Indicators

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

Examples of Indicators

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

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

Choosing an Indicator Indicator Range of colour change Bromothymol blue 6.0 - 7.6 Methyl orange 3.0 - 4.4 Phenolphthalein 8.2 - 10.0 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.

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

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

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

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

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-

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?

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

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

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-]

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?

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

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.

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