Acids and Bases (HL) - Lesson 12 Indicators

Refresh A 25.0 cm3 solution of a weak monoprotic acid, HA(aq), is titrated with 0.155 mol dm–3 sodium hydroxide, NaOH(aq), and the following graph is obtained. Determine the pH at the equivalence point. Explain, using an equation, why the equivalence point is not at pH = 7. Estimate, using data from the graph, the dissociation constant, Ka, of the weak acid, HA, showing your working.

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Lesson 11: Indicators Objectives: Understand how indicators work Determine suitable indicators for a reaction Make indicators from a range of natural products

Indicators An indicator is a compound whose colour depends on pH For example: Phenolphthalein pH< 8: COLOURLESS pH>10: PINK Methyl orange pH<3.2: RED pH>4.4: ORANGE Bromothymol blue: pH<6.0: YELLOW pH>7.6: BLUE Most indicators change colour only once (sometimes twice). The obvious exception is universal indicator which is actually a mixture of several indicators with different colour change ranges. Note: the pH generally changes over a small range, but the logarithmic nature of pH means often equates to a single drop.

How Indicators Work (what you need): Indicators are weak acids/bases in their own right In solution indicators form an equilibrium: In + H+ ⇌ InH+ COULOUR 1 COLOUR 2 Where: ‘In’ stands for indicator As [H+] changes, the equilibrium moves to the left or right, thus changing the colour The structure of indicators change depending on the pH: Higher pH can cause weak-acid groups to deprotonate Low pH can cause weak-base groups to protonate This can have knock-on effects on the structure, of the indicator molecule which changes its colour

How Indicators Work (in detail): Colour in many organic compounds comes from having overlapping π-systems, with many delocalised electrons For example: phenolphthalein The colour derives from conjugated π-systems which are radically altered by the changes in structure The double bond on the central carbon atom allows the π-systems in the three benzene rings to interact with each other, leading to the pink colour This is more detail than required by the IB!!!  + 2H+ pink colourless

pH range and pKa Most indicators change range within ±1.0 of their pKa pKa data for indicators can be found in your data booklet Indicator pKa pH Range Colour Change Acid Alkali methyl orange 3.46 3.2–4.4 Red Yellow bromophenol blue 4.10 3.0–4.6 Yellow Blue bromocresol green 4.90 3.8–5.4 methyl red 5.00 4.8–6.0 bromothymol blue 7.30 6.0–7.6 Yellow Blue phenol red 8.00 6.6–8.0 Yellow Red phenolphthalein 9.50 8.2–10.0 Colourless Pink

Choosing an indicator Check out the simulation here: http://chem- ilp.net/labTechniques/AcidBaseIdicatorSimulation.htm From the simulation you can see the indicator should be chosen such that the colour change happens as close as possible to the equivalence point of the titration

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