How to make methyl orange indicator
Doc Brown's Chemistry – Advanced Level Chemistry Notes
Theoretical–Physical Advanced Level Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 6.2
6.2 Indicator theory of acid–base titrations, pH curves & pKind
What is the theory behind how an acid–alkali indicator works? Why are some weak acids good indicators in acid–base titrations? What is the pH range of an indicator? How does the pH change throughout an weak/strong acid–soluble weak/strong base titration?
GCSE/IGCSE reversible reactions–equilibrium notes *
GCSE/IGCSE notes on acids and bases
6.2 Theory of acid–base titration indicators, pH curves and pKind selection of a suitable indicator and how pH changes when acids react with bases e.g. in a titration All of Equilibria part 5 "pH and weak–strong acids and bases in aqueous solution" should have been studied prior to tackling Part 6. 6.2.1 Acid–base titration indicators are quite often weak acids in which the unionised acid (lets call it HIn ) and its 'de–protonated' form, or conjugate base, the anion (In – ), have different colours. One form can be colourless e.g. phenolphthalein in acid–neutral solutions. The equilibrium can be simply expressed as. HIn – (aq, colour 1)
H + (aq) + In – (aq, colour 2) Applying Le Chatelier's equilibrium principle : Addition of acid favours the formation of more HIn (colour 1) HIn(aq)
H + (aq) + In – (aq) because an increase on the right of [H + ] causes a shift to left increasing [HIn] to minimise 'enforced' rise in [H + ]. Addition of alkali favours the formation of more I – (colour 2): The increase in [OH – ] causes a shift to right because the reaction reduces the [H + ] on the right so more HIn ionises to try to increase the [H + ] i.e. minimising the change in [H + ].
6.2.2 The colour that is observed will depend on the ratio [HIn] / [In – ]. but at pH extremes i.e. very acid, colour 1 will dominate, or in very alkaline solution, colour 2 will dominate. Therefore the maximum colour 'shade' change from one to the other will occur when [HIn] = [In – ]. or [colour 1] = [colour 2]. The pH when [HIn] = [In – ] can be calculated from the dissociation constant, Kind (ka ), for the weak acid indicator. Kind = [H + (aq) ] [In – (aq) ] / [HIn(aq) ]. but when [HIn] = [In – ] the equilibrium expression simplifies to. Kind = [H + (aq) ], so at this point the pH = –log(Kind ) and is referred to as the pKind value. 6.2.3 pH titration curves and choice of indicator A simple pH curve for an acid–alkali is explained on one of my GCSE acid–alkali notes pages and is well worth reading first, before tackling all the possibilities described and explained below. The greatest change in indicator colour (per volume of reagent added), will occur at the equivalence point in the titration. Therefore you need to choose an indicator with a pKind close to the pH at the equivalence point (theory above). In fact acid–base titration indicators are usually effective over a range of several pH units but it is essential for accurate titrations that the colour change is sharp at the equivalence point with a small addition of acidic or alkaline titration solution. Universal indicator is NOT suitable for quantitative analysis and the indicator choices tabulated below are explained via the sets of pH graphs shown further down using Graphs A to D. The effective pH range of the indicator is where there is sufficient colour change to give a good sharp end–point and can be above or below, but close to the pKind value of the indicator. Some effective pH ranges for selected indicators are given below.
Indicator colour change, from acid to alkali