Chapter 17

Calculate the change in pH that results from adding (a) \(0.100 \mathrm{mol} \mathrm{NaNO}_{2}\) to \(1.00 \mathrm{L}\) of \(0.100 \mathrm{M} \mathrm{HNO}_{2}(\mathrm{aq})\) (b) \(0.100 \mathrm{mol} \mathrm{NaNO}_{3}\) to \(1.00 \mathrm{L}\) of \(0.100 \mathrm{M} \mathrm{HNO}_{3}(\mathrm{aq})\) Why are the changes not the same? Explain.

Calculate \(\left[\mathrm{OH}^{-}\right]\) in a solution that is (a) \(0.0062 \mathrm{M}\) \(\mathrm{Ba}(\mathrm{OH})_{2}\) and \(0.0105 \mathrm{M} \mathrm{BaCl}_{2} ;\) (b) \(0.315 \mathrm{M}\left(\mathrm{NH}_{4}\right)_{2} \mathrm{SO}_{4}\) and \(0.486 \mathrm{M} \mathrm{NH}_{3} ;\) (c) \(0.196 \mathrm{M} \mathrm{NaOH}\) and \(0.264 \mathrm{M}\) \(\mathrm{NH}_{4} \mathrm{Cl}\)

What concentration of ammonia, \(\left[\mathrm{NH}_{3}\right],\) should be present in a solution with \(\left[\mathrm{NH}_{4}^{+}\right]=0.732 \mathrm{M}\) to produce a buffer solution with \(\mathrm{pH}=9.12 ?\) For \(\mathrm{NH}_{3}\) \(K_{\mathrm{h}}=1.8 \times 10^{-5}\)

Calculate the \(\mathrm{pH}\) of a buffer that is (a) \(0.012 \mathrm{M} \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\left(K_{\mathrm{a}}=6.3 \times 10^{-5}\right)\) and 0.033 \(\mathrm{M} \mathrm{NaC}_{6} \mathrm{H}_{5} \mathrm{COO}\) (b) \(0.408 \mathrm{M} \mathrm{NH}_{3}\) and \(0.153 \mathrm{M} \mathrm{NH}_{4} \mathrm{Cl}\)

Lactic acid, \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{COOH},\) is found in sour milk. A solution containing \(1.00 \mathrm{g} \mathrm{NaCH}_{3} \mathrm{CH}_{2} \mathrm{COO}\) in 100.0 \(\mathrm{mL}\) of \(0.0500 \mathrm{M} \mathrm{HC}_{3} \mathrm{H}_{5} \mathrm{O}_{3}\) has a \(\mathrm{pH}=4.11 .\) What is \(K_{\mathrm{a}}\) of lactic acid?

Indicate which of the following aqueous solutions are buffer solutions, and explain your reasoning. (a) \(0.100 \mathrm{M} \mathrm{NaCl}\) (b) \(0.100 \mathrm{M} \mathrm{NaCl}-0.100 \mathrm{M} \mathrm{NH}_{4} \mathrm{Cl}\) (c) \(0.100 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{2}-0.150 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{3}^{+} \mathrm{Cl}^{-}\) (d) \(0.100 \mathrm{M} \mathrm{HCl}-0.050 \mathrm{M} \mathrm{NaNO}_{2}\) (e) \(0.100 \mathrm{M} \mathrm{HCl}-0.200 \mathrm{M} \mathrm{NaCH}_{3} \mathrm{COO}\) (f) \(0.100 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH}-0.125 \mathrm{M} \mathrm{NaCH}_{3} \mathrm{CH}_{2} \mathrm{COO}\)

The \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}-\mathrm{HPO}_{4}^{2-}\) combination plays a role in maintaining the pH of blood. (a) Write equations to show how a solution containing these ions functions as a buffer. (b) Verify that this buffer is most effective at \(\mathrm{pH} 7.2\) (c) Calculate the \(\mathrm{pH}\) of a buffer solution in which \(\left[\mathrm{H}_{2} \mathrm{PO}_{4}\right]=0.050 \mathrm{M}\) and \(\left[\mathrm{HPO}_{4}^{2-}\right]=0.150 \mathrm{M} .[\)Hint: Focus on the second step of the phosphoric acid ionization.]

What is the \(\mathrm{pH}\) of a solution obtained by adding \(1.15 \mathrm{mg}\) of aniline hydrochloride \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{3}^{+} \mathrm{Cl}^{-}\right)\) to \(3.18 \mathrm{L}\) of \(0.105 \mathrm{M}\) aniline \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\right) ?\)

What is the pH of a solution prepared by dissolving \(8.50 \mathrm{g}\) of aniline hydrochloride \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{3}^{+} \mathrm{Cl}^{-}\right)\) in \(750 \mathrm{mL}\) of \(0.215 \mathrm{M}\) aniline, \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\right) ?\) Would this solution be an effective buffer? Explain.

You are asked to prepare a buffer solution with a pH of 3.50. The following solutions, all \(0.100 \mathrm{M},\) are available to you: HCOOH, CH \(_{3} \mathrm{COOH}, \mathrm{H}_{3} \mathrm{PO}_{4}, \mathrm{NaCHOO}\) \(\mathrm{NaCH}_{3} \mathrm{COO},\) and \(\mathrm{NaH}_{2} \mathrm{PO}_{4} . \quad\) Describe how you would prepare this buffer solution. [Hint: What volumes of which solutions would you use?

\(\begin{array}{lll}\text { Given } & 1.00 & \mathrm{L}\end{array}\) of a solution that is \(0.100 \mathrm{M}\) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{COOH}\) and \(0.100 \mathrm{M} \mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{COO}\) (a) Over what pH range will this solution be an effective buffer? (b) What is the buffer capacity of the solution? That is, how many millimoles of strong acid or strong base can be added to the solution before any significant change in pH occurs?

Given \(125 \mathrm{mL}\) of a solution that is \(0.0500 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{2}\) and \(0.0500 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{3}^{+} \mathrm{Cl}^{-}\) (a) Over what pH range will this solution be an effective buffer? (b) What is the buffer capacity of the solution? That is, how many millimoles of strong acid or strong base can be added to the solution before any significant change in pH occurs?

A handbook lists various procedures for preparing buffer solutions. To obtain a \(\mathrm{pH}=9.00,\) the handbook says to mix \(36.00 \mathrm{mL}\) of \(0.200 \mathrm{M} \mathrm{NH}_{3}\) with \(64.00 \mathrm{mL}\) of \(0.200 \mathrm{M} \mathrm{NH}_{4} \mathrm{Cl}\) (a) Show by calculation that the pH of this solution is 9.00. (b) Would you expect the \(\mathrm{pH}\) of this solution to remain at \(\mathrm{pH}=9.00\) if the \(100.00 \mathrm{mL}\) of buffer solution were diluted to 1.00 L? To 1000 L? Explain. (c) What will be the pH of the original \(100.00 \mathrm{mL}\) of buffer solution if \(0.20 \mathrm{mL}\) of \(1.00 \mathrm{M} \mathrm{HCl}\) is added to it? (d) What is the maximum volume of \(1.00 \mathrm{M} \mathrm{HCl}\) that can be added to \(100.00 \mathrm{mL}\) of the original buffer solution so that the pH does not drop below \(8.90 ?\)

An acetic acid-sodium acetate buffer can be prepared by the reaction \(\mathrm{CH}_{3} \mathrm{COO}^{-}+\mathrm{H}_{3} \mathrm{O}^{+} \longrightarrow \mathrm{CH}_{3} \mathrm{COOH}+\mathrm{H}_{2} \mathrm{O}\) (From \(\mathrm{NaCH}_{3} \mathrm{COO}\) )(From HCl) (a) If \(12.0 \mathrm{g} \mathrm{NaCH}_{3} \mathrm{COO}\) is added to \(0.300 \mathrm{L}\) of 0.200 M HCl, what is the pH of the resulting solution? (b) If \(1.00 \mathrm{g} \mathrm{Ba}(\mathrm{OH})_{2}\) is added to the solution in part (a), what is the new pH? (c) What is the maximum mass of \(\mathrm{Ba}(\mathrm{OH})_{2}\) that can be neutralized by the buffer solution of part (a)? (d) What is the pH of the solution in part (a) following the addition of \(5.50 \mathrm{g} \mathrm{Ba}(\mathrm{OH})_{2} ?\)

In the use of acid-base indicators, (a) Why is it generally sufficient to use a single indicator in an acid-base titration, but often necessary to use several indicators to establish the approximate pH of a solution? (b) Why must the quantity of indicator used in a titration be kept as small as possible?

The indicator methyl red has a \(\mathrm{pK}_{\mathrm{HIn}}=4.95 . \mathrm{It}\) changes from red to yellow over the pH range from 4.4 to 6.2 (a) If the indicator is placed in a buffer solution of \(\mathrm{pH}=4.55,\) what percent of the indicator will be pre- sent in the acid form, HIn, and what percent will be present in the base or anion form, In \(^{-2}\) (b) Which form of the indicator has the "stronger" (that is, more visible) color- -the acid (red) form or base (yellow) form? Explain.

Phenol red indicator changes from yellow to red in the pH range from 6.6 to \(8.0 .\) Without making detailed calculations, state what color the indicator will assume in each of the following solutions: (a) \(0.10 \mathrm{M} \mathrm{KOH}\) (b) \(0.10 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH} ;\) (c) \(0.10 \mathrm{M} \mathrm{NH}_{4} \mathrm{NO}_{3} ;\) (d) \(0.10 \mathrm{M}\) HBr; (e) \(0.10 \mathrm{M} \mathrm{NaCN} ;\) (f) \(0.10 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH}-0.10 \mathrm{M}\) \(\mathrm{NaCH}_{3} \mathrm{COO}\).

Thymol blue indicator has \(t w o\) pH ranges. It changes color from red to yellow in the pH range from 1.2 to 2.8, and from yellow to blue in the pH range from 8.0 to 9.6. What is the color of the indicator in each of the following situations? (a) The indicator is placed in \(350.0 \mathrm{mL}\) of \(0.205 \mathrm{M} \mathrm{HCl}\) (b) To the solution in part (a) is added \(250.0 \mathrm{mL}\) of \(0.500 \mathrm{M} \mathrm{NaNO}_{2}\) (c) To the solution in part (b) is added \(150.0 \mathrm{mL}\) of \(0.100 \mathrm{M} \mathrm{NaOH}\) (d) To the solution in part (c) is added \(5.00 \mathrm{g} \mathrm{Ba}(\mathrm{OH})_{2}\)

In the titration of \(10.00 \mathrm{mL}\) of \(0.04050 \mathrm{M} \mathrm{HCl}\) with \(0.01120 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\) in the presence of the indicator 2,4-dinitrophenol, the solution changes from colorless to yellow when 17.90 mL of the base has been added. What is the approximate value of \(\mathrm{p} K_{\mathrm{HIn}}\) for 2,4 -dinitrophenol? Is this a good indicator for the titration?

Solution (a) is \(100.0 \mathrm{mL}\) of \(0.100 \mathrm{M} \mathrm{HCl}\) and solution (b) is \(150.0 \mathrm{mL}\) of \(0.100 \mathrm{M} \mathrm{NaCH}_{3} \mathrm{COO}\). A few drops of thymol blue indicator are added to each solution. What is the color of each solution? What is the color of the solution obtained when these two solutions are mixed?

A \(25.00 \mathrm{mL}\) sample of \(\mathrm{H}_{3} \mathrm{PO}_{4}(\text { aq) requires } 31.15 \mathrm{mL}\) of \(0.2420 \mathrm{M}\) KOH for titration to the second equivalence point. What is the molarity of the \(\mathrm{H}_{3} \mathrm{PO}_{4}(\mathrm{aq}) ?\)

A 20.00 mL sample of \(\mathrm{H}_{3} \mathrm{PO}_{4}(\mathrm{aq})\) requires \(18.67 \mathrm{mL}\) of \(0.1885 \mathrm{M} \mathrm{NaOH}\) for titration from the first to the second equivalence point. What is the molarity of the \(\mathrm{H}_{3} \mathrm{PO}_{4}(\mathrm{aq}) ?\)

Two aqueous solutions are mixed: \(50.0 \mathrm{mL}\) of 0.0150 \(\mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4}\) and \(50.0 \mathrm{mL}\) of \(0.0385 \mathrm{M} \mathrm{NaOH} .\) What is the pH of the resulting solution?

Two solutions are mixed: \(100.0 \mathrm{mL}\) of \(\mathrm{HCl}(\mathrm{aq})\) with \(\mathrm{pH} 2.50\) and \(100.0 \mathrm{mL}\) of \(\mathrm{NaOH}(\text { aq) with } \mathrm{pH} 11.00\) What is the pH of the resulting solution?

Calculate the \(\mathrm{pH}\) at the points in the titration of \(25.00 \mathrm{mL}\) of 0.160 M HCl when (a) 10.00 mL and (b) 15.00 mL of 0.242 M KOH have been added.

Calculate the pH at the points in the titration of \(25.00 \mathrm{mL}\) of \(0.160 \mathrm{M} \mathrm{HCl}\) when (a) \(10.00 \mathrm{mL}\) and \((\mathrm{b}) 15.00 \mathrm{mL}\) of 0.242 M KOH have been added.

Calculate the \(\mathrm{pH}\) at the points in the titration of \(25.00 \mathrm{mL}\) of \(0.132 \mathrm{M} \mathrm{HNO}_{2}\) when (a) \(10.00 \mathrm{mL}\) and (b) 20.00 mL of 0.116 M \(\mathrm{HNO}_{2}\) when \((\mathrm{a}) 10.00 \mathrm{mL}\) and (b) \(20.00 \mathrm{mL}\) of \(0.116 \mathrm{M}\) NaOH have been added. For \(\mathrm{HNO}_{2}, K_{\mathrm{a}}=7.2 \times 10^{-4}\) \(\mathrm{HNO}_{2}+\mathrm{OH}^{-} \longrightarrow \mathrm{H}_{2} \mathrm{O}+\mathrm{NO}_{2}^{-}\)

Explain why the volume of \(0.100 \mathrm{M} \mathrm{NaOH}\) required to reach the equivalence point in the titration of \(25.00 \mathrm{mL}\) of \(0.100 \mathrm{M}\) HA is the same regardless of whether HA is a strong or a weak acid, yet the \(\mathrm{pH}\) at the equivalence point is not the same.

Explain whether the equivalence point of each of the following titrations should be below, above, or at pH 7: (a) \(\mathrm{NaHCO}_{3}(\text { aq) titrated with } \mathrm{NaOH}(\mathrm{aq}) ; \text { (b) } \mathrm{HCl}(\mathrm{aq})\) titrated with \(\mathrm{NH}_{3}(\mathrm{aq}) ;\) (c) KOH(aq) titrated with HI(aq).

Determine the following characteristics of the titration curve for \(20.0 \mathrm{mL}\) of \(0.275 \mathrm{M} \mathrm{NH}_{3}(\mathrm{aq})\) titrated with \(0.325 \mathrm{M} \mathrm{HI}(\mathrm{aq})\) (a) the initial \(\mathrm{pH}\) (b) the volume of \(0.325 \mathrm{M} \mathrm{HI}(\mathrm{aq})\) at the equivalence point (c) the \(\mathrm{pH}\) at the half-neutralization point (d) the \(\mathrm{pH}\) at the equivalence point

In the titration of \(20.00 \mathrm{mL}\) of \(0.175 \mathrm{M} \mathrm{NaOH},\) calculate the number of milliliters of \(0.200 \mathrm{M} \mathrm{HCl}\) that must be added to reach a pH of (a) \(12.55,\) (b) \(10.80,\) (c) 4.25

In the titration of \(25.00 \mathrm{mL}\) of \(0.100 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH}\) calculate the number of milliliters of \(0.200 \mathrm{M} \mathrm{NaOH}\) that must be added to reach a pH of (a) \(3.85,\) (b) 5.25 (c) 11.10.

For the titration of \(25.00 \mathrm{mL}\) of \(0.100 \mathrm{M} \mathrm{NaOH}\) with \(0.100 \mathrm{M} \mathrm{HCl},\) calculate the \(\mathrm{pOH}\) at a few representative points in the titration, sketch the titration curve of pOH versus volume of titrant, and show that it has exactly the same form as Figure \(17-9 .\) Then, using this curve and the simplest method possible, sketch the titration curve of pH versus volume of titrant.

Is a solution that is \(0.10 \mathrm{M} \mathrm{Na}_{2} \mathrm{S}(\mathrm{aq})\) likely to be acidic, basic, or pH neutral? Explain.

Is a solution of sodium dihydrogen citrate, \(\mathrm{NaH}_{2} \mathrm{Cit}\) likely to be acidic, basic, or neutral? Explain. Citric \(\mathrm{acid}, \mathrm{H}_{3} \mathrm{Cit}, \mathrm{is} \mathrm{H}_{3} \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{O}_{7}\)

Sodium phosphate, \(\mathrm{Na}_{3} \mathrm{PO}_{4},\) is made commercially by first neutralizing phosphoric acid with sodium carbonate to obtain \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\). The \(\mathrm{Na}_{2} \mathrm{HPO}_{4}\) is further neutralized to \(\mathrm{Na}_{3} \mathrm{PO}_{4}\) with \(\mathrm{NaOH}\) (a) Write net ionic equations for these reactions. (b) \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) is a much cheaper base than is \(\mathrm{NaOH}\) Why do you suppose that \(\mathrm{NaOH}\) must be used as well as \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) to produce \(\mathrm{Na}_{3} \mathrm{PO}_{4} ?\)

Both sodium hydrogen carbonate (sodium bicarbonate) and sodium hydroxide can be used to neutralize acid spills. What is the pH of \(1.00 \mathrm{M} \mathrm{NaHCO}_{3}(\mathrm{aq})\) and of \(1.00 \mathrm{M} \mathrm{NaOH}(\mathrm{aq}) ?\) On a per-liter basis, do these two solutions have an equal capacity to neutralize acids? Explain. On a per-gram basis, do the two solids, \(\mathrm{NaHCO}_{3}(\mathrm{s})\) and \(\mathrm{NaOH}(\mathrm{s}),\) have an equal capacity to neutralize acids? Explain. Why do you suppose that \(\mathrm{NaHCO}_{3}\) is often preferred to \(\mathrm{NaOH}\) in neutralizing acid spills?

The pH of a solution of \(19.5 \mathrm{g}\) of malonic acid in \(0.250 \mathrm{L}\) is \(1.47 .\) The pH of a \(0.300 \mathrm{M}\) solution of sodium hydrogen malonate is 4.26. What are the values of \(K_{\mathrm{a}_{1}}\) and \(K_{\mathrm{a}_{2}}\) for malonic acid?

The ionization constants of ortho-phthalic acid are \(K_{\mathrm{a}_{1}}=1.1 \times 10^{-3}\) and \(K_{\mathrm{a}_{2}}=3.9 \times 10^{-6}\) 1. \(\mathrm{C}_{6} \mathrm{H}_{4}(\mathrm{COOH})_{2}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}+\mathrm{HC}_{8} \mathrm{H}_{4} \mathrm{O}_{4}^{-}\) 2\. \(\mathrm{HC}_{8} \mathrm{H}_{4} \mathrm{O}_{4}^{-}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}+\mathrm{C}_{6} \mathrm{H}_{4}\left(\mathrm{COO}^{-}\right)_{2}\) What are the pH values of the following aqueous solutions: (a) 0.350 M potassium hydrogen orthophthalate; (b) a solution containing 36.35 g potassium ortho-phthalate per liter? (f) \(0.68 \mathrm{M} \mathrm{KCl}, 0.42 \mathrm{M} \mathrm{KNO}_{3}, 1.2 \mathrm{M} \mathrm{NaCl},\) and \(0.55 \mathrm{M}\) \(\mathrm{NaCH}_{3} \mathrm{COO},\) with \(\mathrm{pH}=6.4\)

What stoichiometric concentration of the indicated substance is required to obtain an aqueous solution with the pH value shown: (a) \(\mathrm{Ba}(\mathrm{OH})_{2}\) for \(\mathrm{pH}=11.88 ;(\mathrm{b})\) \(\mathrm{CH}_{3} \mathrm{COOH}\) in \(0.294 \mathrm{M} \mathrm{NaCH}_{3} \mathrm{COO}\) for \(\mathrm{pH}=4.52 ?\)

What stoichiometric concentration of the indicated substance is required to obtain an aqueous solution with the pH value shown: (a) aniline, \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\), for \(\mathrm{pH}=8.95 ;(\mathrm{b}) \mathrm{NH}_{4} \mathrm{Cl}\) for \(\mathrm{pH}=5.12 ?\)

Using appropriate equilibrium constants but without doing detailed calculations, determine whether a solution can be simultaneously: (a) \(0.10 \mathrm{M} \mathrm{NH}_{3}\) and \(0.10 \mathrm{M} \mathrm{NH}_{4} \mathrm{Cl},\) with \(\mathrm{pH}=6.07\) (b) \(0.10 \mathrm{M} \mathrm{NaC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) and \(0.058 \mathrm{M} \mathrm{HI}\) (c) \(0.10 \mathrm{M} \mathrm{KNO}_{2}\) and \(0.25 \mathrm{M} \mathrm{KNO}_{3}\) (d) \(0.050 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\) and \(0.65 \mathrm{M} \mathrm{NH}_{4} \mathrm{Cl}\) (e) \(0.018 \mathrm{M} \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\) and \(0.018 \mathrm{M} \mathrm{NaC}_{6} \mathrm{H}_{5} \mathrm{COO}\) with \(\mathrm{pH}=4.20\) (f) \(0.68 \mathrm{M} \mathrm{KCl}, 0.42 \mathrm{M} \mathrm{KNO}_{3}, 1.2 \mathrm{M} \mathrm{NaCl},\) and \(0.55 \mathrm{M}\) \(\mathrm{NaCH}_{3} \mathrm{COO},\) with \(\mathrm{pH}=6.4\)

This single equilibrium equation applies to different phenomena described in this or the preceding chapter. \(\mathrm{CH}_{3} \mathrm{COOH}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}+\mathrm{CH}_{3} \mathrm{COO}^{-}\) Of these four phenomena, ionization of pure acid, common-ion effect, buffer solution, and hydrolysis, indicate which occurs if (a) \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) and \(\left[\mathrm{CH}_{3} \mathrm{COOH}\right]\) are high, but \(\left[\mathrm{CH}_{3} \mathrm{COO}^{-}\right]\) is very low. (b) \(\left[\mathrm{CH}_{3} \mathrm{COO}^{-}\right]\) is high, but \(\left[\mathrm{CH}_{3} \mathrm{COOH}\right]\) and \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) are very low. (c) \(\left[\mathrm{CH}_{3} \mathrm{COOH}\right]\) is high, but \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) and \(\left[\mathrm{CH}_{3} \mathrm{COO}^{-}\right]\) are low. (d) \(\left[\mathrm{CH}_{3} \mathrm{COOH}\right]\) and \(\left[\mathrm{CH}_{3} \mathrm{COO}^{-}\right]\) are high, but \(\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\) is low.

You are given \(250.0 \mathrm{mL}\) of \(0.100 \mathrm{M} \mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{COOH}\) (propionic acid, \(\left.K_{\mathrm{a}}=1.35 \times 10^{-5}\right) .\) You want to adjust its \(\mathrm{pH}\) by adding an appropriate solution. What volume would you add of (a) \(1.00 \mathrm{M} \mathrm{HCl}\) to lower the \(\mathrm{pH}\) to \(1.00 ;\) (b) \(1.00 \mathrm{M} \mathrm{NaCH}_{3} \mathrm{CH}_{2} \mathrm{COO}\) to raise the \(\mathrm{pH}\) to \(4.00 ;(\mathrm{c})\) water to raise the \(\mathrm{pH}\) by 0.15 unit?

Even though the carbonic acid-hydrogen carbonate buffer system is crucial to the maintenance of the \(\mathrm{pH}\) of blood, it has no practical use as a laboratory buffer solution. Can you think of a reason(s) for this?

Rather than calculate the \(\mathrm{pH}\) for different volumes of titrant, a titration curve can be established by calculating the volume of titrant required to reach certain \(\mathrm{pH}\) values. Determine the volumes of \(0.100 \mathrm{M} \mathrm{NaOH}\) required to reach the following pH values in the titration of \(20.00 \mathrm{mL}\) of \(0.150 \mathrm{M} \mathrm{HCl}: \mathrm{pH}=\) (a) 2.00 (b) \(3.50 ;\) (c) \(5.00 ;\) (d) \(10.50 ;\) (e) \(12.00 .\) Then plot the titration curve.

You are asked to prepare a \(\mathrm{KH}_{2} \mathrm{PO}_{4}-\mathrm{Na}_{2} \mathrm{HPO}_{4}\) solu- tion that has the same \(\mathrm{pH}\) as human blood, 7.40 (a) What should be the ratio of concentrations \(\left[\mathrm{HPO}_{4}^{2-}\right] /\left[\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\right]\) in this solution? (b) Suppose you have to prepare \(1.00 \mathrm{L}\) of the solution described in part (a) and that this solution must be isotonic with blood (have the same osmotic pressure as blood). What masses of \(\mathrm{KH}_{2} \mathrm{PO}_{4}\) and of \(\mathrm{Na}_{2} \mathrm{HPO}_{4} \cdot 12 \mathrm{H}_{2} \mathrm{O}\) would you use? [Hint: Refer to the definition of isotonic on page \(580 .\) Recall that a solution of \(\mathrm{NaCl}\) with \(9.2 \mathrm{g} \mathrm{NaCl} / \mathrm{L}\) solution is isotonic with blood, and assume that \(\mathrm{NaCl}\) is completely ionized in aqueous solution.]

You are asked to bring the \(\mathrm{pH}\) of \(0.500 \mathrm{L}\) of \(0.500 \mathrm{M}\) \(\mathrm{NH}_{4} \mathrm{Cl}(\mathrm{aq})\) to 7.00 How many drops \((1 \text { drop }=0.05 \mathrm{mL})\) of which of the following solutions would you use: \(10.0 \mathrm{M} \mathrm{HCl}\) or \(10.0 \mathrm{M} \mathrm{NH}_{3} ?\)

Because an acid-base indicator is a weak acid, it can be titrated with a strong base. Suppose you titrate \(25.00 \mathrm{mL}\) of a \(0.0100 \mathrm{M}\) solution of the indicator \(p\) -nitrophenol, \(\mathrm{HOC}_{6} \mathrm{H}_{4} \mathrm{NO}_{2},\) with \(0.0200 \mathrm{M} \mathrm{NaOH}\) The \(\mathrm{p} K_{\mathrm{a}}\) of \(p\) -nitrophenol is \(7.15,\) and it changes from colorless to yellow in the pH range from 5.6 to 7.6 (a) Sketch the titration curve for this titration. (b) Show the pH range over which \(p\) -nitrophenol changes color. (c) Explain why \(p\) -nitrophenol cannot serve as its own indicator in this titration.

The neutralization of \(\mathrm{NaOH}\) by \(\mathrm{HCl}\) is represented in equation (1), and the neutralization of \(\mathrm{NH}_{3}\) by HCl in equation (2). 1. \(\mathrm{OH}^{-}+\mathrm{H}_{3} \mathrm{O}^{+} \rightleftharpoons 2 \mathrm{H}_{2} \mathrm{O} \quad K=?\) 2\. \(\mathrm{NH}_{3}+\mathrm{H}_{3} \mathrm{O}^{+} \rightleftharpoons \mathrm{NH}_{4}^{+}+\mathrm{H}_{2} \mathrm{O} \quad K=?\) (a) Determine the equilibrium constant \(K\) for each reaction. (b) Explain why each neutralization reaction can be considered to go to completion.