1. Indicators

2. Indicators – uses
Used for acid-base titrations to determine when neutralisation has happened
All indicators change colour over a certain pH range

3. Examples of Indicators

4. Definitions you need to know
Equivalence Point:
The point in a neutralisatoin reaction when the acid and base have been added in exactly the right proportions for the chemicals to be neutralised
End Point:
The stage in a titration when the indicator has changed colour
The end point of an indicator should be as close to
the equivalence point as possible

5. Half-equivalence point:
Anatomy of an acid-base titration curve: In this case weak-acid/strong-base
Equivalence point:
i.e. The point of inflection (where the gradient starts to drop again). Represents the point at which the acid has just been neutralised.
Half-equivalence point:
i.e. Half the volume added at the equivalence point
The pH of this point is equal to the pKa of the weak acid.
Buffer Region:
A lot of base has to be added to result in only a small change of pH
pH Intercept:
Higher pH if a weaker acid used

6. Choosing an Indicator
Indicator
Range of
colour change
Bromothymol blue
Methyl orange
Phenolphthalein
Which indicator would be most suitable
for this neutralisation reaction?
The pH range when an indicator will change colour
should lie within the vertical part of the graph.

7. Strong Acid/Strong Base
The pH range when an indicator will change colour
should lie within the vertical part of the graph.

8. Strong Acid/Weak Base
The pH range when an indicator will change colour
should lie within the vertical part of the graph.

9. Weak Acid/Strong Base
The pH range when an indicator will change colour
should lie within the vertical part of the graph.

10. Weak acid + weak base
The pH range when an indicator will change colour
should lie within the vertical part of the graph.

14. Indicators (Hln) Hln (aq)  H+ (aq) + ln- (aq)
Generally weak acids that dissociate:
(acid has a different colour to the conjugate base)
Hln (aq)  H+ (aq) + ln- (aq)
Acid Conjugate base
Good indicators have very different colours of Hln and ln-

15. Colour of the Indicator
Hln (aq)  H+ (aq) + ln- (aq)
Acid Conjugate base
The colour of the indicator is determined by the relative concentrations of HIn and In-
What happens when:
An acid is added?
An alkali is added?

16. Addition of acid Hln (aq)  H+ (aq) + ln- (aq) Increase in [H+]
System tries to decrease [H+]
Reacts H+ and In-
More HIn forms
Solution turns red

17. Addition of an alkali Hln (aq)  H+ (aq) + ln- (aq) OH- reacts with H+
[H+] decreases
System tries to increase [H+]
HIn dissociates into H+ and In-
More In- produced
Solution turns blue

18. Hln (aq)  H+ (aq) + ln- (aq)
In terms of the relative concentrations of Hln and In-,
explain why the indicator reaches an intermediary colour
A colour change begins when [Hln] = [ln-]

19. Equilibrium Constant Hln (aq)  H+ (aq) + ln- (aq)
Kin = [H+] [In-]
[HIn]
A colour change begins when [Hln] = [ln-]
At the colour change what will the expression be?

20. Using Indicators Indicators are NOT USED to measure pH
Colour change is not an accurate measurement
Range of colour change limits the measurements that could be taken
pH probes are very accurate
Indicators ARE USED:
To determine the end-point of reactions in titrations
To give a ‘rough and ready’ idea of pH

21. Making Indicators
In this experiment you will make an indicator from a natural product and design an experiment to determine the pH range over which it changes colour.
Follow the instructions on the sheet
BONUS EXPERIMENT (IF YOU FINISH EARLY):
Chop up some onion as finely as possible
Place it in a small beaker with ~25 cm3 sodium hydroxide
Add 50 cm3 hydrochloric acid in 10 cm3 portions, smelling before and after each one.

22. Key Points Indicators change colour over a narrow pH range
The colour change generally occurs at a pH within ± 1.0 of the indicator’s pKa
Indicators should be chosen with a pH range that matches the expected equivalence point as closely as possible