pH scale of acidity and alkalinity, acids, bases-alkalis, salts and neutralisation

7. The pH changes in a neutralisation reaction

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 Sub-index: Index of all pH, Acids, Alkalis, Salts Notes 1. Examples of acid-alkali chemistry : 2. pH scale, indicators, ionic theory of acids-alkali neutralisation : 3. pH examples of acid, neutral or alkaline solutions : 4. Acid reactions with metals/oxides/hydroxides/carbonates and neutralisation reactions : 5. Reactions of bases-alkalis like sodium hydroxide : 6. Four methods of making salts : 7. Changes in pH in a neutralisation : 8. Important formulae, salt solubility and water of crystallisation : 9. Further examples of word/symbol equations for salt preparations : 10. More on Acid-Base Theory and Weak and Strong Acids

sodium hydroxide + hydrochloric acid ==> sodium chloride + water

NaOH_(aq) + HCl_(aq) ==> NaCl_(aq) + H₂O_(l)

sodium hydroxide + sulphuric acid ==> sodium sulphate + water

2NaOH_(aq) + H₂SO_4(aq) ==> Na₂SO_4(aq) + 2H₂O_(l)

potassium hydroxide + nitric acid ==> potassium nitrate + water

KOH_(aq) + HNO_3(aq) ==> KNO_3(aq) + 2H₂O_(l)

Apart from the water, all the actual species are ions eg

H⁺, Cl^- and SO₄^2- ions from the acids, Na⁺, K⁺ and OH^- from the alkali.

The only ions that change are

the hydrogen ion (H⁺ causes acidity), if pH <7 acid, H⁺ concentration is more than OH^- concentration

the hydroxide ion (OH^- causes alkalinity), if pH >7 alkaline, OH^- concentration is more than H⁺ concentration

these two ions combine to form neutral water in these particular neutralisation reactions -

this is the real neutralisation reaction

H⁺_(aq)  + OH^-_(aq)  ==> H₂O_(l)

Note: At pH 7 there are very tiny, but equal, concentrations of H^- and OH^- ions.

This means the Cl^-, SO₄^2-, NO₃^-, Na⁺ and K⁺ ions do NOT change and are called spectator ions and their concentrations only fall on mixing due to the obvious dilution effect of having a bigger volume they are dissolved in.

These are important chemical points when appreciating what is going on when the pH of a solution changes ie understanding what a pH curve represents.

The graphs 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.

The curves

This what is happening in the salt preparation method (a) above.

Note: you can prepare a salt by doing the 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).

Red graph line: If you add acid to an alkali (univ. ind. = blue), the pH starts at about 13 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 alkali to an acid (univ. ind. = red), the pH starts at about 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.

Examples of more accurate and 'specialised' indicators are:

• Apparatus used in titrations - pipette, conical flask and a burette
• Note that the first mentioned is in the flask and the second is in the burette.
• titrating a strong alkali with a strong acid (or vice versa):
  • e.g. for sodium hydroxide (NaOH) - hydrochloric/sulphuric acid (HCl/H₂SO₄) titrations, use ...
  • phenolphthalein indicator (pink in alkali, colourless in acid-neutral solutions), the end-point is the pink <==> colourless change.
  • Litmus works too, the end point is the red <==> purple/blue colour change.
• titrating a weak alkali with a strong acid:
  • e.g. for titrating ammonia (NH₃) with hydrochloric/sulfuric acid (HCl/H₂SO₄), use ...
  • methyl orange indicator (red in acid, yellowish-orange in neutral-acid), the end-point is an 'orange' colour, not easy to see accurately.
  • screened methyl orange indicator is a slightly different dye-indicator mixture that is reckoned to be easier to see than methyl orange, the end-point is a sort of 'greyish orange', but still not easy to do accurately.
• titrating a weak acid with a strong alkali:
  • e.g. for titrating ethanoic acid (CH₃COOH) with sodium hydroxide (NaOH), use ...
  • phenolphthalein indicator (pink in alkali, colourless in acid-neutral solutions, pink in alkali), the end-point is the first permanent pink.
  • methyl red indicator (red in acid, yellow in neutral-alkaline), the end-point is 'orange'.
• titrating a weak acid with a weak alkali (or vice versa):
  • These are NOT practical titrations because the pH changes at the end-point are not great enough to give a sharp colour change with any indicator.

•  Section (2) lists common indicators
• Acid-alkali titration calculations for GCSE/IGCSE students and how to carry out a titration
• Further theory, and examples of strong/weak acids/alkalis (soluble bases) are described in section 10
• Advanced level theory of indicators and titrations
• and Advanced Level acid-alkali titration questions

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