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Chapter 17: Problem 31

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) \(0.10 \mathrm{M} \mathrm{CH}_{3} \mathrm{COOH}-0.10 \mathrm{M}\) \(\mathrm{NaCH}_{3} \mathrm{COO}\).

Short Answer

Expert verified
The colors phenol red would change to in each solution are: (a) red, (b) yellow, (c) orange, (d) yellow, (e) red, and (f) yellow.

Step by step solution

01

Significance of KOH

KOH is a strong base. Therefore, its solution will have a pH greater than 7, and specifically for this exercise, it will bring the pH over 8.0. Therefore, in a \(0.10 M KOH\) solution, the phenol red indicator will turn red.
02

Significance of CH3COOH

CH3COOH is a weak acid, which means that it will not fully ionize in solution, thereby resulting in a solution with a pH of less than 7, and specifically for this exercise, a pH below 6.6. Therefore, in a \(0.10 M CH3COOH\) solution, the phenol red indicator will turn yellow.
03

Significance of NH4NO3

NH4NO3, a salt of a weak base (NH4) and a strong acid (NO3), results in a neutral solution. The phenol red indicator will neither be yellow nor red but will be orange, which is the indicator color at neutral pH.
04

Significance of HBr

HBr is a strong acid. When in solution, it will cause the pH to fall below 6.6. Thus, in a \(0.10 M HBr\) solution, the phenol red indicator will turn yellow.
05

Significance of NaCN

NaCN, being a salt of a weak acid (HCN) and a strong base (NaOH), will create a weakly basic solution. This means the pH will be over 8.0. Therefore, adding the phenol red indicator to a \(0.10 M NaCN\) solution will result in a red color.
06

Significance of CH3COOH-NaCH3COO

The solution of \(0.10 M CH3COOH-0.10 M NaCH3COO\) is a buffer solution with a pH less than 7 and, for this exercise, a pH below 6.6. This will result in the phenol red indicator turning yellow.

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Most popular questions from this chapter

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).

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 \(\mathrm{H}_{3} \mathrm{PO}_{4}(\mathrm{aq}) ?\)

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.

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 \(0.0100 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\)

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?
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