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Acid-Base Titration Indicators

Key Concepts

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Indicator Colours

The table below gives the approximate colour of a number of different acid-base indicators in aqueous solutions of varying pH at 25°C:

25oC Acid neutral Base  
[H+(aq)]
(mol L-1)
100 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-11 10-12 10-13 10-14  
pH 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH range
Universal indicator red red orange-
red
orange pale
orange
orange-
yellow
pale
yellow
green-
yellow
green dark-
green
blue blue blue blue blue  
cyanidin
(red cabbage water)
red red red cerise purple blue blue blue aqua-
marine
emerald-
green
lime lime yellow yellow yellow  
blue litmus indicator red red red red red red red blue blue blue blue blue blue blue blue 5.0 - 8.0
red litmus indicator red red red red red red red red blue blue blue blue blue blue blue 5.0 - 8.0
phenolphthalein indicator colour
-less
colour
-less
colour
-less
colour
-less
colour
-less
colour
-less
colour
-less
colour
-less
pink pink pink pink pink pink pink 8.3 - 10.0
thymol blue indicator yellow yellow yellow yellow yellow yellow yellow yellow yellow blue blue blue blue blue blue 8.0 - 9.6
phenol red indicator yellow yellow yellow yellow yellow yellow yellow yellow red red red red red red red 6.8 - 8.4
bromothymol blue indicator yellow yellow yellow yellow yellow yellow yellow blue blue blue blue blue blue blue blue 6.2 - 7.6
methyl red indicator pink pink pink pink pink pink yellow yellow yellow yellow yellow yellow yellow yellow yellow 4.4 - 6.0
bromocresol green indicator yellow yellow yellow yellow yellow pale blue-
green
blue-
green
blue-
green
blue-
green
blue-
green
blue-
green
blue-
green
blue-
green
blue-
green
blue-
green
3.8 - 5.4
methyl orange indicator red red red red yellow yellow yellow yellow yellow yellow yellow yellow yellow yellow yellow 3.1 - 4.4
bromophenol blue yellow yellow yellow yellow blue blue blue blue blue blue blue blue blue blue blue 3.0 - 4.6
cresol red red red red yellow yellow yellow yellow yellow yellow yellow yellow yellow yellow yellow yellow 0.2 - 1.8
pH 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH range

Choosing an appropriate indicator for a titration:

An appropriate indicator will change colour at the equivalence point of the titration.
Some acid-base indicators are not suitable for use in titrations:

In order to select an appropriate indicator to use for a given acid-base titration:

  1. Determine what species are present at the equivalence point and deduce the approximate pH at the equivalence point.
    The table below gives you some idea of the relative pH at the equivalence point for acid-base titrations based on the relative strength of the acids and strength of the bases2:

    pH of salts formed from
    reactions of acids & bases
    (25oC)
    Strong Base Weak Base
    Strong Acid pH = 7 pH < 7
    Weak Acid pH > 7 pH ≈ 73

  2. Use the table of indicators above to choose an indicator which changes colour over a pH range that includes the equivalence point

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Indicators for Strong Acid - Strong Base Titrations

An aqueous solution of hydrochloric acid, HCl(aq), is a strong acid.
An aqueous solution of sodium hydroxide, NaOH(aq), is a strong base.
The balanced chemical equation below represents the neutralisation reaction between HCl(aq) and NaOH(aq):
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

At the equivalence point of the neutralisation reaction the only species present will be NaCl(aq) and H2O(l)
The aqueous solution of a salt of a strong acid and a strong base will have a pH = 7 at 25oC.
NaCl(aq) will have a pH = 7
Consider bromothymol blue (pH range 6.2 - 7.6) and phenol red (pH range 6.8 - 8.4) as possible indicators for this neutralisation reaction:

pHBromothymol Blue

mL strong acid added to strong base
The titration curve shown in orange shows the changes in pH that occur as HCl(aq) is added to NaOH(aq).

The equivalence point for the reaction is represented by the blue line at pH = 7

The background colour represents the colour of the solution containing the bromothymol blue indicator over the same range of pH values.

  • pH < 6.2 the solution appears to be yellow
  • pH > 7.6 the solution appears to be blue
  • at the end point, between pH 6.2 and 7.6, the solution appears to be green (an equimolar mixture of blue and yellow)
Since the equivalence point for the titration (pH = 7) occurs within the pH range for the visible colour change of the indicator (the end point between pH 6.2 and 7.6), this indicator can be used for this titration.

pHPhenol Red

mL strong acid added to strong base
The titration curve shown in orange shows the changes in pH that occur as HCl(aq) is added to NaOH(aq).

The equivalence point for the reaction is represented by the blue line at pH = 7

The background colour represents the colour of the solution containing the phenol red indicator over the same range of pH values.

  • pH < 6.8 the solution appears to be yellow
  • pH > 8.4 the solution appears to be red
  • at the end point, between pH 6.8 and 8.4, the solution appears to be orange (an equimolar mixture of red and yellow)

Since the equivalence poin-t for the titration (pH = 7) occurs within the pH range for the visible colour change of the indicator (the end point between pH 6.8 and 8.4), this indicator can be used for this titration.

A suitable indicator for this strong acid - strong base titration would be bromothymol blue (pH range 6.2 - 7.6) or phenol red (pH range 6.8 - 8.4).

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Indicators for Weak Acid - Strong Base Titrations

An aqueous solution of acetic (ethanoic) acid, CH3COOH(aq), is a weak acid.
An aqueous solution of sodium hydroxide, NaOH(aq), is a strong base.
Below is the balanced chemical reaction for the reaction between CH3COOH(aq) and NaOH(aq):
CH3COOH(aq) + NaOH(aq) → CH3COONa(aq) + H2O(l)

At the equivalence point CH3COONa(aq), the salt of a weak acid and a strong base, is present so a solution of CH3COONa will have a pH > 7 (CH3COO- is a weak base)
Consider thymol blue (pH range 8.0 - 9.6) or phenolphthalein (8.3 - 10.0) as suitable indicators.

pHThymol blue

mL weak acid added to strong base
The titration curve shown in orange shows the changes in pH that occur as CH3COOH(aq) is added to NaOH(aq).

The equivalence point for the reaction is represented by the blue line at pH = 8.7

The background colour represents the colour of the solution containing the indicator over the same range of pH values.

  • pH < 8.0 the solution appears to be yellow
  • pH > 9.6 the solution appears to be blue
  • at the end point, between pH 8.0 and 9.6, the solution appears to be green (an equimolar mixture of yellow and blue)

Since the equivalence point for the titration (pH = 8.7) occurs within the pH range for the visible colour change of the indicator (the end point between pH 8.0 and 9.6), this indicator can be used for this titration.

pHPhenolphthalein

mL weak acid added to strong base
The titration curve shown in orange shows the changes in pH that occur as CH3COOH(aq) is added to NaOH(aq).

The equivalence point for the reaction is represented by the blue line at pH = 8.7

The background colour represents the colour of the solution containing the phenolphthalein indicator over the same range of pH values.

  • pH < 8.3 the solution appears to be colourless
  • pH > 10.0 the solution appears to be pink
  • at the end point, between pH 8.3 and 10.0, the solution appears to be pale pink (an equimolar mixture of colourless and pink)

Since the equivalence point for the titration (pH = 8.7) occurs within the pH range for the visible colour change of the indicator (the end point between pH 8.3 and 10.0), this indicator can be used for this titration.

A suitable indicator for the titration of the weak acid CH3COOH(aq) and the strong base NaOH(aq) would be either thymol blue (pH range 8.0 - 9.6) or phenolphthalein (pH range 8.3 - 10.0).

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Indicators for Strong Acid - Weak Base Titrations

An aqueous solution of hydrochloric acid, HCl(aq), is a strong acid.
An aqueous solution of ammonia, NH3(aq), is a weak base.
The balanced chemical reaction below represents the reaction between HCl(aq) and NH3(aq):
HCl(aq) + NH3(aq) → NH4Cl(aq)

NH4Cl is the salt of a strong acid and a weak base, so a solution of NH4Cl will have a pH < 7 (NH4+ is a weak acid)
A suitable indicator would be methyl red (pH range 4.4 - 6.0)

pHMethyl Red

mL strong acid added to weak base
The titration curve shown in orange shows the changes in pH that occur as HCl(aq) is added to NH3(aq).

The equivalence point for the reaction is represented by the blue line at pH = 5.28

The background colour represents the colour of the solution containing the methyl red indicator over the same range of pH values.

  • pH < 4.4 the solution appears to be pink
  • pH > 6.0 the solution appears to be yellow
  • at the end point, between pH 4.4 and 6.0, the solution appears to be orange (an equimolar mixture of pink and yellow)

Since the equivalence point for the titration (pH = 5.28) occurs within the pH range for the visible colour change of the indicator (the end point between pH 4.4 and 6.0), this indicator can be used for this titration.

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1. 1 equivalent of an acid is the quantity of that acid which will donate 1 mole of H+.
1 equivalent of a base is the quantity which supplies 1 mole of OH-.
At the equivalence point, 1 equivalent of acid neutralises 1 equivalent of base.

2. The pH at the equivalence can be determined from the titration curve, or by calculating the pH of the salt solution as a result of hydrolysis.

3. The pH at the equivalence point of a weak acid-weak base titration would depend on the relative weakness of the acids and bases used.
If the weak acid is a stronger electrolyte than the weak base, the solution will be acidic, if the weak base is a stronger electrolyte than the weak acid then the solution will be basic.
If the strength of the weak acid and the weak base are similar then the pH of the solution will be 7.
In general, you are more likely to use a back titration (indirect titration) method rather than a direct titration method so you would not need an indicator for a weak acid-weak base titration.