Choosing Acid-Base Titration Indicators
Key Concepts
- An acid-base indicator is a dye that changes colour when pH changes.
- An indicator is actually a Brønsted-Lowry conjugate acid-base pair in which the acid is a different colour to the base.
- The end point (end-point) of a titration is when the indicator changes colour during a titration.
- The equivalence point(1) of an acid-base reaction is when the amount of acid and of base is just sufficient to cause complete consumption of both the acid and the base.
⚛ At the equivalence point neither the acid nor the base is in excess.
⚛ At the equivalence point neither the acid nor the base is the limiting reagent.
- The pH of the solution at the equivalence point depends on the relative strength of the acid and strength of the base used in the titration.
For aqueous solutions at 25°C:
⚛ strong acid + weak base, resultant solution pH < 7 (acidic)
⚛ weak acid + strong base, resultant solution pH > 7 (basic)
⚛ strength of acid = strength of base, resultant solution pH = 7 (neutral)
- An appropriate indicator for an acid-base titration will change colour at the same pH as the equivalence point of the acid-base reaction.
That is, the end point (end-point) of the titration as indicated by the indicator must be the same as the equivalence point of the acid-base reaction.
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Choosing an appropriate indicator for a titration
In an acid-base titration, the following chemical reaction (Arrhenius neutralisation) takes place:
| acid |
+ |
base |
→ |
salt |
+ |
water |
| HA(aq) |
+ |
MOH(aq) |
→ |
MA(aq) |
+ |
H2O(l) |
When all the acid has reacted with all of the base the resultant solution is an aqueous salt solution, MA(aq).
The nature of this aqueous salt solution determines the pH of the resultant solution.
In general, we can apply the following generalisation for aqueous solutions at 25°C based on the relative strength of the acid and strength of the base used in the titration:
- If acid is stronger than base, salt solution has a pH < 7 (acidic)
- If acid is weaker than base, salt solution has a pH > 7 (basic)
- If acid strength is the same as base strength, salt solution has a pH = 7 (neutral)
An appropriate acid-base indicator will change colour at the equivalence point of the titration.
- Strong acid + strong base → salt (pH=7)
Choose an indicator that changes colour at pH=7
- Strong acid + weak base → salt (pH < 7)
Choose an indicator that changes colour at pH < 7
- Weak acid + strong base → salt (pH > 7)
Choose an indicator that changes colour at pH > 7
An appropriate indicator for an acid-base titration will change colour over a narrow pH range, and have distinctive colour at lower pH and a different, distinctive colour at higher pH.
The below shows the approximate colour of some acid-base indicators at different pH values and the type of titrations they are useful for:
| pH |
0 |
1 |
2 |
3 |
4 |
5 |
6 |
7 |
8 |
9 |
10 |
11 |
12 |
13 |
14 |
Use |
|---|
| cresol red |
< 0.2 |
|
> 1.8 |
strong acid + weak base |
|---|
thymol blue
(1st change) |
< 1.2 |
|
> 2.8 |
strong acid + weak base |
|---|
| bromophenol blue |
< 3.0 |
|
> 4.6 |
strong acid + weak base |
|---|
| methyl orange |
< 3.1 |
|
> 4.4 |
strong acid + weak base |
|---|
| bromocresol green |
< 3.8 |
|
> 5.4 |
strong acid + weak base |
|---|
| methyl red |
< 4.4 |
|
> 6.2 |
strong acid + weak base |
|---|
| litmus |
< 5.0 |
|
> 8.0 |
|
|---|
| bromothymol blue |
< 6.0 |
|
> 7.6 |
strong acid + strong base |
|---|
| phenol red |
< 6.8 |
|
> 8.4 |
strong acid + strong base |
|---|
thymol blue
(2nd change) |
< 8.0 |
|
> 9.6 |
weak acid + strong base |
|---|
| phenolphthalein |
< 8.3 |
|
> 10.0 |
weak acid + strong base |
|---|
Not all acid-base indicators are suitable for use in acid-base titrations:
- Litmus is not used in titrations because the pH range over which it changes colour is too great (pH range is 5.0 - 8.0) .
- Universal indicator, which is actually a mixture of several indicators, displays a variety of colours over a wide pH range so it can be used to determine an approximate pH of a solution but is not used for titrations.
There are two steps in deciding which indicator to use for a particular acid-base titration:
- Determine the pH of the solution at the equivalence point:
(a) You may be told the pH of the solution (eg, in an exam question).
(b) You may be able to approximate the pH of the salt solution using the relative strength of acid and base as shown above.
(c) You may be expected to calculate the pH of the solution. (see Calculating pH of Aqueous Salt Solutions).
(d) You may be given a titration curve to use to determine which indicator you would use (examples of this are shown in the next section).
- Use a table of indicator colour and pH range to choose an indicator which changes colour over a pH range that includes the equivalence point.
Examples of Choosing an Appropriate Indicator for Titrations
Choosing an Appropriate Indicator for a Strong Acid - Strong Base Titration
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 25°C.
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:
| pH | Bromothymol 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 < 6.0 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.
|
| pH | Phenol 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 point for the titration (pH=7) occurs within the pH range for the visible colour change of the indicator (the end point between pH 6.8 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).
Choosing an Appropriate Indicator for a Weak Acid - Strong Base Titration
An aqueous solution of acetic acid (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.
| pH | Thymol 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.
|
| pH | Phenolphthalein
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 magenta
- at the end point, between pH 8.3 and 10.0, the solution appears to be pale pink (an equimolar mixture of colourless and magenta)
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).
Choosing an Appropriate Indicator for a Strong Acid - Weak Base Titration
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)
| pH | Methyl 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.
Consequences of Using the Wrong Indicator in an Acid-Base Titration
In the discussion above, we decided that we could use bromothymol blue or phenol red as indicators for the titration of NaOH(aq) (a strong base) with HCl(aq) (a strong acid) because these indicators change colour over a range of pH values that includes the pH of NaCl(aq) (the salt produced in the neutralisation reaction):
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
At the equivalence point the only species in solution is NaCl(aq)
NaCl(aq) has pH=7
bromothymol blue changes colour between pH 6.0 and 7.6
phenol red changes colour between pH 6.8 and 8.4
What would happen if we used a different indicator instead?
What would happen if we used phenolphthalein?
Imagine we are adding NaOH(aq) to HCl(aq) in a conical flask.
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
Initially there is a large excess of acid, the solution is acidic, and the phenolphthalein indicator is colourless.
At the equivalence point the only species in solution is NaCl(aq) which has pH=7
BUT phenolphthalein changes colour between pH 8.3 and 10.0, so, at the equivalence point the phenolphthalein remains colourless.
We will have to add an excess of NaOH(aq) to the HCl(aq) to make phenolphthalein change colour, in other words, the end point as indicated by the indicator will occur AFTER the equivalence point for the acid-base reaction.
What would happen if we used methyl orange?
Imagine we are adding NaOH(aq) to HCl(aq) in a conical flask.
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
Initially there is a large excess of acid, the solution is acidic, and the methyl orange indicator is red.
As we slowly add NaOH (base) to the acid, the pH gradually increases.
When the pH increases to about 3.1, the colour of the indicator starts to look more orange than red so the end point of the titration as indicated by the indicator has been reached.
BUT the equivalence point of the titration will not occur until well after this colour change, at pH=7, so the end point occurs BEFORE the equivalence point.
It is essential that we choose an indicator that changes colour over a range that includes the pH of salt solution formed as a result of the neutralisation reaction (titration reaction).
Ideally: pH at the end point = pH at the equivalence point
| excess acid |
equivalence point moles H+ = moles OH- pH of salt solution |
excess base |
| low pH indicator colour |
end point (pH range) |
high pH indicator colour |
Problem Solving: Choosing an Acid-Base Indicator for a Titration
Question: Chris the Chemist has been asked to determine the concentration of acetic acid (ethanoic acid) in Gran's homemade apple cider vinegar.
Because acetic acid is a weak acid, Chris has decided to titrate it with an aqueous solution of sodium hydroxide because this is a strong base.
Chris has the following acid-base indicators currently available in the lab:
| Indicator name |
pH Range |
Colour change
(low pH → high pH) |
|---|
| methyl orange |
3.1 - 4.4 |
red → yellow |
| bromothymol blue |
6.0 - 7.6 |
yellow → blue |
| phenolphthalein |
8.3 - 10.0 |
colourless → pink |
Which indicator should Chris use for the titration?
Solution:
(using the StoPGoPS approach to problem solving)
| STOP |
STOP! State the Question.
|
| |
What is the question asking you to do?
Name the acid-base indicator to be used
|
| PAUSE |
PAUSE to Prepare a Game Plan
|
| |
(1) What information (data) have you been given in the question?
Acetic acid = a weak acid
Sodium hydroxide = a strong base
pH Range of indicators:
methyl orange: 3.1 - 4.4
bromothymol blue: 6.0 - 7.6
phenolphthalein: 8.3 - 10.0
(2) What is the relationship between what you know and what you need to find out?
(i) Decide the pH at the equivalence point of the titration:
For an acid-base titration, the pH of the final solution depends on the relative strength of the acid and the strength of the base:
acid stronger than base: pH(equivalence) < 7
strength of acid = strength of base: pH(equivalence) = 7
base stronger than acid: pH(equivalence) > 7
(ii) Decide on the pH range of the indicator and hence name the most suitable acid-base indicator to use:
Ideally, indicator's colour change at the end point should occur at the same pH as the equivalence point of the neutralisation reaction.
pH(equivalence) = pH(end point)
In practice, pH(equivalence) occurs within the pH range of the indicator:
pH(lower limit colour change) < pH(equivalence point) < pH(upper limit colour change)(
|
| GO |
GO with the Game Plan |
| |
(i) Decide the pH at the equivalence point of the titration:
Since acetic acid is a weak acid and sodium hydroxide is a strong base, that is, base is stronger than acid:
pH(equivalence) > 7
(ii) Decide on the pH range of the indicator and hence name the most suitable acid-base indicator to use:
pH(end point) = pH(equivalence)
Therefore: pH(end point) > 7
Phenolphthalein would be the best choice because its whole pH range is greater than 7, that is, its pH range is 8.3 - 10.0
|
| PAUSE |
PAUSE to Ponder Plausibility |
| |
Have you answered the question?
Yes, we have named one of the indicators given.
Is your answer plausible?
Work backwards: Assume we use phenolpthalein, what sort of acid and base would be used in the titration?
pH range (phenolphthalein) is 8.3 - 10.0
pH range > 7 therefore the base must be stronger than the acid.
Weak acid + strong base → salt + water
Strong bases: Hydroxides of Group 1 and 2 elements, so sodium hydroxide is a strong base (sodium is a Group 1 element)
Strong acids are: hydrohalic acids (except HF), sulfuric acid, nitric acid and perchloric acid, so acetic acid is a weak acid.
Since the relative strength of the acid and base determined by working backwards agrees with the information given in the question we are reasonably confident that our answer is plausible.
|
| STOP |
STOP! State the Solution |
| |
Chris should use the phenolphthalein indicator.
|
Footnotes:
(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.
You probably won't be using "equivalents" as a measure of quantity in your high school chemistry course, but it is useful to understand where the term "equivalence point" comes from.
