Experiment_9_-_pH_indicators

Experiment_9_-_pH_indicators - 9 Acid-Base Indicators and...

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9 Acid-Base Indicators and pH Introduction Acid-base indicators can be used to determine the pH of a solution. In this experiment, you will prepare your own acid-base indicator by extracting a colored plant pigment. Then you will use your indicator and others available in the lab to estimate the pH of several other solutions. Acid-base indicators are substances that exist in different colored forms depending on pH. This behavior is possible because each acid-base indicator exists in conjugate acid and base forms. The following equilibrium between the acid form (“HIn”) and the base form (“In ”) of a hypothetical indicator exists in aqueous solutions: HIn + H 2 O H 3 O + + In (blue) (red) Suppose that HIn is blue and In is red. We can use Le Châtelier’s principle to explain the color of this indicator in strongly acidic or basic solutions. If we add acid to the indicator solution, we will be adding excess H 3 O + . The reaction will proceed to the left to restore equilibrium, and red In will be converted to blue HIn. Therefore, this indicator will be blue in acidic solutions (with low pH). If we add base to the indicator solution, we will be removing H 3 O + . The reaction will proceed to the right to restore equilibrium, and blue HIn will be converted to red In . Therefore, this indicator will be red in basic solutions (with high pH). The exact pH range in which the indicator changes color depends on the equilibrium constant for the equilibrium between the acid and base forms of the indicator. We can make our discussion more sophisticated by recognizing that indicators are weak acids and bases. Suppose that our hypothetical indicator has K in = 3 × 10 –5 . Then, the following reactant-favored equilibrium between blue HIn and red In exists in acidic aqueous solutions: HIn + H 2 O H 3 O + + In K in = 3 × 10 –5 (blue) (red) Chemistry 1B Experiment 9 33
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Since the dissociation of HIn in water is reactant-favored, the indicator will be blue in acidic solutions. In basic aqueous solutions, the following equilibrium would exist: Chemistry 1B Experiment 9 34
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OH + HIn In + H 2 O K = Chemistry 1B Experiment 9 35
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Chemistry 1B Experiment 9 36
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= 3 × 10 +9 Since the reaction of HIn with base is product-favored, the indicator will be red in basic solutions. Suppose you put a tiny amount of this indicator into a solution that already has a certain pH. If the concentration of indicator is very small, the indicator itself will not noticeably change the pH of the solution. Rather, the relative amounts of the acid and base forms of the indicator will be determined by the pH and the value of K in . We can see this by rearranging the equilibrium constant expression for the indicator: K in = so = You can derive the same formula for the relative amounts of the acid and base forms in basic solutions, using the chemical equation for the equilibrium with OH , and remembering that [H 3 O + ] is equal to K w divided by [OH
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Experiment_9_-_pH_indicators - 9 Acid-Base Indicators and...

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