16.8
Indicators
Certain organic substances change color in dilute solution when the hydronium ion concentration reaches a particular value. For example, phenolphthale…
Indicators change color depending on pH and can be used to monitor the pH of a solution.
A pH indicator contains a weak acid, represented by HIn, which changes color when converted to its conjugate base, the In ion. The color changes over the indicator’s pH interval, which depends on its pKa.
If a solution has a pH below the pKa of the indicator, the hydronium concentration is high. The equilibrium of the reaction shifts towards the left, increasing the concentration of HIn.
When a base is added, the concentration of hydronium decreases. This pH change converts HIn into the In ion, and the solution’s color shifts.
The HIn color dominates when the pH is at least one lower than the pKa; the In ion color dominates when the pH is at least one higher, and the color is a mix at pHs in between.
There are a variety of pH indicators which can be used to monitor the pH that change color over distinct pH ranges.
The point at which an indicator changes color is called the endpoint. The ideal indicator for a titration has an endpoint near the equivalence point.
The titration curve of a strong acid and base is steep due to the rapid pH change near the equivalence point. These titrations can use a relatively broad range of indicators.
During the titration of 0.1 molar hydrochloric acid with 0.1 molar sodium hydroxide, the pH quickly increases from 3 to 11. Either phenolphthalein, an indicator with the pH interval of 8.3 to 10, or methyl red, an indicator with an interval of 4.2 to 6, could be used here as their endpoints overlap with the steep portion of the curve.
In contrast, a titration involving a weak acid or base has a less steep curve than one with only a strong acid and a strong base; therefore, a relatively narrow range of indicators can be used.
During the titration of 0.1 molar acetic acid with 0.1 molar sodium hydroxide, the pH increases quickly from 7 to 11. Phenolphthalein can be used here as its endpoint overlaps with the equivalence point, but methyl red would be an ineffective indicator.
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Q1: How does an acid-base indicator change color in solution?
An acid-base indicator contains a weak acid (HIn) that converts to its conjugate base (In−) as pH changes. Each form has a distinct color. When hydronium concentration is high, the equilibrium shifts left, favoring HIn and its color. When base is added, hydronium decreases, shifting equilibrium right toward In− and its color. The perceived color depends on the ratio of these two species.
Q2: What is the pH interval of an indicator and how does it relate to pKa?
The pH interval is the range of pH values over which an indicator visibly changes color, typically spanning approximately pKa ± 1. When pH is one unit below pKa, the HIn color dominates. When pH is one unit above pKa, the In− color dominates. Between these points, the solution displays a mixed color. This relationship is described by the henderson hasselbalch equation calculating buffers.
Q3: Why is indicator selection important for strong acid-strong base titrations?
Strong acid-strong base titrations produce steep titration curves with rapid pH changes near the equivalence point. This steep region allows a broad range of indicators to work effectively. For example, during titration of 0.1 M hydrochloric acid with 0.1 M sodium hydroxide, pH jumps from 3 to 11, so both phenolphthalein (pH 8.3–10) and methyl red (pH 4.2–6) are suitable because their endpoints overlap the steep portion.
Q4: Why must weak acid-base titrations use more selective indicators?
Weak acid-base titrations produce less steep titration curves than strong acid-strong base reactions, creating a narrower region of rapid pH change. This limits which indicators can be used effectively. During titration of 0.1 M acetic acid with 0.1 M sodium hydroxide, pH rises from 7 to 11. Phenolphthalein works because its endpoint overlaps the equivalence point, but methyl red fails because its endpoint misses this critical region.
Q5: What is the endpoint of a titration and how does it relate to indicator choice?
The endpoint is the point at which an indicator visibly changes color during titration. The ideal indicator has an endpoint positioned near the equivalence point, where the analyte is completely neutralized. For accurate titration results, the indicator's pH interval must overlap the steep portion of the titration curve. Misalignment between endpoint and equivalence point leads to systematic errors in volume measurements.
Q6: How do universal indicators differ from single-color indicators?
Universal indicators and pH paper contain a mixture of multiple indicators, each with different pH intervals and colors. This combination allows them to display distinct colors across a wide range of pH values, enabling approximate pH determination of unknown solutions through visual comparison. Single-color indicators like phenolphthalein or methyl red change color over a narrow, specific pH range and are better suited for precise titration endpoints.
Q7: How does Le Châtelier's principle explain indicator behavior when acid or base is added?
An indicator's equilibrium between HIn and In− responds to changes in hydronium concentration according to Le Châtelier's principle. Adding acid increases hydronium, shifting equilibrium left toward HIn and its associated color. Adding base decreases hydronium, shifting equilibrium right toward In− and its color. This reversible color change makes indicators useful for monitoring pH changes in titrations and other acid-base reactions.