Chapter 11 Problems - E i i E i i r 1 Problems “Pepsin...

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Unformatted text preview: E i i E i i r 1 Problems “Pepsin and Antacid Therapy: A Dilemma,” W. B. Batson and P. H. Laswick, J. Chem. Educ. 56, 484 (1979). “What is the Energy Difference Between H2NCH2COOH and +H3NCH2C0§?” P. Haberfield, J. Chem. Educ. 57, 346 (1980). Problems _____—____———-————————-———— Acids, Bases, Dissociation Constants, and pH 11.1 Classify each of the following species as a Bronsted acid or base, or both: (a) H20, (b) 0H‘, (0) 1130+, (d) NH3, (6) NH1, (1") NHL (g) N05, (11) C033 0) HBr, (i) HCN, (k) HCOa— - 11.2 Write the formulas for the conjugate bases of the following acids: (21) HI, (b) H2804, (c) H2S, (d) HCN, (e) HCOOH (formic acid). 11.3 Classify each of the following species as a weak or strong acid: (a) HN03, (b) HF, (c) H2804, ((1) H801, (e) H2CO3, (f) HCO;, (g) HCl, (h) HCN, (i) HNO2. 11.4 Classify each of the following species as a weak or strong base: (a) LiOH, (b) CN‘, (c) H20, (d) C10; (e) NH; . 11.5 Calculate the pH of the following solutions: (a) 1.0 M HCl, (b) 0.10 M HCl, (c) 1.0 x 10"2 M HCl, (d) 1.0 x 10‘2 M NaOH, (e) 1.0 x 10‘2 M Ba(0H)2. Assume ideal behavior. 11.6 A 0.040 M solution of a monoprotic acid is 13.5% dissociated. What is the dissociation constant of the acid? 11.7 Write the equation relating Ka for a weak acid and Kb for its conjugate base. Use NH; and its conjugate acid NH: to derive the relationship between Ka and Kb. 11.8 The dissociation constant of a monoprotic acid at 298 K is 1.47 X 10‘3. Calculate the degree of dissociation by (a) assuming ideal behavior and (b) using a mean activity coefficient yi = 0.93. The concentration of the acid is 0.010 M. 11.9 The ion product of D20 is 1.35 x 10’15 at 25 °C. (a) Calculate the value of pD for pure D20 where pD = —log[D+]. (b) For what values of pD will a solution be acidic in D20? (c) Derive a relation between pD and pOD. 11.10 HF is a weak acid, but its strength increases with concentration. Explain. (Hint: F' reacts with HF to form HFZ'. The equilibrium constant for this reaction is 5.2 at 25 °C.) 11.11 When the concentration of a strong acid is not substantially higher than 1.0 x 10‘7 M, the ionization of water must be taken into account in the calculation of the solution’s pH. (a) Derive an expression for the pH of a strong acid solution, including the contribution to [H+] from H20. (b) Calculate the pH of a 1.0 x 10"7 M HCl solution. 11.12 What are the concentrations of HSOI, 803—, and H+ in a 0.20 M KHSO4 solution? (Hint: H2804 is a strong acid; Ka for H80; 2 1.3 x 10‘2.) 11.13 Calculate the concentrations of H+, HCO3_, and C03’ in a 0.025 M H2C03 solution. 11.14 To which of the following would the addition of an equal volume of 0.60 M NaOH lead to a solution having a lower pH? (a) Water, (b) 0.30 M HCl, (c) 0.70 M KOH, (d) 0.40 M NaN03. 11.15 A solution contains a weak monoprotic acid, HA, and its sodium salt, NaA, both at 0.1 M concentration. Show that [OH‘] = Kw/Ka. 11.16 A solution of methylamine (CH3NH2) has a pH of 10.64. How many grams of methylamine are in 100.0 mL of the solution? 439 440 Chapter 11: Acids and Bases 11.17 Hydrocyanic acid (HCN) is a weak acid and a deadly poisonous compound that is used in gas chambers in the gaseous form (hydrogen cyanide). Why is it dangerous to treat sodium cyanide with acids (such as HCl) without proper ventilation? 11.18 Novocaine, used as a local anesthetic by dentists, is a weak base (Kb = 8.91 x 10‘6). What is the ratio of the concentration of the base to that of its acid in the blood plasma (pH = 7.40) of a patient? 11.19 Calculate the,percent dissociation of HF at the following concentrations: (a) 0.50 M and (b) 0.050 M ." Comment on your results. 11.20 Explain why phenol is a stronger acid than methanol: OOH CHa—OH Phenol Methanol 11.21 Calculate the concentrations of all species in a 0.100 M H3PO4 solution. 11.22 The disagreeable odor of fish is mainly due to organic compounds (RNHz) containing an amino group, —NH2, where R is the rest of the molecule. Amines are bases just like ammonia. Explain why putting some lemon juice on fish can greatly reduce the odor. Salt Hydrolysis 11.23 Specify which of the following salts will undergo hydrolysis: KF, NaNO3, NH4N02, MgSO4, KCN, C6H5COONa, Rbl, Na2C03, CaClz, HCOOK. 11.24 Calculate the pH of a 0.10 M NH4Cl solution. 11.25 Calculate the pH and percent hydrolysis of a 0.36 M CH3COONa solution. Acid—Base Titration 11.26 A student added NaOH solution from a buret to an Erlenmeyer flask containing HCl solution and used phenolphthalein as indicator. At the equivalence point of the titration, she observed a faint reddish-pink color. However, after a few minutes, the solution gradually turned colorless. What do you suppose happened? 11.27 The ionization constant, [(3, of an indicator, HIn, is 1.0 x 10‘s. The color of the nonionized form is red and that of the ionized form is yellow. What is the color of this indicator in a solution whose pH is 4.00? 11.28 The Ka of a certain indicator is 2.0 X 10‘6. The color of HIn is green, and that of In‘ is red. A few drops of the indicator are added to a HCl solution, which is then titrated against a NaOH solution. At what pH will the indicator change color? 11.29 The pKa of the indicator methyl orange is 3.46. Over what pH range does this indicator change from 90% HIn to 90% In‘? 11.30 A 200-mL volume of NaOH solution was added to 400 mL of a 2.00 M HNO; solution. The pH of the mixed solution was 1.50 units greater than that of the original acid solution. Calculate the molarity of the NaOH solution. 11.31 A volume of 25.0 mL of 0.100 M HCl is titrated with a 0.100 M CH3NH2 solution. Calculate the pH values of the solution (a) after 10.0 mL of CH3NH2 solution have been added, (b) after 25.0 mL of CH3NH2 solution have been added, and (c) after 35.0 mL of CH3NH2 solution have been added. 11.32 Phenolphthalein is the common indicator for the titration of a strong acid with a strong base. (a) If the pKa of phenolphthalein is 9.10, what is the ratio of the nonionized form of the indicator (colorless) to the ionized form (reddish pink) at pH 8.00? (b) If 2 drops of 0.060 M Problems phenolphthalein are used in a titration involving a 50.0—mL volume, what is the concentration of the ionized form at pH 8.00? (Assume that 1 drop = 0.050 mL.) 11.33 Shown below is a titration curve for carbonic acid versus sodium hydroxide. Fill in the missing species and the pH and pKa values. 14 12 1O Volume of NaOH added Buffer Solutions 11.34 Specify which of the following systems can be classified as a buffer system: (a) KCl/HCI, (b) NH3/NH4NO3, (C) NazHPO4/N3H2PO4, KNOz/HNoz, (e) KHSO4/H2804, (f) HCOOK/HCOOH. 11.35 Derive the Henderson—Hasselbalch equation for the buffer system NHj/NHg. 11.36 Calculate the pH of the 0.20 M NH3/0.20 M NH4C1 bufl‘er. What is the pH of the bulTer after the addition of 10.0 mL of 0.10 M HCI to 65.0 mL of the bulTer? 11.37 Calculate the pH of 1.00 L of the buiTer 1.00 M CH3COONa/1.00 M CH3COOH before and after the addition of (a) 0.080 mol NaOH and (b) 0.12 mol HCl. (Assume that there is no change in volume.) 11.38 A quantity of 26.4 mL of a 0.45 M acetic acid solution is added to 31.9 mL of a 0.37 M sodium hydroxide solution. What is the pH of the final solution? 11.39 What is the pH of the buffer 0.10 M NazHPO4/0.10 M KH2P04? Calculate the concentration of all the species in solution. 11.40 A phosphate buffer has a pH equal to 7.30. (a) What is the predominant conjugate pair present in this buffer? (b) If the concentration of this bufl‘er is 0.10 M , what is the new pH after the addition of 5.0 mL of 0.10 M HCl to 20.0 mL of this buffer solution? 11.41 Tris[tris(hydroxymethyl)aminomethane] is a common buffer for studying biological systems HOCH2 HOCH2 1 + pKa=a.1 I HOCHz—(ID—NHa HOCHz—CID—NHZ + H+ HOCH2 HOCH2 (a) Calculate the pH of the tris bulfer after mixing 15.0 mL of 0.10 M HCl solution with 25.0 mL of 0.10 M tris. (b) This buffer was used to study an enzyme-catalyzed reaction. As a result of the reaction, 0.00015 mole of H+ was consumed. What is the pH of the buffer at the end of the reaction? (c) What would be the final pH if no buffer were present? Chapter 11: Acids and Bases 11.42 Describe the number of dilTerent ways to prepare 1 liter of a 0.050 M phosphate buffer with a pH of 7.8. 11.43 Calculate the concentration of all the species present in a solution that is 0.12 M in HCN and 0.34 M in NaCN. What is the pH of the solution? Does the solution possess buffer capacity? 11.44 The pH of blood plasma is 7.40. Assuming the principal buffer system is HCO;/H2CO3, r calculate the ratio [HCO;]/[H2CO3]. Is this buffer more effective against an added acid or an added base? / 11.45 A student is asked to prepare a buffer solution with pH = 8.60, using one of the following weak acids: HA (Ka = 2.7 x 10—3), HB (Ka = 4.4 x 10—6), HC (Ka = 2.6 x 10’9). Which acid should she choose? 11.46 The buffer range is defined by the equation pH = pKa i 1. Calculate the range of the ratio [conjugate base]/[acid] that corresponds to this equation. 11.47 Describe how you would prepare 1 L of 0.20 M CH3COONa/0.20 M CH3COOH bufi‘er system by (a) mixing a solution of CH3COOH with a solution of CH3COONa, (b) reacting a solution of CH3COOH with a solution of NaOH, and (c) reacting a solution of CH3COONa with a solution of HCl. 11.48 How many milliliters of 1.0 M NaOH must be added to 200 mL of 0.10 M NaH2P04 to make a bufler solution with a pH of 7.50? 11.49 Suggest two chemical tests that would allow you to distinguish an acid solution and a bulTer solution both at pH = 3.5. 11.50 How would you prepare a CH3COOH/CH3COONa buffer with a pH of 4.40 and an ionic strength of 0.050 m? Treat molarity the same as molality. 11.51 The pH of a phosphate buffer is 7.10 at 25 °C. What is the pH of the buffer at 37 °C? The A,H° for the relevant dissociation step is 3.75 kJ mol". Amino Acids 11.52 Which of the amino acids listed in Table 11.4 have a buffer capacity in the physiological region of pH 7? 11.53 Calculate the ionic strength of a 0.035 M serine buffer at pH 9.15. 11.54 From the pKa values listed in Table 11.4, calculate the p] value for amino acids lysine and valine. 11.55 Sketch the titration curve for 100 mL of 0.1 M aspartic acid hydrogen chloride titrated with sodium hydroxide. 11.56 At neutral pH, amino acids exist as dipolar ions. Using glycine as an example, and given that the pKa of the carboxyl group is 2.3 and that of the ammonium group is 9.6, predict the predominant form of the molecule at pHs of 1, 7, and 12. Justify your answers using Equation 11.16. Additional Problems 11.57 Describe a procedure that would allow you to compare the strength of Lewis acids. 11.58 From the dependence of KW on temperature (see p. 403), calculate the enthalpy of dissociation for water. 11.59 Freshly distilled, deionized water has a pH of 7. Left standing in air, however, the water gradually becomes acidic. Calculate the pH of the “solution” at equilibrium. (Hint: First calculate the solubility of C02 in water according to Example 7.3. Assume the partial pressure of C02 is 0.00030 atm.) 11.60 Show that the acid dissociation constant, Ka, of a weak monoprotic acid in water is related to its concentration, c (mol L"), and its degree of dissociation, oz, by Ka = azc/(l — or) if the Problems self-dissociation of water is ignored. If the latter is taken into account, show that Ka 2%(12611 + (1 + 4Kwoc‘2c”2)l/2]/(l — a). 11.61 To correct for the effect of ionic strength, we can write the dissociation constant of an acid as 0.509x/T 1 + x/T where Ka is the acid dissociation at zero ionic strength and K; the corresponding value at ionic strength, 1. Calculate the dissociation constant of acetic acid in a 0.15 m KCl solution at 298 K. You may neglect the ionic strength contribution due to the dissociation of the acid itself. 19K; = pKa — 11.62 Depending on the pH of the solution, ferric ions (Fe3+) may exist in the free-ion form or form the insoluble precipitate Fe(OH)3 (KSp = 1.0 x 10—36). Calculate the pH at which 90% of the Fe3+ ions in a 4.5 x 10—5 M Fe3+ solution would be precipitated. What conclusion can you draw about the Fe3+ ion concentration in blood plasma whose pH is 7.40? 11.63 A 0.020 M aqueous solution of benzoic acid has a freezing point of —0.0392 °C. Calculate the dissociation constant of benzoic acid. Assume ideal behavior, and assume that molality is equal to molality at this low concentration. 11.64 The pH of gastric juice is about 1.00 and blood plasma is 7.40. Calculate the Gibbs energy required to secrete a mole of HJr ions from blood plasma to the stomach at 37 °C. Assume ideal behavior. 11.65 Chemical analysis shows that 20.0 mL of a certain sample of blood yields 12.5 mL of C02 gas (measured at 25 °C and 1 atm) when treated with an acid. Calculate (a) the number of moles of CO2 originally present in the blood, (b) the concentration of CO2 and HCO; at equilibrium, and (c) the partial pressure of CO2 over the blood solution at equilibrium. Assume ideal behavior. The pH of blood is 7.40, and the Henry’s law constant for C02 in blood is 29.3 atm mol—1 (kg H2O). 11.66 Calcium oxalate is a major component of kidney stones. From the dissociation constants listed in Table 11.1 and given that the solubility product of CaC2O4 is 3.0 X 10—9, predict whether the formation of kidney stones can be minimized by increasing or decreasing the pH of the fluid present in the kidneys. The pH of normal kidney fluid is about 8.2. 11.67 What is the pH of a 0.050 M glycine solution at 298 K? 11.68 From the dissociation constant of formic acid listed in Table 11.1, calculate the Gibbs energy and the standard Gibbs energy for the dissociation of formic acid at 298 K. 11.69 (a) Calculate the percent ionization of a 0.20 M solution of the monoprotic acetylsalicylic acid (aspirin, C9H304), for which K3 2 3.0 x 104. (b) The pH of gastric juice in the stomach of a certain individual is 1.00. After a few aspirin tablets have been swallowed, the concentration of acetylsalicylic acid in the stomach is 0.20 M. Calculate the percent ionization of the acid under these conditions. What effect does the nonionized acid have on the membranes lining the stomach? 11.70 A 0.400 M formic acid (HCOOH) solution freezes at —0.758 °C. Calculate the value of Ka at that temperature. (Hint: Assume that molarity is equal to molality.) 11.71 Explain the action of smelling salts, which is ammonium nitrate [(NH4)2CO3]. (Hint: The thin film of aqueous solution that lines the nasal passage is slightly basic.) 11.72 Acid—base reactions usually go to completion. Confirm this statement by calculating the equilibrium constant for each of the following cases: (a) a strong acid reacting with a strong base, (b) a strong acid reacting with a weak base (NH3), (c) a weak acid (CH3COOH) reacting with a strong base, and (d) a weak acid (CH3COOH) reacting with a weak base (NH3). (Hint: Strong acids exist as H+ ions, and strong bases exist as OH‘ ions in solution. You need to look up K3, Kb, and Kw values.) 443 E i 1 1 ‘ Chapter 11: Acids and Bases 11.73 When lemon juice is squirted into tea, the color becomes lighter. In part, the color change is due to dilution, but the main reason for the change is an acid—base reaction. What is the reaction? (Hint: Tea contains “polyphenols,” which are weak acids, and lemon juice contains citric acid.) 11.74 One of the most common antibiotics is penicillin G (benzylpenicillinic acid), which has the following structure: , \C/OH ' H \/ O / CH3 /c\N—C/ H \c I I I CHa/ \S/(f—(ID—N—(If—CmO H H o It is a weak monoprotic acid: HP T—‘ H+ +P‘ Ka = 1.64 x10—3 where HP denotes the parent acid and P‘ the conjugate base. Penicillin G is produced by growing molds in fermentation tanks at 25 CC and a pH range of 4.5 to 5.0. The crude form of this antibiotic is obtained by extracting the fermentation broth with an organic solvent in which the acid is soluble. (a) Identify the acidic hydrogen atom. (b) In one stage of purification, the organic extract of the crude penicillin G is treated with a bufler solution at pH = 6.50. What is the ratio of the conjugate base of penicillin G to the acid at this pH? Would you expect the conjugate base to be more soluble in water than the acid? (c) Penicillin G is not suitable for oral administration, but the sodium salt (NaP) is because it is soluble. Calculate the pH of a 0.12 M NaP solution formed when a tablet containing the salt is dissolved in a glass of water. all. guawaw Mm” . I M ...
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Chapter 11 Problems - E i i E i i r 1 Problems “Pepsin...

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