6.0 3 1029 mol/L. Calculate the Ksp value for Pb(IO3)2(s)

A 50.0-mL sample of 0.0413 M AgNO3(aq) is added to 50.0 mL of 0.100 M NaIO3(aq). Calculate the [Ag1] at equilibrium in the resulting solution. [Ksp for AgIO3(s) 5 3.17 3 1028.]

The Hg21 ion forms complex ions with I2 as follows: Hg21 1aq2 1 I 2 1aq2mHgI1 1aq2 K1 5 1.0 3 108 HgI1 1aq2 1 I 2 1aq2mHgI2 1aq2 K2 5 1.0 3 105 HgI2 1aq2 1 I 2 1aq2mHgI3 2 1aq2 K3 5 1.0 3 109 HgI3 2 1aq2 1 I 2 1aq2mHgI4 22 1aq2 K4 5 1.0 3 108 A solution is prepared by dissolving 0.088 mole of Hg(NO3)2 and 5.00 mole of NaI in enough water to make 1.0 L of solution. a. Calculate the equilibrium concentration of [HgI4 22]. b. Calculate the equilibrium concentration of [I2]. c. Calculate the equilibrium concentration of [Hg21]

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%?

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?

For solutions containing salts of the form NH4X, the pH is determined by using the equation pH 5 pKa 1NH4 12 1 pKa 1HX2 2 a. Derive this equation. (Hint: Review Section 8.7 on the pH of solutions containing amphoteric species.) b. Use this equation to calculate the pH of the following solutions: ammonium formate, ammonium acetate, and ammonium bicarbonate. See Appendix 5 for Ka values. c. Solutions of ammonium acetate are commonly used as pH 5 7.0 buffers. Write equations to show how an ammonium acetate solution neutralizes added H1 and OH2.

Consider the titration of 100.0 mL of a solution that contains a mixture of 0.050 M H2SO4 and 0.20 M H2C6H6O6. Calculate the pH a. before any 0.10 M NaOH has been added. b. after a total of 100.0 mL of 0.10 M NaOH has been added. c. after a total of 300.0 mL of 0.10 M NaOH has been added. d. after a total of 500.0 mL of 0.10 M NaOH has been added.

The copper(I) ion forms a complex ion with CN2 according to the following equation: Cu1(aq) 1 3CN(aq)2 34 Cu(CN)3 22(aq) Kf 5 1.0 3 1011 a. Calculate the solubility of CuBr(s) (Ksp 5 1.0 3 1025) in 1.0 L of 1.0 M NaCN. b. Calculate the concentration of Br2 at equilibrium. c. Calculate the concentration of CN2 at equilibrium.

Calcium oxalate (CaC2O4) is relatively insoluble in water (Ksp 5 2 3 1029). However, calcium oxalate is more soluble in acidic solution. How much more soluble is calcium oxalate in 0.10 M H1 than in pure water? In pure water, ignore the basic properties of C2O4 22.

a. Calculate the molar solubility of SrF2 in water, ignoring the basic properties of F2. (For SrF2, Ksp 5 7.9 3 10210.) b. Would the measured molar solubility of SrF2 be greater than or less than the value calculated in part a? Explain. c. Calculate the molar solubility of SrF2 in a solution buffered at pH 5 2.00. (Ka for HF is 7.2 3 1024.)

What is the maximum possible concentration of Ni21 ion in water at 258C that is saturated with 0.10 M H2S and maintained at pH 3.0 with HCl?

A mixture contains 1.0 3 1023 M Cu21 and 1.0 3 1023 M Mn21 and is saturated with 0.10 M H2S. Determine a pH where CuS precipitates but MnS does not precipitate. Ksp for CuS 5 8.5 3 10245 and Ksp for MnS 5 2.3 3 10213.

Consider 1.0 L of an aqueous solution that contains 0.10 M sulfuric acid to which 0.30 mole barium nitrate is added. Assuming no change in volume of the solution, determine the pH, the concentration of barium ions in the final solution, and the mass of solid formed.

Calculate the solubility of AgCN(s) (Ksp 5 2.2 3 10212) in a solution containing 1.0 M H1. (Ka for HCN is 6.2 3 10210.)

Consider the titration of 100.0 mL of a 1.00 3 1024 M solution of an acid HA (Ka 5 5.0 3 10210) with 1.00 3 1023 M NaOH. Calculate the pH for the following conditions. a. before any NaOH has been added b. after 5.00 mL of NaOH has been added c. at the stoichiometric point

Consider a solution formed by mixing 200.0 mL of 0.250 M Na3PO4, 135.0 mL of 1.000 M HCl, and 100.0 mL of 0.100 M NaCN. a. Calculate the pH of this solution. b. Calculate the concentration of HCN in 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, 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

Calculate the pH of a solution prepared by mixing 500.0 mL of 0.50 M Na3PO4 and 500.0 mL of 0.10 M H2SO4.

Consider the titration of 100.0 mL of 0.10 M phosphoric acid with 0.10 M NaOH. a. Determine the pH at the third half-equivalence point by assuming it is a special point (see Fig. 8.11). b. Calculate the pH at the third equivalence point. c. Why must the answer to part a be incorrect? Why cant we use the special point on the graph? (Explain the assumption made in using the special point and why it is not valid in this case.) d. Calculate the pH at the third half-equivalence point.

In the titration of 100.0 mL of a 0.0500 M solution of acid H3A (Ka1 5 1.0 3 1023, Ka2 5 5.0 3 1028, Ka3 5 2.0 3 10212), calculate the volume of 1.00 M NaOH required to reach pH values of 9.50 and 4.00.

Consider the titration curve in Exercise 91 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?

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?

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.

Aluminum ions react with the hydroxide ion to form the precipitate Al(OH)3(s), but can also react to form the soluble complex ion Al(OH)4 2. In terms of solubility, Al(OH)3(s) will be more soluble in very acidic solutions as well as more soluble in very basic solutions. a. Write equations for the reactions that occur to increase the solubility of Al(OH)3(s) in very acidic solutions and in very basic solutions. b. Lets study the pH dependence of the solubility of Al(OH)3(s) in more detail. Show that the solubility of Al(OH)3, as a function of [H1], obeys the equation S 5 3H1 4 3 Ksp/Kw 3 1 KKw/3H1 4 where S 5 solubility 5 [Al31] 1 [Al(OH)4 2] and K is the equilibrium constant for Al1OH2 3 1s2 1 OH2 1aq2mAl1OH2 4 2 1aq2 c. The value of K is 40.0 and Ksp for Al(OH)3 is 2 3 10232. Plot the solubility of Al(OH)3 in the pH range 412.