pH changes in acid–alkali neutralisations & indicator choice
Doc Brown's Chemistry GCSE/IGCSE Science–Chemistry Revision Notes
pH scale of acidity and alkalinity, acids, bases–alkalis, salts and neutralisation
7. The pH changes in a neutralisation reaction and choice of indicator for a titration
pH titration curves for a neutralisation reaction: How does the pH change during a neutralisation reaction? What indicator do you use for a particular acid–alkali titration? Litmus, phenolphthalein, methyl orange, methyl red colour changes are given for determining the end–point of a titration.
GCSE/IGCSE Acids & Alkalis revision notes sub–index: Index of all pH, Acids, Alkalis, Salts Notes 1. Examples of everyday acids, alkalis, salts, pH of solution, hazard warning signs : 2. pH scale, indicators, ionic theory of acids–alkali neutralisation : 4. Reactions of acids with metals/oxides/hydroxides/carbonates, neutralisation reactions : 5. Reactions of bases–alkalis like ammonia & sodium hydroxide : 6. Four methods of making salts : 7. Changes in pH in a neutralisation, choice and use of indicators : 8. Important formulae of compounds, salt solubility and water of crystallisation : 10. More on Acid–Base Theory and Weak and Strong Acids
7. What pH changes go on in a neutralisation reaction?
Typical neutralisation reactions involving mixing a soluble acid with a soluble base (alkali) include
These two graphs, on the same set of axis, show how the pH changes when an alkali (soluble base) and an acid neutralise each other and what you see visually using universal indicator (univ. ind.). These simple curves represent what happens when eg hydrochloric acid and sodium hydroxide are mixed or nitric acid and potassium hydroxide (1 : 1 molar equations) BUT the curves are complicated for acids like sulfuric acid where the molar ratio is NOT a 1 : 1 molar ratio with the alkali. You need to consult more advanced notes via links at the end of the page.
Strictly speaking, they only apply to a strong acid and strong soluble base (alkali), but this pattern of pH change illustrated by the graph is what is happening in the salt preparation method (a) or in acid and alkali titrations.
Note: You can prepare a salt, or analyse an acid or alkaline solution by doing an acid–alkali addition either way round but in either case the volume of acid or alkali needed for neutralisation = the volume reading X at pH 7 (univ. ind. green). This strictly speaking only applies if it is a strong acid reacting with a strong alkali. At first on adding one to the other, the pH only changes gradually, but then you get a much more dramatic change as you approach the end-point i.e. the complete neutralisation point. It is at this point, the end-point, you get the sharpest change in pH, and, any indicator you choose to use, MUST change colour the most sharply at this point on the pH curve, to get the end-point accurately, hence the titration volume of acid or alkali accurately.
Red graph line: If you add a strong acid to a strong alkali (univ. ind. = blue), the pH starts at about 13-14 and only falls little at first as the colour changes from purple ==> blue. Then the pH falls much more steeply as the indicator colour changes from 'bluey' green ==> dark green ==> pale green. The solution is then neutralised at pH 7. This is the point where the salt is 100% formed. With further addition of excess acid, the pH falls and then levels out to about pH 1 as the colour changes further from green ==> yellow ==> orange ==> red
In terms of H+ and OH– ions: Initially a high concentration of OH–, so solution very alkaline, but as the H+ is steadily added, the OH– ions are neutralised to water. Therefore the OH– concentration steadily falls as does the pH because the solution becomes less alkaline. At pH 7, neutral there are very tiny equal concentrations of H+ and OH–. If excess acid is added, the pH steadily falls to around 1 as the concentration of H+ from the acid rises.
Blue graph line: If you add a strong alkali to a strong acid (univ. ind. = red), the pH starts at about 0-1 and only rises a little at first with the colour still quite red. Then on further addition of alkali the pH rises more sharply as the colour changes from red ==> orange ==> yellow and eventually at the neutralisation point at pH 7 the univ. ind. is green. This is the point where the salt is 100% formed. With excess alkali the pH continues to rise and then levels out to about 13 as the indicator colour changes through dark green ==> blue ==> purple.
In terms of H+ and OH– ions: Initially a high concentration of H+, very acid, but as the OH– of the alkali is steadily added, the H+ ions are neutralised to water. Therefore the H+ concentration steadily falls and the pH rises as the solution becomes less acid. At pH 7, neutral there are very tiny equal concentrations of H+ and OH–. If excess alkali is added the pH steadily rises from 7 to around 13 as the concentration of OH– from the alkali rises ie becoming a much more alkaline solution.
Universal indicator, and most other acid–base indicators, work for strong acid and alkali titrations, but universal indicator is a somewhat crude indicator for other acid–alkali titrations because it gives such a range of colours for different pH's. So, to get accurate titration results you need to use a special indicator for a particular acid–alkali titration. The complications arise because not all acids and soluble bases (alkalis) are as strong as each other. There is more on weak/strong acids in section 10. More on Acid–Base Theory and Weak and Strong Acids
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