To determine the indicator range of some acid-base indicators

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  1. fill burette with deionized water, NaOH, HCl + white tile
  2. 1st beaker: 25 HCl + 10 water into beaker
  3. 2nd beaker: 25 NaOH +10 water
  4. 3rd beaker: 25 buffer + 2 drop of indicator

Add more indicator if too pale

  1. Add same no of drop of indicator to 1st and 2nd 
  2. Dilute HCl with water, vol same as 3rd 

Dilute NaOH with water, vol same as 3rd

  1. Add NaOH 1cm3 at a time to 3rd, mix
  2. Measure pH with pH meter just when color change (compare to 1st)
  3. Add NaOH 1cm3 at a time to 3rd, mix
  4. Measure pH until color change is complete (compare with 2nd )

11. Repeat with other indicator

S.K.H. Lam Woo Memorial Secondary School

F.7 Chemistry Laboratory Report

Name: Chan Ching Wai     Class: F.7H (2)

Experiment 4: Indicator

Date of Experiment: 16-11-2010

Objective

To determine the indicator range of some acid-base indicators

Introduction

In this experiment, the indicator ranges of some acid-base indicators were determined. Indicators are chemicals that would change color as the pH of the solution in which they are dissolved changes within the indicator range.

Indicators are commonly used in acid-base equilibrium in order to determine the concentration of a solution. In the titration, the equivalence point, which is the point at which equal quantities of acid and base have reacted, needed to be determined. As the suitable indicators change color obviously or reach end point, near the equivalence point, the end point could signify the equivalence point. However, indicators would only be applicable within the indicator range. If the pH of the solution at equivalence point was out of this range, the indicator shows no change in color and hence failed to serve its function.

 

The principle behind the color change of indicator mainly lies on the ratio between the indicator and its conjugate counterparts which had different color. The indicator was a weak acid or a weak base itself which dissociated in water to set up an ionic equilibrium.

HIn (aq)+ H2O (l)        H3O+ (aq) + In- (aq)     for weak acid

HIn+ (aq) + H2O     H3O+ (aq) + In (aq)       for weak base

Both the indicator and its conjugated counterparts are colored. With the change in equilibrium position, the ratio of the indicator and conjugated counterparts changes so as the ratio of different color. For example, in acidic medium, the equilibrium would shift to the left and the solution. The concentration of conjugated counterpart is higher than the indicator and hence its color masked the original color. The solution would show the color of the dominant species which is HIn for weak acid and HIn+ for weak base. On the other hand, in alkaline medium, the equilibrium would shift to right and the solution would show the color of the dominant species which is In- for weak acid and In for weak base.

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To be specific, a particular ratio of indicator to conjugated counterparts would give a distinct color.

∵ KC = [H3O+ (aq)][ In- (aq)] / [HIn (aq)][ H2O (l)]

   KIn = [H3O+ (aq)][ In- (aq)] / [HIn (aq)]

   [H3O+ (aq)] = KIn X [HIn (aq)] / [ In- (aq)]

   -log [H3O+ (aq)] = -log KIn – log [HIn (aq)] / [ In- (aq)]

   pH = pK In + log [ In- (aq)]/ [HIn (aq)]

The ratio of [ In- (aq)]/ [HIn (aq)] determined the color of the solution.

When [ In- (aq)]/ [HIn (aq)] < 1/10, color of [HIn (aq)] is observed and there ...

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