Introduction to General, Organic and Biochemistry

Problem 4

For each of the following, tell whether the acid is strong or weak.
(a) Acetic acid
(b) $\mathrm{HCl}$
(c) $\mathrm{H}_{3} \mathrm{PO}_{4}$
(d) $\mathrm{H}_{2} \mathrm{SO}_{4}$
(e) HCN
(f) $\mathrm{H}_{2} \mathrm{CO}_{3}$

David Collins

Numerade Educator

00:58

Problem 5

For each of the following, tell whether the base is strong or weak.
(a) $\mathrm{NaOH}$
(b) Sodium acetate
(c) KOH
(d) Ammonia
(e) Water

David Collins

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02:11

Problem 6

Which of these acids are monoprotic, which are diprotic, and which are triprotic? Which are amphiprotic?
(a) $\mathrm{H}_{2} \mathrm{PO}_{4}^{-}$
(b) $\mathrm{HBO}_{3}^{2-}$
(c) $\mathrm{HClO}_{4}$
(d) $\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}$
(e) $\mathrm{HSO}_{3}^{-}$
(f) HS"
$(\mathrm{g}) \mathrm{H}_{2} \mathrm{CO}_{3}$

David Collins

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00:10

Problem 7

Define (a) a Bronsted-Lowry acid and (b) a Bronsted-Lowry base.

David Collins

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00:45

Problem 8

Write the formula for the conjugate base of each acid.
(a) $\mathrm{H}_{2} \mathrm{SO}_{4}$
(b) $\mathrm{H}_{3} \mathrm{BO}_{3}$
(c) HI
(d) $\mathrm{H}_{3} \mathrm{O}^{+}$
(e) $\mathrm{NH}_{4}^{+}$
(f) $\mathrm{HPO}_{4}^{2-}$

David Collins

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00:28

Problem 9

Write the formula for the conjugate base of each acid.
(a) $\mathrm{H}_{2} \mathrm{PO}_{4}^{-}$
(b) $\mathrm{H}_{2} \mathrm{S}$
(c) $\mathrm{HCO}_{3}^{-}$
(d) $\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}$
(e) $\mathrm{H}_{2} \mathrm{O}$

David Collins

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Problem 10

Write the formula for the conjugate acid of each base.
(a) $\mathrm{OH}^{-}$
(b) $\mathrm{HS}^{-}$
(c) $\mathrm{NH}_{3}$
(d) $\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{O}^{-}$
(e) $\mathrm{CO}_{3}^{2-}$
(f) $\mathrm{HCO}_{3}^{-}$

David Collins

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00:33

Problem 11

Write the formula for the conjugate acid of each base.
(a) $\mathrm{H}_{2} \mathrm{O}$
(b) $\mathrm{HPO}_{4}^{2-}$
(c) $\mathrm{CH}_{3} \mathrm{NH}_{2}$
(d) $\mathrm{PO}_{4}^{3-}$

David Collins

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00:35

Problem 12

Show how the amphiprotic ion hydrogen carbonate, $\mathrm{HCO}_{3}^{-},$ can react as both an acid and a base.

David Collins

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00:24

Problem 13

Draw the acid and base reactions for the amphiprotic ion $\mathrm{HPO}_{3}^{2-}$.

David Collins

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01:15

Problem 14

For each equilibrium, label the stronger acid, stronger base, weaker acid, and weaker base. For which reaction(s) does the position of equilibrium lie toward the right? For which does it lie toward the left?
(a) $\mathrm{H}_{3} \mathrm{PO}_{4}+\mathrm{OH}^{-} \rightleftharpoons \mathrm{H}_{2} \mathrm{PO}_{4}^{-}+\mathrm{H}_{2} \mathrm{O}$
(b) $\mathrm{H}_{2} \mathrm{O}+\mathrm{Cl}^{-} \rightleftharpoons \mathrm{HCl}+\mathrm{OH}^{-}$
(c) $\quad \mathrm{HCO}_{3}^{-}+\mathrm{OH}^{-} \rightleftharpoons \mathrm{CO}_{3}^{2-}+\mathrm{H}_{2} \mathrm{O}$

David Collins

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01:27

Problem 15

For each equilibrium, label the stronger acid, stronger base, weaker acid, and weaker base. For which reaction(s) does the position of equilibrium lie toward the right? For which does it lie toward the left?
(a) $\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{OH}+\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{O}^{-} \rightleftharpoons \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{O}^{-}+\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}$
(b) $\mathrm{HCO}_{3}^{-}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{H}_{2} \mathrm{CO}_{3}+\mathrm{OH}^{-}$
(c) $\mathrm{CH}_{3} \mathrm{COOH}+\mathrm{H}_{2} \mathrm{PO}_{4}^{-} \rightleftharpoons \mathrm{CH}_{3} \mathrm{COO}^{-}+\mathrm{H}_{3} \mathrm{PO}_{4}$

David Collins

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01:45

Problem 16

Will carbon dioxide be evolved as a gas when sodium bicarbonate is added to an aqueous solution of each compound? Explain.
(a) Sulfuric acid
(b) Ethanol, $\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}$
(c) Ammonium chloride, $\mathrm{NH}_{4} \mathrm{Cl}$

David Collins

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00:29

Problem 17

Which has the larger numerical value?
(a) The $\mathrm{p} K_{\mathrm{a}}$ of a strong acid or the $\mathrm{p} K_{\mathrm{a}}$ of a weak acid
(b) The $K_{\mathrm{a}}$ of a strong acid or the $K_{\mathrm{a}}$ of a weak acid

David Collins

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01:21

Problem 18

In each pair, select the stronger acid.
(a) Pyruvic acid $\left(\mathrm{p} K_{\mathrm{a}}=2.49\right)$ or lactic acid $\left(p K_{s}=3.08\right)$
(b) Citric acid $\left(\mathrm{p} K_{\mathrm{a}}=3.08\right)$ or phosphoric acid $\left(p K_{a}=2.10\right)$
(c) Benzoic acid $\left(K_{\mathrm{a}}=6.5 \times 10^{-5}\right)$ or lactic acid $\left(K_{\mathrm{a}}=8.4 \times 10^{-4}\right)$
(d) Carbonic acid $\left(K_{2}=4.3 \times 10^{-7}\right)$ or boric acid $\left(K_{\mathrm{a}}=7.3 \times 10^{-10}\right)$

David Collins

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01:41

Problem 19

Which solution will be more acidic; that is, which will have a lower pH?
(a) $0.10 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH}$ or $0.10 \mathrm{M} \mathrm{HCl}$
(b) $0.10 M \mathrm{CH}_{3} \mathrm{COOH}$ or $0.10 \mathrm{M} \mathrm{H}_{3} \mathrm{PO}_{4}$
(c) $0.010 M \mathrm{H}_{2} \mathrm{CO}_{3}$ or $0.010 \mathrm{M} \mathrm{NaHCO}_{3}$
(d) $0.10 M \mathrm{NaH}_{2} \mathrm{PO}_{4}$ or $0.10 \mathrm{M} \mathrm{Na}_{2} \mathrm{HPO}_{4}$
(e) $0.10 M$ aspirin $\left(\mathrm{p} K_{\mathrm{a}}=3.47\right)$ or $0.10 \mathrm{M}$ acetic acid

David Collins

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01:59

Problem 20

Which solution will be more acidic; that is, which will have a lower pH?
(a) $0.10 M \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{OH}$ (phenol) or $0.10 \mathrm{M} \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}$
(ethanol)
(b) $0.10 M \mathrm{NH}_{3}$ or $0.10 \mathrm{M} \mathrm{NH}_{4} \mathrm{Cl}$
(c) $0.10 M \mathrm{NaCl}$ or $0.10 \mathrm{M} \mathrm{NH}_{4} \mathrm{Cl}$
(d) $0.10 M \mathrm{CH}_{3} \mathrm{CH}(\mathrm{OH}) \mathrm{COOH}$ (lactic acid) or $0.10 \mathrm{M}$
$\mathrm{CH}_{3} \mathrm{COOH}$
(e) $\left.0.10 M \text { ascorbic acid (vitamin } C, p K_{n}=4.1\right)$ or $0.10 M$ acetic acid

David Collins

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02:09

Problem 21

Write an equation for the reaction of HCl with each compound. Which are acid-base reactions? Which are redox reactions?
(a) $\mathrm{Na}_{2} \mathrm{CO}_{3}$
(b) $\mathrm{Mg}$
(c) $\mathrm{NaOH}$
(d) $\operatorname{Fe}_{2} \mathrm{O}_{3}$
(e) $\mathrm{NH}_{3}$
$\begin{array}{llll}\text { (f) } & \mathrm{CH}_{3} \mathrm{NH}_{2}(\mathrm{g}) & \mathrm{NaHCO}_{3} & \text { (h) } \mathrm{Al}\end{array}$

David Collins

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01:15

Problem 22

When a solution of sodium hydroxide is added to a solution of ammonium carbonate and then heated, ammonia gas, $\mathrm{NH}_{3}$, is released. Write a net ionic equation for this reaction. Both $\mathrm{NaOH}$ and $\left(\mathrm{NH}_{4}\right)_{2} \mathrm{CO}_{3}$ exist as dissociated ions in aqueous solution.

David Collins

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Problem 23

Given the following values of $\left[\mathrm{H}_{3} \mathrm{O}^{+}\right],$ calculate the cor responding value of $\left[\mathrm{OH}^{-}\right]$ for each solution.
(a) $10^{-11} M$
(b) $10^{-4} M$
(c) $10^{-7} M$
(d) $10 M$

David Collins

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00:47

Problem 24

Given the following values of $\left[\mathrm{OH}^{-}\right]$, calculate the corresponding value of $\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]$ for each solution.
(a) $10^{-10} M$
$\begin{array}{ll}\text { (b) } & 10^{-2} M\end{array}$
(c) $10^{-7} M$
(d) $10 M$

David Collins

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00:33

Problem 25

What is the pH of each solution given the following values of $\left[\mathrm{H}_{3} \mathrm{O}^{+}\right] ?$ Which solutions are acidic, which are basic, and which are neutral?
(a) $10^{-8} M$
(b) $10^{-10} M$
(c) $10^{-2} M$
(d) $10^{\circ} M$
(e) $10^{-7} M$

David Collins

Numerade Educator

Problem 26

What is the $\mathrm{pH}$ and $\mathrm{pOH}$ of each solution given the following values of $\left[\mathrm{OH}^{-}\right] ?$ Which solutions are acidic, which are basic, and which are neutral?
(a) $10^{-3} M$
(b) $10^{-1} M$
(c) $10^{-5} M$
$(d) 10^{-7} M$

David Collins

Numerade Educator

Problem 27

What is the $\mathrm{pH}$ of each solution given the following values of $\left[\mathrm{H}_{3} \mathrm{O}^{+}\right] ?$ Which solutions are acidic, which are basic, and which are neutral?
(a) $3.0 \times 10^{-9} M$
(b) $6.0 \times 10^{-2} M$
(c) $8.0 \times 10^{-12} M$
(d) $5.0 \times 10^{-7} M$

David Collins

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00:23

Problem 28

Which is more acidic, a beer with $\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=3.16 \times 10^{-5} \mathrm{or}$ a wine with $\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]=5.01 \times 10^{-4} ?$

David Collins

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01:19

Problem 29

What is the $\left[\mathrm{OH}^{-}\right]$ and $\mathrm{p} \mathrm{OH}$ of each solution?
(a) $0.10 M \mathrm{KOH}, \mathrm{pH}=13.0$
(b) $0.10 M \mathrm{Na}_{2} \mathrm{CO}_{3}, \mathrm{pH}=11.6$
(c) $0.10 M \mathrm{Na}_{3} \mathrm{PO}_{4}, \mathrm{pH}=12.0$
(d) $0.10 M \mathrm{NaHCO}_{3}, \mathrm{pH}=8.4$

David Collins

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00:12

Problem 30

What is the purpose of an acid-base titration?

David Collins

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00:24

Problem 31

What is the molarity of a solution made by dissolving $12.7 \mathrm{g}$ of $\mathrm{HCl}$ in enough water to make $1.00 \mathrm{L}$ of solution?

David Collins

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00:55

Problem 32

What is the molarity of a solution made by dissolving $3.4 \mathrm{g}$ of $\mathrm{Ba}(\mathrm{OH})_{2}$ in enough water to make $450 \mathrm{mL}$ of solution? Assume that $\mathrm{Ba}(\mathrm{OH})_{2}$ ionizes completely in water to $\mathrm{Ba}^{2+}$ and $\mathrm{OH}^{-}$ ions. What is the pH of the solution?

David Collins

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02:26

Problem 33

Describe how you would prepare each of the following solutions (in each case, assume that the base is a solid).
(a) $400.0 \mathrm{mL}$ of $0.75 \mathrm{M} \mathrm{NaOH}$
(b) $1.0 \mathrm{L}$ of $0.071 \mathrm{MBa}(\mathrm{OH})_{2}$
(c) $500.0 \mathrm{mL}$ of $0.1 \mathrm{M} \mathrm{KOH}$
(d) $2.0 \mathrm{L}$ of $0.3 \mathrm{M}$ sodium acetate

David Collins

Numerade Educator

Problem 34

If $25.0 \mathrm{mL}$ of an aqueous solution of $\mathrm{H}_{2} \mathrm{SO}_{4}$ requires $19.7 \mathrm{mL}$ of $0.72 M \mathrm{NaOH}$ to reach the end point, what is the molarity of the $\mathrm{H}_{2} \mathrm{SO}_{4}$ solution?

David Collins

Numerade Educator

Problem 35

A sample of $27.0 \mathrm{mL}$ of $0.310 \mathrm{M} \mathrm{NaOH}$ is titrated with $0.740 M \mathrm{H}_{2} \mathrm{SO}_{4} .$ How many milliliters of the $\mathrm{H}_{2} \mathrm{SO}_{4}$
solution are required to reach the end point?

David Collins

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00:47

Problem 36

A $0.300 M$ solution of $\mathrm{H}_{2} \mathrm{SO}_{4}$ was used to titrate $10.00 \mathrm{mL}$ of $\mathrm{NaOH} ; 15.00 \mathrm{mL}$ of acid was required to neutralize the basic solution. What was the molarity of the base?

David Collins

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00:20

Problem 37

A solution of NaOH base was titrated with $0.150 M$ $\mathrm{HCl},$ and $22.0 \mathrm{mL}$ of acid was needed to reach the end point of the titration. How many moles of the base were in the solution?

David Collins

Numerade Educator

Problem 38

The usual concentration of $\mathrm{HCO}_{3}^{-}$ ions in blood plasma is approximately 24 millimoles per liter $(\mathrm{mmol} / \mathrm{L}) .$ How would you make up $1.00 \mathrm{L}$ of a solution containing this concentration of $\mathrm{HCO}_{3}^{-}$ ions?

David Collins

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00:24

Problem 39

What is the end point of a titration?

David Collins

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00:29

Problem 40

Why does a titration not tell us the acidity or basicity of a solution?

David Collins

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00:30

Problem 41

Write equations to show what happens when, to a buffer solution containing equimolar amounts of $\mathrm{CH}_{3} \mathrm{COOH}$ and $\mathrm{CH}_{3} \mathrm{COO}^{-},$ we add:
(a) $\mathrm{H}_{3} \mathrm{O}^{+}$
(b) $\mathrm{OH}^{-}$

David Collins

Numerade Educator

00:30

Problem 42

Write equations to show what happens when, to a buffer solution containing equimolar amounts of $\mathrm{HPO}_{4}^{2-}$ and $\mathrm{H}_{2} \mathrm{PO}_{4}^{-},$ we add
(a) $\mathrm{H}_{3} \mathrm{O}^{+}$
(b) $\mathrm{OH}^{-}$

David Collins

Numerade Educator

00:36

Problem 43

We commonly refer to a buffer as consisting of approximately equal molar amounts of a weak acid and its conjugate base - for example, $\mathrm{CH}_{3} \mathrm{COOH}$ and $\mathrm{CH}_{3} \mathrm{COO}^{-} .$ Is it also possible to have a buffer consisting of approximately equal molar amounts of a weak base and its conjugate acid? Explain.

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00:34

Problem 44

What is meant by buffer capacity?

David Collins

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00:35

Problem 45

How can you change the pH of a buffer? How can you change the capacity of a buffer?

David Collins

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01:03

Problem 46

What is the connection between buffer action and Le Chatelier's principle?

David Collins

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00:28

Problem 47

Give two examples of a situation where you would want a buffer to have unequal amounts of the conjugate acid and the conjugate base.

David Collins

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00:16

Problem 48

How is the buffer capacity affected by the ratio of the conjugate base to the conjugate acid?

David Collins

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01:04

Problem 49

Can $100 \mathrm{mL}$ of $0.1 \mathrm{M}$ phosphate buffer at $\mathrm{pH} 7.2$ act as an effective buffer against $20 \mathrm{mL}$ of $1 \mathrm{M} \mathrm{NaOH} ?$

David Collins

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00:31

Problem 50

What is the pH of a buffer solution made by dissolving $0.10 \mathrm{mol}$ of formic acid, $\mathrm{HCOOH}$, and $0.10 \mathrm{mol}$ of sodium formate, HCOONa, in 1 L of water?

David Collins

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01:03

Problem 51

The pH of a solution made by dissolving $1.0 \mathrm{mol}$ of propanoic acid and 1.0 mol of sodium propanoate in $1.0 \mathrm{L}$ of water is 4.85.
(a) What would the pH be if we used $0.10 \mathrm{mol}$ of each (in $1 \mathrm{L}$ of water) instead of $1.0 \mathrm{mol} ?$
(b) With respect to buffer capacity, how would the two solutions differ?

David Collins

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00:26

Problem 52

A $0.15 M \mathrm{HNO}_{2}$ aqueous solution is mixed with a $0.20 M \mathrm{NaNO}_{2}$ aqueous solution, where the $\mathrm{pK}_{\mathrm{a}}$ of $\mathrm{HNO}_{2}$ is equal to 3.37 . What is the $\mathrm{pH}$ of the resulting solution?

David Collins

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00:45

Problem 53

A $0.040 M \mathrm{NaCH}_{3} \mathrm{CO}_{2}$ aqueous solution is mixed with a $0.080 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH}$ aqueous solution, where the $\mathrm{p} K_{\mathrm{a}}$
of $\mathrm{CH}_{3} \mathrm{COOH}$ is equal to $4.75 .$ What is the $\mathrm{pH}$ of the resulting solution?

David Collins

Numerade Educator

00:19

Problem 54

Show that when the concentration of the weak acid, $[\mathrm{HA}],$ in an acid-base buffer equals that of the conjugate base of the weak acid, $\left[\mathrm{A}^{-}\right],$ the $\mathrm{pH}$ of the buffer solution is equal to the $\mathrm{p} K_{\mathrm{a}}$ of the weak acid.

David Collins

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00:28

Problem 55

Show that the pH of a buffer is 1 unit higher than its $\mathrm{p} K_{\mathrm{a}}$ when the ratio of $\mathrm{A}^{-}$ to $\mathrm{HA}$ is 10 to 1.

David Collins

Numerade Educator

01:15

Problem 56

Calculate the pH of an aqueous solution containing the following:
(a) $0.80 M$ lactic acid and $0.40 \mathrm{M}$ lactate ion
(b) $0.30 M \mathrm{NH}_{3}$ and $1.50 \mathrm{M} \mathrm{NH}_{4}^{+}$

David Collins

Numerade Educator

00:35

Problem 57

The pH of $0.10 \mathrm{M} \mathrm{HCl}$ is $1.0 .$ When $0.10 \mathrm{mol}$ of sodium acetate, $\mathrm{CH}_{3} \mathrm{COONa}$, is added to this solution, its pH changes to $2.9 .$ Explain why the pH changes and why it changes to this particular value.

David Collins

Numerade Educator

00:54

Problem 58

If you have $100 \mathrm{mL}$ of a $0.1 \mathrm{M}$ buffer made of $\mathrm{NaH}_{2} \mathrm{PO}_{4}$ and $\mathrm{Na}_{2} \mathrm{HPO}_{4}$ that is at $\mathrm{pH} 6.8$ and you add
10 mL of $1 M$ HCl, will you still have a usable buffer? Why or why not?

David Collins

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00:15

Problem 59

Write an equation showing the reaction of TRIS in the acid form with sodium hydroxide (do not write out the chemical formula for TRIS).

David Collins

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00:21

Problem 60

What is the pH of a solution that is $0.1 \mathrm{M}$ in TRIS in the acid form and $0.05 M$ in TRIS in the basic form?

David Collins

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00:17

Problem 61

Explain why you do not need to know the chemical formula of a buffer compound to use it.

David Collins

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00:16

Problem 62

If you have a HEPES buffer at pH $4.75,$ will it be a usable buffer? Why or why not?

David Collins

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00:26

Problem 63

Which of the compounds listed in Table 8.6 would be the most effective for making a buffer at pH $8.15 ?$ Why?

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00:32

Problem 64

Which of the compounds listed in Table 8.6 would be the most effective for making a buffer at pH $7.0 ?$

David Collins

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00:06

Problem 65

(Chemical Connections $8 \mathrm{A}$ ) Which weak base is used as a flame retardant in plastics?

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00:08

Problem 66

(Chemical Connections $8 \mathrm{B}$ ) Name the most common bases used in over-the-counter antacids.

David Collins

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00:19

Problem 67

(Chemical Connections $8 \mathrm{C}$ ) What causes
(a) respiratory acidosis and
(b) metabolic acidosis?

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00:26

Problem 68

(Chemical Connections 8 D) Explain how the sprinter's trick works. Why would an athlete want to raise the $\mathrm{pH}$ of his or her blood?

David Collins

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00:53

Problem 69

(Chemical Connections 8 D) Another form of the sprinter's trick is to drink a sodium bicarbonate shake before the event. What would be the purpose of doing so? Give the relevant equations.

David Collins

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00:53

Problem 70

4-Methylphenol, $\mathrm{CH}_{3} \mathrm{C}_{6} \mathrm{H}_{4} \mathrm{OH}\left(\mathrm{p} K_{\mathrm{a}}=10.26\right),$ is
only slightly soluble in water, but its sodium salt, $\mathrm{CH}_{3} \mathrm{C}_{6} \mathrm{H}_{4} \mathrm{O}^{-} \mathrm{Na}^{+},$ is quite soluble in water. In which of
the following solutions will 4 -methylphenol dissolve more readily than in pure water?
(a) Aqueous $\mathrm{NaOH}$
(b) Aqueous $\mathrm{NaHCO}_{3}$
(c) Aqueous $\mathrm{NH}_{3}$

David Collins

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00:38

Problem 71

Benzoic acid, $\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\left(\mathrm{p} K_{\mathrm{a}}=4.19\right),$ is only slightly
soluble in water, but its sodium salt, $\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COO}^{-} \mathrm{Na}^{+}$,is quite soluble in water. In which of the following solutions will benzoic acid dissolve more readily than in pure water?
(a) Aqueous NaOH
(b) Aqueous $\mathrm{NaHCO}_{3}$
(c) Aqueous $\mathrm{Na}_{2} \mathrm{CO}_{3}$

David Collins

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00:38

Problem 72

Assume that you have a dilute solution of HCl $(0.10 M)$ and a concentrated solution of acetic acid $(5.0 M) .$ Which solution is more acidic? Explain.

David Collins

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00:56

Problem 73

Which of the two solutions from Problem 8.72 would take a greater amount of $\mathrm{NaOH}$ to hit a phenolphthalein end point assuming you had equal volumes of the two? Explain.

David Collins

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Problem 74

What is the pH of a solution if you mix $300 . \mathrm{mL}$ of $0.30 M$ TRIS in the base form with $250 .$ mL of $0.15 M$ TRIS in the acidic form?

David Collins

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00:25

Problem 75

What is the molarity of a solution made by dissolving $0.583 \mathrm{g}$ of the diprotic acid oxalic acid, $\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4},$ in enough water to make $1.75 \mathrm{L}$ of solution?

David Collins

Numerade Educator

01:16

Problem 76

Following are three organic acids and the $\mathrm{p} K_{\mathrm{a}}$ of each:
butanoic acid, $4.82 ;$ barbituric acid, $5.00 ;$ and lactic acid, 3.85.
(a) What is the $K_{\mathrm{a}}$ of each acid?
(b) Which of the three is the strongest acid, and which is the weakest?
(c) What information do you need to predict which of the three acids would require the most NaOH to reach a phenolphthalein end point?

David Collins

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00:46

Problem 77

The $\mathrm{p} K_{\mathrm{n}}$ value of barbituric acid is $5.0 .$ If the $\mathrm{H}_{3} \mathrm{O}^{+}$ and barbiturate ion concentrations are each $0.0030 M,$ what is the concentration of the undissociated barbituric acid?

David Collins

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00:36

Problem 78

If pure water self-ionizes to give $\mathrm{H}_{3} \mathrm{O}^{+}$ and $\mathrm{OH}^{-}$ ions, why doesn't pure water conduct an electric current?

David Collins

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00:25

Problem 79

Can an aqueous solution have a pH of zero? Explain your answer using aqueous HCl as your example.

David Collins

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00:47

Problem 80

If an acid, HA, dissolves in water such that the $K_{\mathrm{a}}$ is 1000, what is the $\mathrm{p} K_{\mathrm{s}}$ of that acid? Is this scenario possible?

David Collins

Numerade Educator

00:16

Problem 81

A scale of $K_{\mathrm{b}}$ values for bases could be set up in a manner similar to that for the $K_{\mathrm{a}}$ scale for acids. However, this setup is generally considered unnecessary. Explain.

David Collins

Numerade Educator

00:54

Problem 82

Do a $1.0 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH}$ solution and a $1.0 \mathrm{M} \mathrm{HCl}$
solution have the same pH? Explain.

David Collins

Numerade Educator

00:30

Problem 83

Do a $1.0 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH}$ solution and a $1.0 \mathrm{M} \mathrm{HCl}$
solution require the same amount of $1.0 \mathrm{M} \mathrm{NaOH}$ to hit a titration end point? Explain.

David Collins

Numerade Educator

02:05

Problem 84

Suppose you wish to make a buffer whose pH is 8.21. You have available $1 \mathrm{L}$ of $0.100 \mathrm{M} \mathrm{NaH}_{2} \mathrm{PO}_{4}$ and solid $\mathrm{Na}_{2} \mathrm{HPO}_{4} .$ How many grams of the solid $\mathrm{Na}_{2} \mathrm{HPO}_{4}$
must be added to the stock solution to accomplish this task? (Assume that the volume remains 1 L.)

David Collins

Numerade Educator

00:51

Problem 85

In the past, boric acid was used to rinse an inflamed eye. What is the $\mathrm{H}_{3} \mathrm{BO}_{3} / \mathrm{H}_{2} \mathrm{BO}_{3}^{-}$ ratio in a borate buffer solution that has a pH of $8.40 ?$

David Collins

Numerade Educator

00:43

Problem 86

Suppose you want to make a $\mathrm{CH}_{3} \mathrm{COOH} / \mathrm{CH}_{3} \mathrm{COO}^{-}$
buffer solution with a pH of $5.60 .$ The acetic acid concentration is to be $0.10 \mathrm{M}$. What should the acetate ion concentration be?

David Collins

Numerade Educator

00:25

Problem 87

For an acid-base reaction, one way to determine the position of equilibrium is to say that the larger of the equilibrium arrow pair points to the acid with the higher value of $\mathrm{p} K_{\mathrm{a}} .$ For example,
Explain why this rule works.

David Collins

Numerade Educator

01:02

Problem 88

When a solution prepared by dissolving $4.00 \mathrm{g}$ of an unknown monoprotic acid in $1.00 \mathrm{L}$ of water is titrated with $0.600 \mathrm{M} \mathrm{NaOH}, 38.7 \mathrm{mL}$ of the $\mathrm{NaOH}$ solution is needed to neutralize the acid. Determine the molarity of the acid solution. What is the molar mass of the unknown acid?

David Collins

Numerade Educator

00:33

Problem 89

Write equations to show what happens when, to a buffer solution containing equal amounts of HCOOH and $\mathrm{HCOO}^{-},$ we add:
(a) $\mathrm{H}_{3} \mathrm{O}^{+}$
(b) $\mathrm{OH}^{-}$

David Collins

Numerade Educator

Problem 90

If we add $0.10 \mathrm{mol}$ of $\mathrm{NH}_{3}$ to $0.50 \mathrm{mol}$ of $\mathrm{HCl}$ dissolved in enough water to make $1.0 \mathrm{L}$ of solution, what happens to the $\mathrm{NH}_{3}$ ? Will any $\mathrm{NH}_{3}$ remain? Explain.

David Collins

Numerade Educator

01:29

Problem 91

Suppose you have an aqueous solution prepared by dissolving $0.050 \mathrm{mol}$ of $\mathrm{NaH}_{2} \mathrm{PO}_{4}$ in $1 \mathrm{L}$ of water. This solution is not a buffer, but suppose you want to make it into one. How many moles of solid $\mathrm{Na}_{2} \mathrm{HPO}_{4}$ must you add to this aqueous solution to make it into:
(a) A buffer of pH 7.21
(b) A buffer of pH 6.21
(c) A buffer of $\mathrm{pH} 8.21$

David Collins

Numerade Educator

Problem 92

The pH of a $0.10 \mathrm{M}$ solution of acetic acid is 2.93. When 0.10 mol of sodium acetate, $\mathrm{CH}_{3} \mathrm{COONa}$, is added to this solution, its pH changes to 4.74 Explain why the pH changes and why it changes to this particular value.

David Collins

Numerade Educator

00:41

Problem 93

Suppose you have a phosphate buffer $\left(\mathrm{H}_{2} \mathrm{PO}_{4}^{-} / \mathrm{HPO}_{4}^{2-}\right)$
of pH 7.21 . If you add more solid $\mathrm{NaH}_{2} \mathrm{PO}_{4}$ to this buffer, would you expect the pH of the buffer to increase, decrease, or remain unchanged? Explain.

David Collins

Numerade Educator

Problem 94

Suppose you have a bicarbonate buffer containing carbonic acid, $\mathrm{H}_{2} \mathrm{CO}_{3}$, and sodium bicarbonate, $\mathrm{NaHCO}_{3}$, and the pH of the buffer is 6.37 . If you add more solid $\mathrm{NaHCO}_{3}$ to this buffer solution, would you expect its pH to increase, decrease, or remain unchanged? Explain.

David Collins

Numerade Educator