The solubility of KCl is 3.7 M at 20 C. Two beakers

What is the purpose of an acid-base buffer?

How do the acid and base components of a buffer function? Why are they often a conjugate acid- base pair of a weak acid?

What is the common-ion effect? How is it related to Le Chteliers principle? Explain with equations that include HF and NaF

The scenes below depict solutions of the same HA/A? buffer (with counterions and water molecules omitted for clarity). (a) Which solution has the greatest buffer capacity? (b) Explain how the pH ranges of the buffers compare. (c) Which solution can react with the largest amount of added strong acid?

What is the difference between buffers with high and low capacities? Will adding 0.01 mol of HCl produce a greater pH change in a buffer with a high or a low capacity? Explain.

Which of these factors influence buffer capacity? How? (a) Conjugate acid-base pair (b) pH of the buffer (c) Concentration of buffer components (d) Buffer range (e) pKa of the acid component

What is the relationship between the buffer range and the buffer-component concentration ratio?

A chemist needs a pH 3.5 buffer. Should she use NaOH with formic acid (Ka ? 1.8?104 ) or with acetic acid (Ka ? 1.8?105 )? Why? What is the disadvantage of choosing the other acid? What is the role of the NaOH?

State and explain the relative change in the pH and in the buffer-component concentration ratio, [NaA]/[HA], for each of the following additions: (a) Add 0.1 M NaOH to the buffer (b) Add 0.1 M HCl to the buffer (c) Dissolve pure NaA in the buffer (d) Dissolve pure HA in the buffer 19.10 Does the pH increase or decrease, and does it do so to a large or small extent, with each of the following additions? (a) 5 drops of 0.1 M NaOH to 100 mL of 0.5 M acetate buffer (b) 5 drops of 0.1 M HCl to 100 mL of 0.5 M acetate buffer (c) 5 drops of 0.1 M NaOH to 100 mL of 0.5 M HCl (d) 5 drops of 0.1 M NaOH to distilled water

What are the [H3O] and the pH of a propanoic acid propanoate buffer that consists of 0.35 M CH3CH2COONa and 0.15 M CH3CH2COOH (Ka of propanoic acid ? 1.3?105 )?

What are the [H3O] and the pH of a benzoic acidbenzoate buffer that consists of 0.33 M C6H5COOH and 0.28 M C6H5COONa (Ka of benzoic acid ? 6.3?105 )?

What are the [H3O] and the pH of a benzoic acidbenzoate buffer that consists of 0.33 M C6H5COOH and 0.28 M C6H5COONa (Ka of benzoic acid ? 6.3?105 )?

What are the [H3O] and the pH of a buffer that consists of 0.55 M HNO2 and 0.75 M KNO2 (Ka of HNO2 ? 7.1?104 )?

What are the [H3O] and the pH of a buffer that consists of 0.20 M HF and 0.25 M KF (Ka of HF ? 6.8?104 )?

Find the pH of a buffer that consists of 0.45 M HCOOH and 0.63 M HCOONa (pKa of HCOOH ? 3.74)

Find the pH of a buffer that consists of 0.95 M HBrO and 0.68 M KBrO (pKa of HBrO ? 8.64).

Find the pH of a buffer that consists of 1.3 M sodium phenolate (C6H5ONa) and 1.2 M phenol (C6H5OH) (pKa of phenol ? 10.00)

Find the pH of a buffer that consists of 0.12 M boric acid (H3BO3) and 0.82 M sodium borate (NaH2BO3) (pKa of boric acid ? 9.24)

Find the pH of a buffer that consists of 0.25 M NH3 and 0.15 M NH4Cl (pKb of NH3 ? 4.75)

Find the pH of a buffer that consists of 0.50 M methylamine (CH3NH2) and 0.60 M CH3NH3Cl (pKb of CH3NH2 ? 3.35)

A buffer consists of 0.22 M KHCO3 and 0.37 M K2CO3. Carbonic acid is a diprotic acid with Ka1 ? 4.5?107 and Ka2 ? 4.7?1011. (a) Which Ka value is more important to this buffer? (b) What is the buffer pH?

A buffer consists of 0.50 M NaH2PO4 and 0.40 M Na2HPO4. Phosphoric acid is a triprotic acid (Ka1 ? 7.2?103 , Ka2 ? 6.3?108 , and Ka3 ? 4.2?1013). (a) Which Ka value is most important to this buffer? (b) What is the buffer pH? 1

What is the component concentration ratio, [Pr]/[HPr], of a buffer that has a pH of 5.44 (Ka of HPr ? 1.3?105 )?

What is the component concentration ratio, [NO2 ]/[HNO2], of a buffer that has a pH of 2.95 (Ka of HNO2 ? 7.1?104 )?

What is the component concentration ratio, [BrO]/[HBrO], of a buffer that has a pH of 7.95 (Ka of HBrO ? 2.3?109 )?

What is the component concentration ratio, [CH3COO]/[CH3COOH], of a buffer that has a pH of 4.39 (Ka of CH3COOH ? 1.8?105 )?

A buffer containing 0.2000 M of acid, HA, and 0.1500 M of its conjugate base, A, has a pH of 3.35. What is the pH after 0.0015 mol of NaOH is added to 0.5000 L of this solution?

A buffer that contains 0.40 M base, B, and 0.25 M of its conjugate acid, BH, has a pH of 8.88. What is the pH after 0.0020 mol of HCl is added to 0.25 L of this solution?

A buffer that contains 0.110 M HY and 0.220 M Y has a pH of 8.77. What is the pH after 0.0015 mol of Ba(OH)2 is added to 0.350 L of this solution?

A buffer that contains 1.05 M B and 0.750 M BH has a pH of 9.50. What is the pH after 0.0050 mol of HCl is added to 0.500 L of this solution?

A buffer is prepared by mixing 204 mL of 0.452 M HCl and 0.500 L of 0.400 M sodium acetate. (See Appendix C.) (a) What is the pH? (b) How many grams of KOH must be added to 0.500 L of the buffer to change the pH by 0.15 units?

A buffer is prepared by mixing 50.0 mL of 0.050 M sodium bicarbonate and 10.7 mL of 0.10 M NaOH. (See Appendix C.) (a) What is the pH? (b) How many grams of HCl must be added to 25.0 mL of the buffer to change the pH by 0.07 units?

Choose specific acid-base conjugate pairs to make the following buffers: (a) pH 4.5; (b) pH 7.0. (See Appendix C.)

Choose specific acid-base conjugate pairs to make the following buffers: (a) [H3O] 1?109 M; (b) [OH] 3?105 M. (See Appendix C.) 19

Choose specific acid-base conjugate pairs to make the following buffers: (a) pH 3.5; (b) pH 5.5. (See Appendix C.)

Choose specific acid-base conjugate pairs to make the following buffers: (a) [OH] 1?106 M; (b) [H3O] 4?104 M. (See Appendix C.) Pr

An industrial chemist studying bleaching and sterilizing prepares several hypochlorite buffers. Find the pH of (a) 0.100 M HClO and 0.100 M NaClO; (b) 0.100 M HClO and 0.150 M NaClO; (c) 0.150 M HClO and 0.100 M NaClO; (d) 1.0 L of the solution in part (a) after 0.0050 mol of NaOH has been added

Oxoanions of phosphorus are buffer components in blood. For a KH2PO4 /Na2HPO4 solution with pH ? 7.40 (pH of normal arterial blood), what is the buffer-component concentration ratio?

How can you estimate the pH range of an indicators color change? Why do some indicators have two separate pH ranges?

Why does the color change of an indicator take place over a range of about 2 pH units?

Why doesnt the addition of an acid-base indicator affect the pH of the test solution?

What is the difference between the end point of a titration and the equivalence point? Is the equivalence point always reached first? Explain.

The scenes below depict the relative concentrations of H3PO4, H2PO4 , and HPO4 2 during a titration with aqueous NaOH, but they are out of order. (Phosphate groups are purple, hydrogens are blue, and Na ions and water molecules are not shown.) (a) List the scenes in the correct order. (b) What is the pH in the correctly ordered second scene (see Appendix C)? (c) If it requires 10.00 mL of the NaOH solution to reach this scene, how much more is needed to reach the last scene? 1

Explain how strong acidstrong base, weak acidstrong base, and weak basestrong acid titrations using the same concentrations differ in terms of (a) the initial pH and (b) the pH at the equivalence point. (The component in italics is in the flask.)

What species are in the buffer region of a weak acidstrong base titration? How are they different from the species at the equivalence point? How are they different from the species in the buffer region of a weak basestrong acid titration?

Why is the center of the buffer region of a weak acidstrong base titration significant?

How does the titration curve of a monoprotic acid differ from that of a diprotic acid?

The indicator cresol red has Ka ? 3.5?109 . Over what approximate pH range does it change color?

The indicator ethyl red has Ka ? 3.8?106 . Over what approximate pH range does it change color?

Use Figure 19.5 to find an indicator for these titrations: (a) 0.10 M HCl with 0.10 M NaOH (b) 0.10 M HCOOH (Appendix C) with 0.10 M NaOH

Use Figure 19.5 to find an indicator for these titrations: (a) 0.10 M CH3NH2 (Appendix C) with 0.10 M HCl (b) 0.50 M HI with 0.10 M KOH

Use Figure 19.5 to find an indicator for these titrations: (a) 0.5 M (CH3)2NH (Appendix C) with 0.5 M HBr (b) 0.2 M KOH with 0.2 M HNO3

Use Figure 19.5 to find an indicator for these titrations: (a) 0.25 M C6H5COOH (Appendix C) with 0.25 M KOH (b) 0.50 M NH4Cl (Appendix C) with 0.50 M NaOH

Calculate the pH during the titration of 40.00 mL of 0.1000 M HCl with 0.1000 M NaOH solution after the following additions of base: (a) 0 mL (b) 25.00 mL (c) 39.00 mL (d) 39.90 mL (e) 40.00 mL (f) 40.10 mL (g) 50.00 mL

Calculate the pH during the titration of 30.00 mL of 0.1000 M KOH with 0.1000 M HBr solution after the following additions of acid: (a) 0 mL (b) 15.00 mL (c) 29.00 mL (d) 29.90 mL (e) 30.00 mL (f) 30.10 mL (g) 40.00 mL

Find the pH during the titration of 20.00 mL of 0.1000 M butanoic acid, CH3CH2CH2COOH (Ka ? 1.54?105 ), with 0.1000 M NaOH solution after the following additions of titrant: (a) 0 mL (b) 10.00 mL (c) 15.00 mL (d) 19.00 mL (e) 19.95 mL (f) 20.00 mL (g) 20.05 mL (h) 25.00 mL

Find the pH during the titration of 20.00 mL of 0.1000 M triethylamine, (CH3CH2)3N (Kb ? 5.2?104 ), with 0.1000 M HCl solution after the following additions of titrant: (a) 0 mL (b) 10.00 mL (c) 15.00 mL (d) 19.00 mL (e) 19.95 mL (f) 20.00 mL (g) 20.05 mL (h) 25.00 mL

Find the pH of the equivalence point(s) and the volume (mL) of 0.0372 M NaOH needed to reach it in titrations of (a) 42.2 mL of 0.0520 M CH3COOH (b) 28.9 mL of 0.0850 M H2SO3 (two equivalence points)

Find the pH of the equivalence point(s) and the volume (mL) of 0.0588 M KOH needed to reach it in titrations of (a) 23.4 mL of 0.0390 M HNO2 (b) 17.3 mL of 0.130 M H2CO3 (two equivalence points)

Find the pH of the equivalence point(s) and the volume (mL) of 0.125 M HCl needed to reach it in titrations of (a) 65.5 mL of 0.234 M NH3 (b) 21.8 mL of 1.11 M CH3NH2

Find the pH and volume (mL) of 0.447 M HNO3 needed to reach the equivalence point(s) in titrations of (a) 2.65 L of 0.0750 M pyridine (C5H5N) (b) 0.188 L of 0.250 M ethylenediamine (H2NCH2CH2NH2)

The molar solubility of M2X is 5?105 M. What is the molarity of each ion? How do you set up the calculation to find Ksp? What assumption must you make about the dissociation of M2X into ions? Why is the calculated Ksp higher than the actual value?

Why does pH affect the solubility of BaF2 but not of BaCl2?

A list of Ksp values like that in Appendix C can be used to compare the solubility of silver chloride directly with that of silver bromide but not with that of silver chromate. Explain

In a gaseous equilibrium, the reverse reaction occurs when Qc ? Kc. What occurs in aqueous solution when Qsp ? Ksp?

Write the ion-product expressions for (a) silver carbonate; (b) barium fluoride; (c) copper(II) sulfide

Write the ion-product expressions for (a) iron(III) hydroxide; (b) barium phosphate; (c) tin(II) sulfide.

Write the ion-product expressions for (a) calcium chromate; (b) silver cyanide; (c) nickel(II) sulfide.

Write the ion-product expressions for (a) lead(II) iodide; (b) strontium sulfate; (c) cadmium sulfide

The solubility of silver carbonate is 0.032 M at 20C. Calculate its Ksp.

The solubility of zinc oxalate is 7.9103? M at 18C. Calculate its Ksp.

The solubility of silver dichromate at 15C is 8.310?3 g/100 mL solution. Calculate its Ksp.

The solubility of calcium sulfate at 30C is 0.209 g/100 mL solution. Calculate its Ksp.

Find the molar solubility of SrCO3 (Ksp ? 5.410?10) in (a) pure water and (b) 0.13 M Sr(NO3)2.

Find the molar solubility of BaCrO4 (Ksp ? 2.110?10) in (a) pure water and (b) 1.510?3 M Na2CrO4.

Calculate the molar solubility of Ca(IO3)2 in (a) 0.060 M Ca(NO3)2 and (b) 0.060 M NaIO3. (See Appendix C.)

Which compound in each pair is more soluble in water? (a) Magnesium hydroxide or nickel(II) hydroxide (b) Lead(II) sulfide or copper(II) sulfide (c) Silver sulfate or magnesium fluoride

Which compound in each pair is more soluble in water? (a) Strontium sulfate or barium chromate (b) Calcium carbonate or copper(II) carbonate (c) Barium iodate or silver chromate

Which compound in each pair is more soluble in water? (a) Barium sulfate or calcium sulfate (b) Calcium phosphate or magnesium phosphate (c) Silver chloride or lead(II) sulfate

Which compound in each pair is more soluble in water? (a) Barium sulfate or calcium sulfate (b) Calcium phosphate or magnesium phosphate (c) Silver chloride or lead(II) sulfate

Which compound in each pair is more soluble in water? (a) Manganese(II) hydroxide or calcium iodate (b) Strontium carbonate or cadmium sulfide (c) Silver cyanide or copper(I) iodide

Write equations to show whether the solubility of either of the following is affected by pH: (a) AgCl; (b) SrCO3

Write equations to show whether the solubility of either of the following is affected by pH: (a) CuBr; (b) Ca3(PO4)2.

Write equations to show whether the solubility of either of the following is affected by pH: (a) Fe(OH)2; (b) CuS.

Write equations to show whether the solubility of either of the following is affected by pH: (a) PbI2; (b) Hg2(CN)2.

Does any solid Cu(OH)2 form when 0.075 g of KOH is dissolved in 1.0 L of 1.010?3 M Cu(NO3)2?

Does any solid PbCl2 form when 3.5 mg of NaCl is dissolved in 0.250 L of 0.12 M Pb(NO3)2?

Does any solid Ba(IO3)2 form when 7.5 mg of BaCl2 is dissolved in 500. mL of 0.023 M NaIO3?

Does any solid Ag2CrO4 form when 2.710?5 g of AgNO3 is dissolved in 15.0 mL of 4.010?4 M K2CrO4?

When blood is donated, sodium oxalate solution is used to precipitate Ca2?, which triggers clotting. A 104-mL sample of blood contains 9.710?5 g Ca2?/mL. A technologist treats the sample with 100.0 mL of 0.1550 M Na2C2O4. Calculate [Ca2?] after the treatment. (See Appendix C for Ksp of CaC2O4H ? 2O.)

A 50.0-mL volume of 0.50 M Fe(NO3)3 is mixed with 125 mL of 0.25 M Cd(NO3)2. (a) If aqueous NaOH is added, which ion precipitates first? (See Appendix C.) (b) Describe how the metal ions can be separated using NaOH. (c) Calculate the [OH?] that will accomplish the separation.

How can a metal cation be at the center of a complex anion?

Write equations to show the stepwise reaction of Cd(H2O)4 2? in an aqueous solution of KI to form CdI4 2.? Show that Kf(overall) ? Kf1 Kf2 Kf3 Kf4. 1

Consider the dissolution of PbS in water: Adding aqueous NaOH causes more PbS to dissolve. Does this violate Le Chteliers principle? Explain.

Write a balanced equation for the reaction of Hg(H2O)4 2? in aqueous KCN

Write a balanced equation for the reaction of Zn(H2O)4 2? in aqueous NaCN.

White a balanced equation for the reaction of Ag(H2O)2 ? in aqueous Na2S2O3

Write a balanced equation for the reaction of Al(H2O)6 3? in aqueous KF

What is [Ag?] when 25.0 mL each of 0.044 M AgNO3 and 0.57 M Na2S2O3 are mixed [Kf of Ag(S2O3)2 3? ? 4.71013]?

Potassium thiocyanate, KSCN, is often used to detect the presence of Fe3? ions in solution through the formation of the red Fe(H2O)5SCN2? (or, more simply, FeSCN2?). What is [Fe3?] when 0.50 L each of 0.0015 M Fe(NO3)3 and 0.20 M KSCN are mixed? Kf of FeSCN2? ? 8.9102

Find the solubility of Cr(OH)3 in a buffer of pH 13.0 [Ksp of Cr(OH)3 ? 6.310?31; Kf of Cr(OH)4 ? ? 8.01029]. 19.102 Find the solubility of AgI in 2.5 M NH3 [Ksp of AgI ? 8.310?17; Kf of Ag(NH3)2 ? ? 1.7107 ]. 19

When 0.84 g of ZnCl2 is dissolved in 245 mL of 0.150 M NaCN, what are [Zn2?], [Zn(CN)4 2?], and [CN]? [Kf of Zn(CN)4 2? ? 4.21019]?

When 0.84 g of ZnCl2 is dissolved in 245 mL of 0.150 M NaCN, what are [Zn2?], [Zn(CN)4 2?], and [CN]? [Kf of Zn(CN)4 2? ? 4.21019]?

When 2.4 g of Co(NO3)2 is dissolved in 0.350 L of 0.22 M KOH, what are [Co2?], [Co(OH)4 2?], and [OH]? [Kf of Co(OH)4 2? ? 5109 ]?

What volumes of 0.200 M HCOOH and 2.00 M NaOH would make 500. mL of a buffer with the pH of one made from 475 mL of 0.200 M benzoic acid and 25 mL of 2.00 M NaOH?

A microbiologist is preparing a medium on which to culture E. coli bacteria. She buffers the medium at pH 7.00 to minimize the effect of acid-producing fermentation. What volumes of equimolar aqueous solutions of K2HPO4 and KH2PO4 must she combine to make 100. mL of the pH 7.00 buffer?

As an FDA physiologist, you need 0.700 L of formic acidformate buffer with a pH of 3.74. (a) What is the required buffer-component concentration ratio? (b) How do you prepare this solution from stock solutions of 1.0 M HCOOH and 1.0 M NaOH? (c) What is the final concentration of HCOOH in this solution?

Tris(hydroxymethyl)aminomethane [(HOCH2)3CNH2, known as TRIS] is a weak base used in biochemical experiments to make buffer solutions in the pH range of 7 to 9. A certain TRIS buffer has a pH of 8.10 at 25 C and a pH of 7.80 at 37 C. Why does the pH change with temperature?

Water flowing through pipes of carbon steel must be kept at pH 5 or greater to limit corrosion. If an 8.0?103 lb/hr water stream contains 10 ppm sulfuric acid and 0.015% acetic acid, how many pounds per hour of sodium acetate trihydrate must be added to maintain that pH?

Gout is caused by an error in metabolism that leads to a buildup of uric acid in body fluids, which is deposited as slightly soluble sodium urate (C5H3N4O3Na) in the joints. If the extracellular [Na] is 0.15 M and the solubility of sodium urate is 0.085 g/100. mL, what is the minimum urate ion concentration (abbreviated [Ur]) that will cause a deposit of sodium urate?

Cadmium ion in solution is analyzed by precipitation as the sulfide, a yellow compound used as a pigment in everything from artists oil paints to glass and rubber. Calculate the molar solubility of cadmium sulfide at 25 C.

In the discussion of cave formation in the Chemical Connections essay (p. 859), the dissolution of CO2 (equation 1) has a Keq of 3.1?102 , and the formation of aqueous Ca(HCO3)2 (equation 2) has a Keq of 1?1012. The fraction by volume of atmospheric CO2 is 3?104 . (a) Find [CO2(aq)] in equilibrium with atmospheric CO2. (b) Determine [Ca2] arising from (equation 2) given current levels of atmospheric CO2. (c) Calculate [Ca2] if atmospheric CO2 doubles. 19.

Phosphate systems form essential buffers in organisms. Calculate the pH of a buffer made by dissolving 0.80 mol of NaOH in 0.50 L of 1.0 M H3PO4.

The solubility of KCl is 3.7 M at 20 C. Two beakers contain 100. mL of saturated KCl solution: 100. mL of 6.0 M HCl is added to the first beaker and 100. mL of 12 M HCl to the second. (a) Find the ion- product constant of KCl at 20 C. (b) What mass, if any, of KCl will precipitate from each beaker?

It is possible to detect NH3 gas over 102 M NH3. To what pH must 0.15 M NH4Cl be raised to form detectable NH3?

Manganese(II) sulfide is one of the compounds found in the nodules on the ocean floor that may eventually be a primary source of many transition metals. The solubility of MnS is 4.7?104 g/100 mL solution. Estimate the Ksp of MnS.

The normal pH of blood is 7.40 ? 0.05 and is controlled in part by the H2CO3/HCO3 buffer system. (a) Assuming that the Ka value for carbonic acid at 25 C applies to blood, what is the [H2CO3]/[HCO3 ] ratio in normal blood? (b) In a condition called acidosis, the blood is too acidic. What is the [H2CO3]/[HCO3 ] ratio in a patient whose blood pH is 7.20? 1

A bioengineer preparing cells for cloning bathes a small piece of rat epithelial tissue in a TRIS buffer (see Problem 19.108). The buffer is made by dissolving 43.0 g of TRIS (pKb ? 5.91) in enough 0.095 M HCl to make 1.00 L of solution. What is the molarity of TRIS and the pH of the buffer?

Sketch a qualitative curve for the titration of ethylenediamine, H2NCH2CH2NH2, with 0.1 M HCl.

A solution contains 0.10 M ZnCl2 and 0.020 M MnCl2. Given the following information, how would you adjust the pH to separate the ions as their sulfides ([H2S] of a saturated aqueous solution at 25 C ? 0.10 M; Kw ? 1.0? 014 at 25 C)?

Amino acids [general formula NH2CH(R)COOH] can be considered polyprotic acids. In many cases, the R group contains additional amine and carboxyl groups. (a) Can an amino acid dissolved in pure water have a protonated COOH group and an unprotonated NH2 group (Ka of COOH group ? 4.47?103 ; Kb of NH2 group ? 6.03?105 )? Use glycine, NH2CH2COOH, to explain why. (b) Calculate [NH3CH2COO]/[NH3CH2COOH] at pH 5.5. (c) The R group of lysine is CH2CH2CH2CH2NH2 (pKb ? 3.47). Draw the structure of lysine at pH 1, physiological pH (~7), and pH 13. (d) The R group of glutamic acid is CH2CH2COOH (pKa ? 4.07). Of the forms of glutamic acid that are shown below, which predominates at (1) pH 1, (2) physiological pH ( 7), and (3) pH 13? 19.

The scene below depicts a saturated solution of MCl2(s) in the presence of dilute aqueous NaCl; each sphere represents 1.0?106 mol of ion, and the volume is 250.0 mL (solid MCl2 is shown as green chunks; Na ions and water molecules are not shown). (a) Calculate the Ksp of MCl2. (b) If M(NO3)2(s) is added, is there an increase, decrease, or nochange in the number of Cl particles? In the Ksp? In the mass of MCl2(s)?

Tooth enamel consists of hydroxyapatite, Ca5(PO4)3OH (Ksp ? 6.8?1037). Fluoride ion added to drinking water reacts with Ca5(PO4)3OH to form the more tooth decayresistant fluorapatite, Ca5(PO4)3F (Ksp ? 1.0?1060). Fluoridated water has dramatically decreased cavities among children. Calculate the solubility of Ca5(PO4)3OH and of Ca5(PO4)3F in water.

The acid-base indicator ethyl orange turns from red to yellow over the pH range 3.4 to 4.8. Estimate Ka for ethyl orange.

Use the values obtained in Problem 19.54 to sketch a curve of [H3O] vs. mL of added titrant. Are there advantages or disadvantages to viewing the results in this form? Explain.

6 Instrumental acid-base titrations use a pH meter to monitor the changes in pH and volume. The equivalence point is found from the volume at which the curve has the steepest slope. (a) Use Figure 19.7 to calculate the slope pH/ V for all pairs of adjacent points and to calculate the average volume (Vavg) for each interval. (b) Plot pH/ V vs. Vavg to find the steepest slope, and thus the volume at the equivalence point. (For example, the first pair of points gives pH ? 0.22, V ? 10.00 mL; hence, pH/ V ? 0.022 mL1 , and Vavg ? 5.00 mL.)

What is the pH of a solution of 6.5?109 mol of Ca(OH)2 in 10.0 L of water [Ksp of Ca(OH)2 ? 6.5?106 ]?

The Henderson-Hasselbalch equation gives a relationship for obtaining the pH of a buffer solution consisting of HA and A. Derive an analogous relationship for obtaining the pOH of a buffer solution consisting of B and BH. 1

Muscle physiologists study the accumulation of lactic acid [CH3CH(OH)COOH] during exercise. Food chemists study its occurrence in sour milk, beer, wine, and fruit. Industrial microbiologists study its formation by various bacterial species from carbohydrates. A biochemist prepares a lactic acidlactate buffer by mixing 225 mL of 0.85 M lactic acid (Ka ? 1.38?104 ) with 435 mL of 0.68 M sodium lactate. What is the buffer pH?

A student wants to dissolve the maximum amount of CaF2 (Ksp ? 3.2?1011) to make 1 L of aqueous solution. (a) Into which of the following solvents should she dissolve the salt? (I) Pure water (II) 0.01 M HF (III) 0.01 M NaOH (IV) 0.01 M HCl (V) 0.01 M Ca(OH)2 (b) Which would dissolve the least amount of salt?

A 500.-mL solution consists of 0.050 mol of solid NaOH and 0.13 mol of hypochlorous acid (HClO; Ka ? 3.0?108 ) dissolved in water. (a) Aside from water, what is the concentration of each species that is present? (b) What is the pH of the solution? (c) What is the pH after adding 0.0050 mol of HCl to the flask?

Calcium ion present in water supplies is easily precipitated as calcite (CaCO3): CaCO3 Ca (s) Because the Ksp decreases with temperature, heating hard water forms a calcite scale, which clogs pipes and water heaters. Find the solubility of calcite in water (a) at 10 C (Ksp ? 4.4?109 ) and (b) at 30 C (Ksp ? 3.1?109).

Calculate the molar solubility of Hg2C2O4 (Ksp ? 1.75?1013) in 0.13 M Hg2(NO3)2.

The well water in an area is hard because it is in equilibrium with CaCO3 in the surrounding rocks. What is the concentration of Ca2 in the well water (assuming the waters pH is such that the CO3 2 ion is not significantly protonated)? (See Appendix C for Ksp of CaCO3.)

Environmental engineers use alkalinity as a measure of the capacity of carbonate buffering systems in water samples: Find the alkalinity of a water sample that has a pH of 9.5, 26.0 mg/L CO3 2, and 65.0 mg/L HCO3 .

Human blood contains one buffer system based on phosphate species and one on carbonate species. Assuming that blood has a normal pH of 7.4, what are the principal phosphate and carbonate species present? What is the ratio of the two phosphate species? (In the presence of the dissolved ions and other species in blood, Ka1 of H3PO4 ? 1.3?102 , Ka2 ? 2.3?107 , and Ka3 ? 6?1012; Ka1 of H2CO3 ? 8?107 and Ka2 ? 1.6?1010.) 19.

Litmus is an organic dye extracted from lichens. It is red below pH 4.5 and blue above pH 8.3. One drop of either 0.1 M HCl (pH 1) or a pH 3 buffer changes blue litmus paper to red, but a drop of 0.001 M HCl (also pH 3) does not. Explain

An environmental technician collects a sample of rainwater. Back in the lab, her pH meter isnt working, so she uses indicator solutions to estimate the pH. A piece of litmus paper turns red, indicating acidity, so she divides the sample into thirds and obtains the following results: thymol blue turns yellow; bromphenol blue turns green; and methyl red turns red. Estimate the pH of the rainwater

A 0.050 M H2S solution contains 0.15 M NiCl2 and 0.35 M Hg(NO3)2. What pH is required to precipitate the maximum amount of HgS but none of the NiS? (See Appendix C.)

Quantitative analysis of Cl ion is often performed by a titration with silver nitrate, using sodium chromate as an indicator. As standardized AgNO3 is added, both white AgCl and red Ag2CrO4 precipitate, but so long as some Cl remains, the Ag2CrO4 redissolves as the mixture is stirred. When the red color is permanent, the equivalence point has been reached. (a) Calculate the equilibrium constant for the reaction (b) Explain why the silver chromate redissolves. (c) If 25.00 cm3 of 0.1000 M NaCl is mixed with 25.00 cm3 of 0.1000 M AgNO3, what is the concentration of Ag remaining in solution? Is this sufficient to precipitate any silver chromate? 1

An ecobotanist separates the components of a tropical bark extract by chromatography. She discovers a large proportion of quinidine, a dextrorotatory isomer of quinine used for control of arrhythmic heartbeat. Quinidine has two basic nitrogens (Kb1 ? 4.0?106 and Kb2 ? 1.0?1010). To measure the concentration, she carries out a titration. Because of the low solubility of quinidine, she first protonates both nitrogens with excess HCl and titrates the acidified solution with standardized base. A 33.85-mg sample of quinidine is acidified with 6.55 mL of 0.150 M HCl. (a) How many milliliters of 0.0133 M NaOH are needed to titrate the excess HCl? (b) How many additional milliliters of titrant are needed to reach the first equivalence point of quinidine dihydrochloride? (c) What is the pH at the first equivalence point?

Some kidney stones form by the precipitation of calcium oxalate monohydrate (CaC2O4H2O, Ksp 2.3?10?9 ). The pH of urine varies from 5.5 to 7.0, and the average [Ca2?] in urine is 2.6?10?3 M. (a) If the [oxalic acid] in urine is 3.0?10?13 M, will kidney stones form at pH 5.5? (b) At pH 7.0? (c) Vegetarians have a urine pH above 7. Are they more or less likely to form kidney stones? 19

A biochemist needs a medium for acid-producing bacteria. The pH of the medium must not change by more than 0.05 pH units for every 0.0010 mol of H3O? generated by the organisms per liter of medium. A buffer consisting of 0.10 M HA and 0.10 M A? is included in the medium to control its pH. What volume of this buffer must be included in 1.0 L of medium?

A 35.00-mL solution of 0.2500 M HF is titrated with a standardized 0.1532 M solution of NaOH at 25?C. (a) What is the pH of the HF solution before titrant is added? (b) How many milliliters of titrant are required to reach the equivalence point? (c) What is the pH at 0.50 mL before the equivalence point? (d) What is the pH at the equivalence point? (e) What is the pH at 0.50 mL after the equivalence point?

Because of the toxicity of mercury compounds, mercury(I) chloride is used in antibacterial salves. The mercury(I) ion (Hg2 2?) consists of two bound Hg? ions. (a) What is the empirical formula of mercury(I) chloride? (b) Calculate [Hg2 2?] in a saturated solution of mercury(I) chloride (Ksp 1.5?10?18). (c) A seawater sample contains 0.20 lb of NaCl per gallon. Find [Hg2 2?] if the seawater is saturated with mercury(I) chloride. (d) How many grams of mercury(I) chloride are needed to saturate 4900 km3 of pure water (the volume of Lake Michigan)? (e) How many grams of mercury(I) chloride are needed to saturate 4900 km3 of seawater?

A lake that has a surface area of 10.0 acres (1 acre 4.840?103 yd2 ) receives 1.00 in. of rain of pH 4.20. (Assume that the acidity of the rain is due to a strong, monoprotic acid.) (a) How many moles of H3O? are in the rain falling on the lake? (b) If the lake is unbuffered (pH 7.00) and its average depth is 10.0 ft before the rain, find the pH after the rain has been mixed with lake water. (Ignore runoff from the surrounding land.) (c) If the lake contains hydrogen carbonate ions (HCO3 ?), what mass of HCO3 ? would neutralize the acid in the rain?

A 35.0-mL solution of 0.075 M CaCl2 is mixed with 25.0 mL of 0.090 M BaCl2. (a) If aqueous KF is added, which fluoride precipitates first? (b) Describe how the metal ions can be separated using KF to form the fluorides. (c) Calculate the fluoride ion concentration that will accomplish the separation

Even before the industrial age, rainwater was slightly acidic due to dissolved CO2. Use the following data to calculate pH of unpolluted rainwater at 25?C: vol % in air of CO2 0.033 vol %; solubility of CO2 in pure water at 25C? and 1 atm 88 mL CO2/100 mL H2O; Ka1 of H2CO3 4.5?107 ? .

Seawater at the surface has a pH of about 8.5. (a) Which of the following species has the highest concentration at this pH: H2CO3; HCO3 ?; CO3 2?? Explain. (b) What are the concentration ratios [CO3 2?]/[HCO3 ?] and [HCO3 ?]/[H2CO3] at this pH? (c) In the deep sea, light levels are low, and the pH is around 7.5. Suggest a reason for the lower pH at the greater ocean depth. (Hint: Consider the presence or absence of plant and animal life, and the effects on carbon dioxide concentrations.)

Ethylenediaminetetraacetic acid (abbreviated H4EDTA) is a tetraprotic acid. Its salts are used to treat toxic metal poisoning by forming soluble complex ions that are then excreted. Because EDTA4? also binds essential calcium ions, it is often administered as the calcium disodium salt. For example, when Na2Ca(EDTA) is given to a patient, the [Ca(EDTA)]2? ions react with circulating Pb2? ions and the metal ions are exchanged: A child has a dangerous blood lead level of 120 ?g/100 mL. If the child is administered 100. mL of 0.10 M Na2Ca(EDTA), assuming the exchange reaction and excretion process are 100% efficient, what is the final concentration of Pb2? in ?g/100 mL blood? (Total blood volume is 1.5 L.)

Buffers that are based on 3-morpholinopropanesulfonic acid (MOPS) are often used in RNA analysis. The useful pH range of a MOPS buffer is 6.5 to 7.9. Estimate the Ka of MOPS.

NaCl is purified by adding HCl to a saturated solution of NaCl (317 g/L). When 28.5 mL of 8.65 M HCl is added to 0.100 L of saturated solution, what mass (g) of pure NaCl precipitates?

Scenes A to D represent tiny portions of 0.10 M aqueous solutions of a weak acid HA (red and blue; Ka 4.5?10?5 ), its conjugate base A?(red), or a mixture of the two (only these species are shown): (a) Which scene(s) show(s) a buffer? (b) What is the pH of each solution? (c) Arrange the scenes in sequence, assuming that they represent stages in a weak acidstrong base titration. (d) Which scene represents the titration at its equivalence point?

Scenes A to C represent aqueous solutions of the slightly soluble salt MZ (only the ions of this salt are shown): (a) Which scene represents the solution just after solid MZ is stirred thoroughly in distilled water? (b) If each sphere represents 2.5106 M of ions, what is the Ksp of MZ? (c) Which scene represents the solution after Na2Z(aq) is added? (d) If Z2 is CO3 2, which scene represents the solution after the pH has been lowered? M2?(aq) ? Z2 MZ(s) (aq) 19.

The solubility of Ag(I) in aqueous solutions containing different concentrations of Cl is based on the following equilibria: When solid AgCl is shaken with a solution containing Cl, Ag(I) is present as both Ag? and AgCl2 . The solubility of AgCl is the sum of the concentrations of Ag? and AgCl2 . (a) Show that [Ag?] in solution is given by ] and that [AgCl2 ] in solution is given by (b) Find the [Cl] at which [Ag?] ? [AgCl2 ]. (c) Explain the shape of a plot of AgCl solubility vs. [Cl]. (d) Find the solubility of AgCl at the [Cl] of part (b), which is the minimum solubility of AgCl in the presence of Cl. 19.156

EDTA binds metal ions to form complex ions (see Problem 19.150), so it is used to determine the concentrations of metal ions in solution: A 50.0-mL sample of 0.048 M Co2? is titrated with 0.050 M EDTA4. Find [Co2?] and [EDTA4] after (a) 25.0 mL and (b) 75.0 mL of EDTA4 are added (log Kf of CoEDTA2? 16.31). M