Ch.17 - Applications of Aqueous Equilibria
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- Is the pH greater than, equal to, or less than 7 after the neutralization of each of the following pairs of acids and bases? (b) NH3 and HClO4 (d) (CH3)3N and HOBr
Problem 48
Problem 48a
Is the pH greater than, equal to, or less than 7 after the neutralization of each of the following pairs of acids and bases? (a) NaOH and benzoic acid (C6H5CO2H)
Problem 48c
Is the pH greater than, equal to, or less than 7 after the neutralization of each of the following pairs of acids and bases? (c) KOH and HI
- Which of the following mixtures has the higher pH? (a) Equal volumes of 0.10 M HCN and 0.10 M NaOH (b) Equal volumes of 0.10 M HClO4 and 0.10 M NaOH
Problem 49
- Phenol (C6H5OH, Ka = 1.3 x 10^-10) is a weak acid used in mouthwashes, and pyridine (C5H5N, Kb = 1.8 x 10^-9) is a weak base used as a solvent. Calculate the value of Kn for the neutralization of phenol by pyridine. Does the neutralization reaction proceed very far toward completion?
Problem 51
- The equilibrium constant Kn for the neutralization of lactic acid (C3H6O3) and urea (CH4N2O) is 2.1 x 10^-4. What is Kb for urea? The Ka of lactic acid is 1.4 x 10^-4.
Problem 53
- The equilibrium constant Kn for the neutralization of boric acid (H3BO3) and caffeine (C8H10N4O2) is 24. What is Kb for caffeine? The Ka of boric acid is 5.8 x 10^-10.
Problem 54
- Does the pH increase, decrease, or remain the same when the substances are added to the solutions? (a) LiF to an HF solution (b) KI to an HI solution (c) NH4Cl to an NH3 solution
Problem 55
- Does the pH increase, decrease, or remain the same on the addition of each of the following? (a) NH4NO3 to an NH3 solution (b) Na2CO3 to an NaHCO3 solution (c) NaClO4 to an NaOH solution
Problem 56
- Calculate the pH of a solution that is 0.25 M in HF and 0.10 M in NaF.
Problem 57
- Calculate the pH of a solution prepared by dissolving 0.10 mol of solid NH4Cl in 0.500 L of 0.40 M NH3. Assume that there is no volume change.
Problem 58
- The pH of a solution of HN3 (Ka = 1.9 x10^-5) and NaN3 is 4.86. What is the molarity of NaN3 if the molarity of HN3 is 0.016 M?
Problem 59
Problem 60
The pH of a solution of NH3 and NH4Br is 8.90. What is the molarity of NH4Br if the molarity of NH3 is 0.016 M?
- Which of the following gives a buffer solution when equal volumes of the two solutions are mixed? (a) 0.10 M HF and 0.10 M NaF (b) 0.10 M HF and 0.10 M NaOH (c) 0.20 M HF and 0.10 M NaOH (d) 0.10 M HCl and 0.20 M NaF
Problem 63
- Which of the following gives a buffer solution when equal volumes of the two solutions are mixed? (a) 0.10 M NH3 and 0.10 M HCl (b) 0.20 M NH3 and 0.10 M HCl (c) 0.10 M NH4Cl and 0.10 M NH3 (d) 0.20 M NH4Cl and 0.10 M NaOH
Problem 64
- Which of the following solutions has the greater buffer capacity: 100 mL of 0.30 M HNO2-0.30 M NaNO2 or 100 mL of 0.10 M HNO2-0.10 M NaNO2? Explain.
Problem 65
- Which of the following solutions has the greater buffer capacity: 50 mL of 0.20 M NH4Br-0.30 M NH3 or 50 mL of 0.40 M NH4Br-0.60 M NH3? Explain.
Problem 66
- Calculate the pH of a buffer solution that is 0.20 M in HCN and 0.12 M in NaCN. Will the pH change if the solution is diluted by a factor of 2? Explain.
Problem 67
- Calculate the pH of a buffer solution prepared by dissolving 4.2 g of NaHCO3 and 5.3 g of Na2CO3 in 0.20 L of water. Will the pH change if the solution volume is increased by a factor of 10? Explain.
Problem 68
- Calculate the pH of 0.250 L of a 0.36 M formic acid–0.30 M sodium formate buffer before and after the addition of (a) 0.0050 mol of NaOH. Assume that the volume remains constant. (b) 0.0050 mol of HCl. Assume that the volume remains constant.
Problem 69
Problem 70a
Calculate the pH of 0.375 L of a 0.18 M acetic acid–0.29 M sodium acetate buffer before and after the addition of (a) 0.0060 mol of KOH. Assume that the volume remains constant.
Problem 70b
Calculate the pH of 0.375 L of a 0.18 M acetic acid–0.29 M sodium acetate buffer before and after the addition of (b) 0.0060 mol of HBr. Assume that the volume remains constant.
- Use the Henderson–Hasselbalch equation to calculate the pH of a buffer solution that is 0.25 M in formic acid (HCO2H) and 0.50 M in sodium formate (HCO2Na).
Problem 71
- A food chemist studying the formation of lactic acid in sour milk prepares a buffer that is 0.58 M in lactic acid (HC3H5O3) and 0.36 M in sodium lactate (NaC3H5O3). Use the Henderson–Hasselbalch equation to calculate the pH of the buffer solution.
Problem 72
- What is the Henderson–Hasselbalch equation to calculate the ratio of HCO3- and H2CO3 in blood with a pH of 7.40, given that the value of Ka for carbonic acid at body temperature (37 degrees Celsius) is 7.9 x 10^-7?
Problem 73
- The ratio of HCO3- to H2CO3 in blood is called the 'bicarb number' and is used as a measure of blood pH in hospital emergency rooms. A newly diagnosed diabetic patient is admitted to the emergency room with ketoacidosis and a bicarb number of 10. Calculate the blood pH. Ka for carbonic acid at room temperature (37 degrees Celsius) os 7.9 x 10^-7).
Problem 74
- In what volume ratio should you mix 1.0 M solutions of NH4Cl and NH3 to produce a buffer solution having pH = 9.80?
Problem 75
- Give a recipe for preparing a CH3CO2H-CH3CO2Na buffer solution that has pH = 4.44.
Problem 76
- You need a buffer solution that has pH = 7.00. Which of the following buffer systems should you choose? Explain. (a) H3PO4 and H2PO4 - (b) H2PO4- and HPO42- (c) HPO42- and PO43-
Problem 77
- Which of the following conjugate acid–base pairs should you choose to prepare a buffer solution that has pH = 4.50? Explain. (a) HSO4- and SO42- (b) HOCl and OCl- (c) C6H5CO2H and C6H5CO2-.
Problem 78