QUESTIONS AND PROBLEMS(3)

2020-09-17 来源：步旅网

QUESTIONS AND PROBLEMS

Chapter Ⅲ Acid-Base Equilibrium and Titrations

3-1. Briefly describe or define and give an example of

*(a) a weak electrolyte.

(b) a Brønsted-Lowry acid.

*(c) the conjugate acid of a Brønsted-Lowry base.

(d) neutralilation, in terms of the Brønsted-Lowry concept.

*(e) an amphiprotic solvent.

(f) a zwitterion.

*( g) autoprotolysis.

(h) a strong acid.

*(i) the Le Chatelier principle.

(j) the common-ion effect.

3-2. Briefly describe or define and give an example of

*(a) an amphiprotic solute. (b) a differentiating solvent.

*(c) a leveling solvent. (d) a mass-action effect.

*3-3. Briefly explain why there is no terrn in an equilibrium constant expression for water or for a pure solid, even though one (or both) appears in the balanced net ionic equation for the equilibrium.

3-4. Identify the acid on the left and its conjugate base on the right in the following equations:

*(a) HOCl + H2O  H3O+ + OCl- (b) HONH2 + H2O  HONH3+ + OH-

*(c) NH4+ + H2O  NH3 + H3O+ (d) 2HCO3-  H2CO3 + CO32-

*(e) PO43- + H2PO4-  2HPO42-

3-5. Identify the base on the left and its cunjugate acid on the right in the equations for Problem 3-4.

3-6. Write expressions for the autoprotolysis of

*(a) H2O. (b) CH3COOH. *(c) CH3NH2. (d) CH3OH.

3-7. Write the equilibriurn-constant expressions and obtain numerical value for each constant in

*(a) the basic dissociation of ethylamine, C2H5NH2.

(b) the acidic dissociation of hydrogen cyanide, HCN.

*(c) the acidic dissociation of pyridine hydrochloride, C5H5NHCl.

(d) the basic dissociation of NaCN.

*(e) the dissociation of H3AsO4 to H3O+ and AsO43-.

(f) the reaction of CO32- with H2O. to give H2CO3 and OH-.

3-8. Generate the solubility-product expression for

*(a) CuI. *(b) PbCIF. *(c) PbI2. (d) BiI3 (e) MgNH4PO4.

3-9. Express the solubility-product constant for each substance in Problem 3-8 in terms of its rnolar solubility S.

3-10. Calculate the solubility-product constant for each of the following substances. given that the rnolar concentrations of their saturated solutions are as indicated:

(a) CuSeO3 (1.42×3-4 M). *(b) Pb(IO3)2 (4.3×3-5 M).

3-11. Calculate the solubility of the solutes in Problem 3-10 for solutions in which the cation concentration is 0.050 M.

3-12. Calculate the solubility of the solutes in Problem 3-10 for solutions in which the anion concentration is 0.050 M.

*3-13. What CrO42- concentration is required to

(a) initiate precipitation of Ag2CrO4 from a solution that is 3.41×3-2 M in Ag+?

(b) lower the concentration of Ag+ in a solution to 2.00×3-6 M?

3-14. What hydroxide concentration is required to

(a) initiate precipitation of A13+ from a 2.50×3-2 M solution of Al2(SO4)3?

(b) lower the A13+ concentration in the foregoing solution to 2.00×3-7M?

*3-15. The solubility-product constant for Ce(IO3)3 is 3.2×3-10. What is the Ce3+ concentration in a solution prepared by mixing 50.0 mL of 0.0250 M Ce3+ with 50.00 rnL of

(a) water? (b) 0.040 M IO3-? (c) 0.250 M IO3-? (d) 0.150 M IO3-?

3-16. The solubility-product constant for K2PdCl6 is 6.0×3-6

(K2PdCl6  2K+ + PdCl62-). What is the K+ concentration of a solution prepared by mixing 50.0 mL of 0.200 M KCl with 50.0 mL of

(a) 0.0500 M PdCl62-? (b) 0.100 M PdCl62-? (c) 0.200 M PdCl62-?

*3-17. The solubility products for a series of iodides are

CuI Ksp =1×3-12 AgI Ksp = 8.3×3-17

PbI2 Ksp = 7.1×3-9 BiI3 Ksp = 8.1×3-19

List these four compounds in order of decreasing molar solubility in

(a) water. (b) 0.10 M NaI (c) a 0.010 M solution of the solute cation.

3-18. The solubility products for a series of hydroxides are

BiOOH Ksp = 4.0×3-10= [BiO+] [OH-]

Be(OH)2 Ksp = 7.0×3-22

Tm(OH)3 Ksp = 3.0×3-24

Hf(OH)4 Ksp = 4.0×3-26

Which hydroxide has

(a) the lowest molar solubility in H2O?

(b) the lowest molar solubility in a solution that is 0.10 M in NaOH?

3-19. Calculate the pH of water at 0°C and 100°C.

3-20. At 25°C, what are the molar H3O+ and OH- concentrations in

*(a) 0.0300 M HOCl? (b) 0.0600 M butanoic acid?

*(c) 0.100 Methylamine? (d) 0.200 M trimethylamine?

*(e) 0.200 M NaOCl? (f) 0.0860 M CH3CH2COONa?

*(g) 0.250 M hydroxylamine hydrochloride? (h) 0.0250 M ethanolamine hydrochloride?

3-21. At 25°C, what is the hydronium ion concentration in

*(a) 0.100 M chloroacetic acid? *(b) 0.100 M sodium chloroacetate?

*(e) 1.00×3-3 M aniline hydrochloride? (f) 0.200 M HIO3 ?

3-22. What is a buffer solution, and what are its properties?

*3-23. Define buffer capacity.

3-24. Which has the greater buffer capacity: (a) a mixture containing 0.100 mol of NH3 and 0.200 mol of NH4Cl or (b) a rnixture containing 0.0500 mol of NH3 and 0.100 mol of NH4Cl?

*3-25. Consider solutions prepared by

(a) dissolving 8.00 mmol of NaOAc in 200 mL of 0.100 M HOAc.

(b) adding 100 rnL of 0.0500 M NaOH to 100 rnL of 0.175 M HOAc.

In what respects do these solutions resemble one another? How do they differ?

3-26. Consult Appendix 3 and pick out a suitable acid/base pair to prepare a buffer with a pH of

*(a) 3.5, (b) 7.6, *(c) 9.3, (d) 5.1.

*3-27. What weight of sodium formate must be added to 400.0 mL of 1.00 M formic acid to produce a buffer solution that has a pH of 3.50?

3-28. What weight of sodium glycolate should be added to 300.0 mL of 1.00 M glycolic acid to produce a buffer solution with a pH of 4.00?

*3-29. What volume of 0.200 M HCI must be added to 250.0 mL of 0.300 M sodium mandelate to produce a buffer solution with a pH of 3.37?

3-30. What volume of 2.00 M NaOH rnust be added to 300.0 mL of 1.00 M glycolic acid to produce a buffer solution having a pH of 4.00?

3-31. Is the following statement true or false. or both? Define your answer with equations. examples, or graphs. “A butfer rnaintains the pH of a solution constant.”

3-32. Challenge Problem: It can be shown1 that the buffer capacity is

+cKHOKTa3+w2.303H3O2++H3OKaH3O

where cT is the molar analytical concentration of the buffer.

(a) Show that

+2.303OHHO3cT01

(b) Use the equation in (a) to explain the shape of Figure 3-6.

(c) Differentiate the equation presented at the beginning of the problern and show that the butfer capacity is at a maximum whenα0 =α1 = 0.5.

(d) Describe the conditions under which these relationships apply.

*3-1. Make a distinction between

(a) activity and activity coefficient.

(b) thermodynamic and concentration equilibrium constants.

3-2. List general properties of activity coefficients.

*3-3. Neglecting any effects caused by volume changes, would you expect the ionic strength to (1) increase, (2) decrease, or (3) remain essentially unchanged by the addition of NaOH to a dilute solution of

(a) magnesium chloride [Mg(OH)2 (s) forms]?

(b) hydrochloric acid? (c) acetic acid?

3-4. Neglecting any effects caused by volume changes, would you expect the

ionic strength to (1) increase, (2) decrease. or (3) remain essentially unchanged by the addition of iron (III) chloride to

(a) HCl? (b) NaOH? (c) AgNO3?

*3-5. Explain why the initial slope for Ca2+ in Figure 3-3 is steeper than that for K+?

3-6. What is the numerical value of the activity coefficient of aqueous ammonia (NH3) at an ionic strength of 0.1 ?

3-7. Calculate the ionic strength of a solution that is

*(a) 0.040 M in FeSO4. (b) 0.20 M in (NH4)2CrO4.

*( c) 0.10 M in FeCl3 and 0.20 M in FeCl2.

(d) 0.060 M in La(NO3)3 and 0.030 M in Fe(NO3)2.

3-8. Use Equation 3-5 to calculate the activity coefficient of

*(a) Fe3+ at μ= 0.075. (b) Pb2+ at μ = 0.012.

*(c) Ce4+ at μ = 0.080. (d) Sn4+ at μ = 0.060.

3-9. Calculate activity coefficients for the species in Problem 3-8 by linear

interpolation of the data in Table 3-2.

3-10. For a solution in whichμis 5.0×10-2 . calculate K'sp for

* (a) AgSCN. (b) PbI2. *(c) La(IO3)3. (d) MgNH4PO4.

*3-11. Use activities to calculate the molar solubility of Zn(OH)2 in

(a) 0.0100 M KCl. (b) 0.0167 M K2SO4 .

(d) the solution that results when you mix 20.0 mL of 0.100 M KOH with 80.0 mL of 0.0250 M ZnCl2.

*3-12. Calculate the solubilities of the following cornpounds in a 0.0333 M solution of Mg(ClO4)2 using (1) activities and (2) molar concentrations:

(a) AgSCN. (b) PbI2 . (c) BaSO4. (d) Cd2Fe(CN)6.

Cd2Fe(CN)6(s)  2Cd2+ + Fe(CN)64- Ksp = 3.2×10-17

*3-13. Calculate the solubilities of the following compounds in a 0.0167 M solution of Ba(NO3)2 using (I) activities and (2) molar concentrations:

(a) AgIO3. (b) Mg(OH)2. (c) BaSO4. (d) La(IO3)3.

*3-14. Calculate the % relative error in solubility by using concentrations instead of activities for the following compounds in 0.05000 M KNO3. using the thermodynamic solubility products listed in Appendix 2.

*(a) CuCl (acu' = 0.3 nrn). (b) Fe(OH)2. *(c) Fe(OH)3. (d) La(IO3)3.

*(e) Ag3AsO4 (aAsO4 = 0.4 nm).

3-15. Calculate the % relative error in hydronium ion concentration by using concentrations instead of activities in calculating the pH of solution of the following species using the thermodynamic constants found in Appendix 3.

*(a) 0.100 M HOAc and 0.200 M NaOAc.

(b) 0.0500 M NH3 and 0.200 M NH4Cl.

13-16. (a) Repeat the computations of Problem 3-15 using a spreadsheet. Vary the concentration of Ba(NO3)2 from 0.0001 M to 1 M in a manner similar to that used in the spreadsheet exercise.

(b) Plot pS versus pc, where pc is the negative logalithm of the concentration of Ba(NO3)2.

3-17. Design and construct a spreadsheet to calculate activity coefficients in a format similar to Table 3-2. Enter values of ax in cells A3, A4, A5, and so forth. and enter ionic charges in cells B3, B4, B5, and so forth. Enter in cells C2:G2 the same set of values for ionic strength listed in Table 3-2. Enter the formula for the activity coefficients in cells C3:G3. Be sure to use absolute cell references for ionic strength in your formulas for the activity coefficients. Finally, copy the formulas for the activity coefficients into the rows below row C by highlighting C3:G3 and dragging the fill handle downward. Compare the activity coefficients that you calculate to those in Table 3-2. Do you find any discrepancies? If so, explain how they arise.

13-18. Challenge Problem. In example 3-5. we neglected the contribution of nitrous acid to the ionic strength. We also used the simplified solution for the hydronium ion concentration,

HO3Kaca (a) Carry out an iterative solution to the problem in which you actually calculate the ionic strength. first without taking into account the dissociation of the acid. Then calculate corresponding activity coefficients for the ions using the Debye-Huckel equation, compute a new Ka, and find a new value for [H3O+]. Repeat the process. but use the concentrations of [H3O+] and [NO2-] along with the 0.05 M NaCl to calculate a new ionic strength; once again. find the activity coefficients. Ka, and a new value for [H3O+]. Iterate until you obtain two consecutive values of [H3O+] that are equal to Within 0.1 %. How many iterations did you need? What is the relative error between your final value and the value

obtained in Example 3-5 with no activity correction? What is the relative error between the first value that you calculated and the last one? You may want to use a spreadsheet to assist you in these calculations.

(b) Now perform the same calculation, except this time calculate the hydronium ion concentration using the quadratic equation or the method of successive approximations each time you compute a new ionic strength. How much irnprovement do you observe over the results that you obtained in (a)?

(c) When are activity corrections like those that you carried out in (a) necessary? What variables rnust be considered in deciding whether to make such corrections?

(d) When are corrections such as those in (b) necessary? What criteria do you use to decide whether these corrections should be made?

(e) Suppose that you are attempting to determine ion concentrations in a complex matrix such as blood serum or urine. Is it possible to make activity corrections in such a system? Explain your answer.

lJ. N. Buller. Ionic Equilibrium: A Mathematical Approach. p. 151. Menlo Park. CA: Addjson-Wele). 1964.

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QUESTIONS AND PROBLEMS(3)