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Chapter 15: Problem 18

Acid-base indicators mark the end point of titrations by "magically" turning a different color. Explain the "magic" behind acid-base indicators.

Short Answer

Expert verified
The "magic" behind acid-base indicators is their ability to change color in response to pH changes in a solution, due to the shift in equilibrium between their acidic (HIn) and basic (In-) forms. This color change allows one to visually identify the end point in a titration, signaling that the reaction has reached neutralization. Each indicator has a specific pH range in which it works effectively, and it is crucial to choose an indicator with an end point that falls within the expected equivalence point of the titration process.

Step by step solution

01

Understanding acid-base indicators

An acid-base indicator is a substance that changes its color in response to changes in pH (the level of acidity or alkalinity) within a solution. These substances can be either weak acids or weak bases that can react with the acids or bases present in solutions during a titration process.
02

Function of acid-base indicators

The ability of these indicators to change color relies on the equilibrium between two different forms: the acidic form (often denoted as HIn) and the basic form (In-). When the pH of the solution changes, the equilibrium between these two forms shifts, causing the dominant form to change, which is marked by a distinct color change. This is due to the different colors exhibited by different forms 鈥� acidic and basic 鈥� when they absorb light.
03

Types of acid-base indicators

There are several different types of acid-base indicators, each with specific pH ranges in which they work effectively. Common acid-base indicators include phenolphthalein, litmus paper (blue and red), bromothymol blue, and universal indicator. Understanding the working pH range of each indicator is essential for its accurate use.
04

Indicators in titrations

In a titration, the end point is marked by the color change of an acid-base indicator, which signals that the reaction has reached neutralization. In an acid-base titration, this happens when the same number of moles of the acid has reacted with the base, or vice versa. The chosen indicator should have an end point that falls within the pH range of the expected equivalence point.
05

Summing up the "magic"

The "magic" behind acid-base indicators lies in their ability to change color in response to pH changes in a solution. This shift in equilibrium between acidic and basic forms allows scientists and researchers to visually identify when the titration process reaches its end point, signaling that the reaction has reached neutralization. This valuable property makes acid-base indicators an indispensable tool in various fields, such as chemistry, biology, and environmental science.

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Most popular questions from this chapter

Calculate the pH of a solution formed by mixing 100.0 \(\mathrm{mL}\) of 0.100 \(\mathrm{M}\) NaF and 100.0 \(\mathrm{mL}\) of 0.025 \(\mathrm{M} \mathrm{HCl} .\)

Derive an equation analogous to the Henderson-Hasselbalch equation but relating pOH and \(\mathrm{p} K_{\mathrm{b}}\) of a buffered solution composed of a weak base and its conjugate acid, such as \(\mathrm{NH}_{3}\) and \(\mathrm{NH}_{4}^{+}\)

One method for determining the purity of aspirin \(\left(\mathrm{C}_{9} \mathrm{H}_{8} \mathrm{O}_{4}\right)\) is to hydrolyze it with NaOH solution and then to titrate the remaining NaOH. The reaction of aspirin with NaOH is as follows: \(\mathrm{C}_{9} \mathrm{H}_{8} \mathrm{O}_{4}(s)+2 \mathrm{OH}^{-}(a q)\) $$ \mathrm{C}_{7} \mathrm{H}_{3} \mathrm{O}_{3}^{-}(a q)+\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{O}_{2}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(l) $$ A sample of aspirin with a mass of 1.427 g was boiled in 50.00 \(\mathrm{mL}\) of 0.500\(M \mathrm{NaOH}\) . After the solution was cooled, it took 31.92 \(\mathrm{mL}\) of 0.289 \(\mathrm{M} \mathrm{HCl}\) to titrate the excess NaOH. Calculate the purity of the aspirin. What indicator should be used for this titration? Why?

Calculate the number of moles of \(\mathrm{HCl}(g)\) that must be added to 1.0 \(\mathrm{L}\) of 1.0 \(\mathrm{M} \mathrm{NaC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) to produce a solution buffered at each pH. $$ \text{(a)}\mathrm{pH}=\mathrm{p} K_{\mathrm{a}} \quad \text { b. } \mathrm{pH}=4.20 \quad \text { c. } \mathrm{pH}=5.00 $$

Calculate the mass of sodium acetate that must be added to 500.0 \(\mathrm{mL}\) of 0.200\(M\) acetic acid to form a \(\mathrm{pH}=5.00\) buffer solution.
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