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Chapter 14: Problem 37

Given three acid-base indicators-methyl orange (end point at \(\mathrm{pH}\) 4), bromthymol blue (end point at \(\mathrm{pH} 7\) ), and phenolphthalein (end point at \(\mathrm{pH}\) 9) - which would you select for the following acid-base titrations? (a) perchloric acid with an aqueous solution of ammonia (b) nitrous acid with lithium hydroxide (c) hydrobromic acid with strontium hydroxide (d) sodium fluoride with nitric acid

Short Answer

Expert verified
Answer: (a) Methyl orange, (b) Phenolphthalein, (c) Bromthymol blue, (d) Methyl orange.

Step by step solution

01

(a) Perchloric acid with an aqueous solution of ammonia

Perchloric acid (HClO4) is a strong acid, and ammonia (NH3) is a weak base. For the titration of a strong acid with a weak base, the pH at the equivalence point is expected to be less than 7. So, we need to choose an indicator with end point pH less than 7. Among the given indicators, methyl orange has an end point at pH 4, which is suitable for this titration.
02

(b) Nitrous acid with lithium hydroxide

Nitrous acid (HNO2) is a weak acid, and lithium hydroxide (LiOH) is a strong base. For the titration of a weak acid with a strong base, the pH at the equivalence point is expected to be greater than 7. Thus, we need to choose an indicator with end point pH greater than 7. Among the given indicators, phenolphthalein has an end point at pH 9, which is suitable for this titration.
03

(c) Hydrobromic acid with strontium hydroxide

Hydrobromic acid (HBr) is a strong acid, and strontium hydroxide (Sr(OH)2) is a strong base. For the titration of a strong acid with a strong base, the pH at the equivalence point is expected to be around 7. Therefore, we should choose an indicator with end point pH approximately 7. Among the given indicators, bromthymol blue has an end point at pH 7, which is suitable for this titration.
04

(d) Sodium fluoride with nitric acid

Sodium fluoride (NaF) is a salt of a strong base (NaOH) and a weak acid (HF). It will react with the strong acid, nitric acid (HNO3), in a titration that is essentially between the weak base (F-) and the strong acid (HNO3). For the titration of a weak base with a strong acid, the pH at the equivalence point is expected to be less than 7. Thus, we need to choose an indicator with end point pH less than 7. Among the given indicators, methyl orange has an end point at pH 4, which is suitable for this titration.

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Most popular questions from this chapter

What is the \(\mathrm{pH}\) of a \(0.1500 \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4}\) solution if (a) the ionization of \(\mathrm{HSO}_{4}^{-}\) is ignored? (b) the ionization of \(\mathrm{HSO}_{4}^{-}\) is taken into account? \(\left(K_{\mathrm{a}}\right.\) for \(\mathrm{HSO}_{4}^{-}\) is \(\left.1.1 \times 10^{-2} .\right)\)

A \(0.1375 \mathrm{M}\) solution of potassium hydroxide is used to titrate \(35.00 \mathrm{~mL}\) of \(0.257 M\) hydrobromic acid. (Assume that volumes are additive.) (a) Write a balanced net ionic equation for the reaction that takes place during titration. (b) What are the species present at the equivalence point? (c) What volume of potassium hydroxide is required to reach the equivalence point? (d) What is the \(\mathrm{pH}\) of the solution before any \(\mathrm{KOH}\) is added? (e) What is the \(\mathrm{pH}\) of the solution halfway to the equivalence point? (f) What is the \(\mathrm{pH}\) of the solution at the equivalence point?

A buffer is prepared using the propionic acid/propionate \(\left(\mathrm{HC}_{3} \mathrm{H}_{5} \mathrm{O}_{2} /\right.\) \(\left.\mathrm{C}_{3} \mathrm{H}_{5} \mathrm{O}_{2}^{-}\right)\) acid-base pair for which the ratio \(\left[\mathrm{HC}_{3} \mathrm{H}_{5} \mathrm{O}_{2}\right] /\left[\mathrm{C}_{3} \mathrm{H}_{5} \mathrm{O}_{2}^{-}\right]\) is \(4.50 .\) \(K_{\mathrm{a}}\) for propionic acid is \(1.4 \times 10^{-5}\). (a) What is the \(\mathrm{pH}\) of this buffer? (b) Enough strong base is added to convert \(27 \%\) of \(\mathrm{HC}_{3} \mathrm{H}_{5} \mathrm{O}_{2}\) to \(\mathrm{C}_{3} \mathrm{H}_{5} \mathrm{O}_{2}^{-} .\) What is the \(\mathrm{pH}\) of the resulting solution? (c) Strong base is added to increase the \(\mathrm{pH}\). What must the acid/base ratio be so that the \(\mathrm{pH}\) increases by exactly one unit (e.g., from 2 to 3 ) from the answer in (a)?

Which of the following would form a buffer if added to \(650.0 \mathrm{~mL}\) of \(0.40 M \mathrm{Sr}(\mathrm{OH})_{2} ?\) (a) \(1.00 \mathrm{~mol}\) of \(\mathrm{HF}\) (b) \(0.75 \mathrm{~mol}\) of \(\mathrm{HF}\) (c) \(0.30 \mathrm{~mol}\) of \(\mathrm{HF}\) (d) \(0.30 \mathrm{~mol}\) of \(\mathrm{NaP}\) (e) \(0.30 \mathrm{~mol}\) of \(\mathrm{HCl}\) Explain your reasoning in each case.

There is a buffer system \(\left(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}-\mathrm{HPO}_{4}{ }^{2-}\right)\) in blood that helps keep the blood \(\mathrm{pH}\) at about \(7.40 .\left(\mathrm{K}_{\mathrm{a}} \mathrm{H}_{2} \mathrm{PO}_{4}^{-}=6.2 \times 10^{-8}\right)\). (a) Calculate the \(\left[\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\right] /\left[\mathrm{HPO}_{4}^{2-}\right]\) ratio at the normal \(\mathrm{pH}\) of blood. (b) What percentage of the \(\mathrm{HPO}_{4}{ }^{2-}\) ions are converted to \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) when the \(\mathrm{pH}\) goes down to \(6.80\) ? (c) What percentage of \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) ions are converted to \(\mathrm{HPO}_{4}{ }^{2-}\) when the \(\mathrm{pH}\) goes up to \(7.80\) ?
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