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Chapter 17: Problem 96

Briefly describe each of the following ideas, phenomena, or methods: (a) the common-ion effect; (b) the use of a buffer solution to maintain a constant \(\mathrm{pH}\) (c) the determination of \(\mathrm{p} K_{\mathrm{a}}\) of a weak acid from a titration curve; (d) the measurement of \(\mathrm{pH}\) with an acid-base indicator.

Short Answer

Expert verified
The common-ion effect is a shift in equilibrium upon addition of an ion already present in the solution. Buffer solutions help maintain a constant pH by balancing concentrations of a weak acid and its conjugate base. The pKa of a weak acid is determined from the pH at the midpoint of its titration curve. Acid-base indicators measure pH by changing color due to structural changes upon donating or accepting protons.

Step by step solution

01

Explain the Common-ion effect

In chemistry, the common-ion effect refers to an effect where the introduction of an ion into a solution that already contains that same ion leads to a suppression of the ionization of a weak electrolyte. This phenomenon is a consequence of Le Chatelier's principle, which states that if a system in equilibrium is disturbed, the equilibrium will shift to counteract the disturbance.
02

Understand the use of buffer solutions

Buffer solutions are solutions that resist changes in their pH when either acid or base is added to them. They find widespread use in labs and industrial processes where a constant pH is necessary. This is achieved by maintaining a balance between a weak acid and its conjugate base (or a weak base and its conjugate acid) in the solution.
03

Learn about the determination of pKa

The pKa of a weak acid is determined from a titration curve, which is a plot of the pH of the solution as a function of the amount of base added. The equivalence point is the point where the amount of base added completely reacts with the weak acid present. Halfway to the equivalence point, the concentrations of the weak acid and conjugate base are equal, and the pH equals the pKa. We can locate this point on the titration curve and read off the pKa value.
04

Discuss the measurement of pH with an acid-base indicator

Acid-base indicators are used to measure the pH of a solution. The indicator, a weak organic acid or base, changes color over a specific pH range. The color change happens because the structure of the indicator molecule changes as it donates or accepts a proton (H+), and different structures have different colors. By comparison with a color chart, one can estimate the pH of the solution.

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Most popular questions from this chapter

If an indicator is to be used in an acid-base titration having an equivalence point in the pH range 8 to 10 , the indicator must (a) be a weak base; (b) have \(K_{\mathrm{a}}=1 \times 10^{-9} ;(\mathrm{c})\) ionize in two steps; (d) be added to the solution only after the solution has become alkaline.

The Henderson-Hasselbalch equation can be written as \(\mathrm{pH}=\mathrm{p} K_{\mathrm{a}}-\log \left(\frac{1}{\alpha}-1\right)\) where \(\alpha=\frac{\left[\mathrm{A}^{-}\right]}{\left[\mathrm{A}^{-}\right]+[\mathrm{HA}]}\) Thus, the degree of ionization \((\alpha)\) of an acid can be determined if both the \(\mathrm{pH}\) of the solution and the \(\mathrm{p} K_{\mathrm{a}}\) of the acid are known. (a) Use this equation to plot the pH versus the degree of ionization for the second ionization constant of phosphoric acid \(\left(K_{\mathrm{a}}=6.3 \times 10^{-8}\right)\) (b) If \(\mathrm{pH}=\mathrm{p} K_{\mathrm{a}}\) what is the degree of ionization? (c) If the solution had a pH of 6.0 what would the value of \(\alpha\) be?

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.]

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 \(\mathrm{pH}=4.52 ?\)

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 \(8.90 ?\)
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