[Author © Dr Phil Brown PhD: Doc Brown's chemistry exam revision notes on pH-reactions-acids-bases-neutralisation-salt preps, suitable for students of UK GCSE level and international IGCSE/O level chemistry courses, ~US grades 9-10 chemistry notes [page updated Mar 11th 2026 *]

Summary of this page

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. These revision notes and pH curve graphs for acid alkali titration neutralisations and their pH changes, should prove useful for the new AQA chemistry, Edexcel chemistry & OCR chemistry GCSE (9–1, 9-5 and 5-1) science courses.

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

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

 

sodium hydroxide + sulfuric acid ==> sodium sulfate + 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 i.e. understanding what a pH curve represents.

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.

be able to explain the shape of the graph i.e. why the pH changes the way it does!

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 (see section 10).

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.

You can do the experiments by measuring out a given volume of an acid/alkali of known concentration, and adding the acid/alkali in small portions from a burette and measuring the pH with a pH meter after every small addition e.g. 1 cm3 of the acid/alkali solution. See another page for the apparatus and how to use it.

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 i.e becoming a much more alkaline solution.

The pH titration curves involving weak acids and weak alkalis are explained in section 10.

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

Examples of more accurate and 'specialised' indicators

• 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/potassium hydroxide (NaOH/KOH) – hydrochloric/sulfuric acid (HCl/H₂SO₄) titrations,
  • the pH at the endpoint is around 7, so you can 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 AND with universal indicator you see green at the endpoint, but these two indicators are NOT as accurate as phenolphthalein and other indicators.
    • They are ok as simple demonstration to illustrate the principles of a titration, BUT, they should not be used for quantitative work.
• titrating a weak soluble base (weak alkali) with a strong acid:
  • e.g. for titrating ammonia (NH₃) with hydrochloric/sulfuric acid (HCl/H₂SO₄),
  • the end-point is typically around pH 5, so you have to use ...
  • methyl orange indicator (red in acid, yellowish–orange in neutral–acid), the end–point is an 'orange' colour, not easy to see accurately,
  • or, 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,
  • or, methyl red indicator (red in acid, yellow in neutral–alkaline), the end–point is 'orange'.
• titrating a weak acid with a strong alkali:
  • e.g. for titrating ethanoic acid (CH₃COOH) with sodium hydroxide (NaOH),
  • the end-point is typically around pH 9, so you have to use ...
  • phenolphthalein indicator (pink in alkali, colourless in acid–neutral solutions, pink in alkali), the end–point is the first permanent pink.
• 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.
• See section 10. for more details on acid-alkali titrations and calculations and procedures.