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Chapter 17: Problem 29

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?

Short Answer

Expert verified
Part (a): A single indicator is enough in an acid-base titration as its color change occurs at the equivalence point, marking the end of the titration. However, to approximate the pH of a solution, we often need multiple indicators because each has a distinct color change at a specific pH range. Part (b): The quantity of indicator used in a titration should be minimized to avoid it interfering with the test results, as it can potentially shift the pH and the equivalence point of the reaction, influencing its accuracy.

Step by step solution

01

Part (a) Step 1: Understanding Single Indicator Usage in Acid-Base Titrations

In an acid-base titration, the purpose is to determine the unknown concentration of an acidic or basic solution. A single indicator is typically sufficient because the color change of an indicator happens very quickly around its transition pH, which corresponds to the equivalence point in the titration. This color change is distinct enough to help determine the end point of the titration (when equal amounts of acid and base have reacted).
02

Part (a) Step 2: Understanding Need for Multiple Indicators to Determine Approximate pH of a Solution

When determining the approximate pH of a solution, however, it's often necessary to use several indicators. This is because each indicator changes its color at a particular pH range, and this range varies across different indicators. Hence, using multiple indicators provides a wider range of pH values, and thus a better approximation of the pH of a solution.
03

Part (b) Step 1: Understanding the Reason for Minimal Usage of Indicators in Titrations

The quantity of indicator used in a titration must be kept at a minimum level for multiple reasons. Primarily, the indicator itself might undergo a chemical reaction and interfere with the results. More specifically, added indicator can potentially shift the pH of the solution. It's vital to remember that indicators are also weak acids or weak bases, and adding them in large amounts could influence the pH of the solution and the equivalence point, leading to less accurate results.

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Most popular questions from this chapter

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 \(5.00 \mathrm{g} \mathrm{Ba}(\mathrm{OH})_{2}\)

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.

A 25.00 -mL sample of \(0.0100 \mathrm{M} \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\left(\mathrm{K}_{\mathrm{a}}=\right.\) \(\left.6.3 \times 10^{-5}\right)\) is titrated with \(0.0100 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\) Calculate the \(\mathrm{pH}\) (a) of the initial acid solution; (b) after the addition of 6.25 mL of \(0.0100 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\) (c) at the equivalence point; (d) after the addition of a total of \(15.00 \mathrm{mL}\) of \(0.0100 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\)

Indicate whether you would expect the equivalence point of each of the following titrations to be below, above, or at \(\mathrm{pH}\) 7. Explain your reasoning. (a) \(\mathrm{NaHCO}_{3}(\mathrm{aq})\) is titrated with \(\mathrm{NaOH}(\mathrm{aq})\) (a) (b) \(\mathrm{HCl}(\mathrm{aq})\) is titrated with \(\mathrm{NH}_{3}(\mathrm{aq}) ;\) (c) \(\mathrm{KOH}(\mathrm{aq})\) is titrated with HI(aq).

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 \(5.50 \mathrm{g} \mathrm{Ba}(\mathrm{OH})_{2} ?\)
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