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Indicators

Key Concepts

An indicator is a dye than changes colour when pH changes
An indicator is actually a Brönsted-Lowry conjugate acid-base pair in which the acid is a different colour to the base

Examples of Indicators:

pH	0	1	2	3	4	5	6	7	8	9	10	11	12	13	14	pH range
	Acid							neutral	Base
[H⁺] (mol/L)	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
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
phenol- phthalein 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
bromo- thymol 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
bromo- cresol 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
bromo- phenol 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.
Litmus is not used in titrations because the pH range over which it changes colour is too great.
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.
Determine what species are present at the equivalence point & deduce the pH at the equivalence point

pH of salts formed from
reactions of acids & bases Strong Base Weak Base

Strong Acid pH = 7 pH < 7

Weak Acid pH > 7 pH = 7
Use the table of indicators to choose an indicator which changes colour over a pH range that includes the equivalence point

Examples:

Strong Acid - Strong Base titration

HCl(aq) + NaOH(aq) -----> NaCl(aq) + H₂O(l)
At equivalence the only species present will be NaCl(aq) & H₂O(l)
The solution of a salt of a strong acid and a strong base will have a pH=7
NaCl(aq) will have a pH=7
A suitable indicator would be bromothymol blue (pH range 6.2 - 7.6) or phenol red (pH range 6.8 - 8.4)

Strong Acid - Weak Base titration

HCl(aq) + NH₃(aq) -----> NH₄Cl(aq)
NH₄Cl is the salt of a strong acid & a weak base, so a solution of NH₄Cl will have a pH < 7 (NH₄⁺ is a weak acid)
A suitable indicator would be methyl orange (pH range 3.1 - 4.4) or methyl red (pH range 4.4 - 6.0)

Weak Acid - Strong Base titration

CH₃COOH(aq) + NaOH(aq) -----> CH₃COONa(aq) + H₂O(l)
CH₃COONa is the salt of a weak acid & a strong base, so a solution of CH₃COONa will have a pH > 7 (CH₃COO^- is a weak base)
A suitable indicator would be phenolphthalein (pH range 8.3 - 10.0) or thymol blue (pH 8.0 - 9.6)

Dissociation constants for indicators:

let HIn be the acid form of the indicator and In^- the base form and K_In the dissociation constant for the indicator

HIn + H₂O

In^- + H₃O⁺

K_In = [In^-][H₃O⁺]
[HIn]

K_In
[H₃O⁺] = [In^-]
[HIn]

The colour of the indicator at any pH is determined by the ratio [In^-]:[HIn]

As pH increases, [H₃O⁺] decreases, [In^-]:[HIn] increases, so the indicator has increasingly more of the colour of the base form and less of the acid form

It is important to use as little indicator as possible during a titration since the indicator itself reacts with the titration's reactants