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Chapter 16: Problem 41

Explain how an acid-base indicator works in a titration. What are the criteria for choosing an indicator for a particular acid-base titration?

Short Answer

Expert verified
An acid-base indicator is a substance that changes color based on the pH of the solution it's in. In a titration, it's used to indicate when the reaction has reached its equivalence point by changing color. The choice of indicator depends on the pH at the equivalence point of the titration, as different indicators change color at different pH values.

Step by step solution

01

Understanding Acid-Base Indicators

An acid-base indicator is a substance that changes its color in response to a change in pH. This means it can signal, or 'indicate,' when the solution it is in has a certain pH.
02

Role of Indicators in Titrations

In a titration experiment, an indicator is used to signal when the reaction has reached its equivalence point, or when the solution is neutral. The indicator does this by changing color.
03

Criteria for Choosing an Indicator

The choice of indicator depends on the pH at the equivalence point of the titration. Different indicators change color at different pH values, so the chosen indicator should have a pKa value close to the pH of the titration's equivalence point for an accurate end point determination.

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Most popular questions from this chapter

A 0.2688 -g sample of a monoprotic acid neutralizes \(16.4 \mathrm{~mL}\) of \(0.08133 \mathrm{M}\) KOH solution. Calculate the molar mass of the acid.

The molar mass of a certain metal carbonate, \(\mathrm{MCO}_{3}\), can be determined by adding an excess of \(\mathrm{HCl}\) acid to react with all the carbonate and then "back titrating" the remaining acid with a \(\mathrm{NaOH}\) solution. (a) Write equations for these reactions. (b) In a certain experiment, \(18.68 \mathrm{~mL}\) of \(5.653 \mathrm{M} \mathrm{HCl}\) were added to a \(3.542-\mathrm{g}\) sample of \(\mathrm{MCO}_{3}\). The excess HCl required \(12.06 \mathrm{~mL}\) of \(1.789 \mathrm{M} \mathrm{NaOH}\) for neutralization. Calculate the molar mass of the carbonate and identify \(\mathrm{M}\).

Calculate the molar solubility of \(\mathrm{AgCl}\) in a \(1.00-\mathrm{L}\) solution containing \(10.0 \mathrm{~g}\) of dissolved \(\mathrm{CaCl}_{2}\).

Sketch titration curves for the following acid-base titrations: (a) \(\mathrm{HCl}\) versus \(\mathrm{NaOH},\) (b) \(\mathrm{HCl}\) versus \(\mathrm{CH}_{3} \mathrm{NH}_{2},\) (c) \(\mathrm{CH}_{3} \mathrm{COOH}\) versus \(\mathrm{NaOH}\). In each case, the base is added to the acid in an Erlenmeyer flask. Your graphs should show \(\mathrm{pH}\) on the \(y\) axis and volume of base added on the \(x\) axis.

From Table 16.2 we see that silver bromide (AgBr) has a larger solubility product than iron(II) hydroxide \(\left[\mathrm{Fe}(\mathrm{OH})_{2}\right] .\) Does this mean that \(\mathrm{AgBr}\) is more soluble than \(\mathrm{Fe}(\mathrm{OH})_{2} ?\)
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