This is a quick and accurate method for determining acidic or basic substances in many samples.
Several inorganic acids and bases.
Hundreds of organic species.
The titrant is typically a strong acid or base.
The sample species can be either a strong or weak acid or base.
COOH COO- K+
O COOH
As a last resort, you can standardize an HCl solution with a standard NaOH solution.
Standardization verses a secondary standard is not recommended in most cases.
Each standardization introduces an error so your results are less reliable if NaOH is used.
While chemical indicators can be used for endpoint detection, a pH electrode is the best way to monitor an acid base titration.
pH
% titration Plot of ml titrant (or
% titration) vs pH will result in a typical titration curve.
mmoles acid - mmoles base total volume
mmoles excess total volume
(100 ml)(0.10 M) - (10 ml)(0.10 M) 100 ml + 10 ml
ml titrant total ml [H3O+] pH 0 100 0.10 1.00 10 110 0.082 1.09 20 120 0.067 1.17 30 130 0.054 1.28 40 140 0.043 1.37 50 150 0.033 1.48 60 160 0.025 1.60 70 170 0.018 1.74 80 180 0.011 1.96 90 190 0.0053 2.28
0 0.5 1 1.5 2 2.5
0 30 60 90 ml NaOH
pH
0 1 2 3 4 5 6 7
0 10 20 30 40 50 60 70 80 90 100
ml NaOH
pH
10 ml 210 ml
ml titrant total V [OH-] pH 110 210 0.0048 11.68 120 220 0.0091 11.96 130 230 0.013 12.11 140 240 0.017 12.23 150 250 0.020 12.30 160 260 0.023 12.36 170 270 0.026 12.41 180 280 0.029 12.46 190 290 0.031 12.49 200 300 0.033 12.52
0 2 4 6 8 10 12 14
0 20 40 60 80 100 120 140 160 180 200
ml NaOH
pH pH
ml titrant basic sample
acidic sample
The concentration of either our sample or titrant can affect the shape of our titration curve.
different acid sample concentrations As the [acid] decreases,
we get a less distinct jump in pH.
We get a similar effect as the concentration of the titrant is reduced.
This is one of the reasons that most strong acid-base titrations are done in the 0.5-0.1 M range.
acid base
First, we’ll only be concerned about the titration of a weak acid with a strong base or a weak base with a strong acid.
! We still have the same four general regions for our titration curve.
! The calculation will require that you use the appropriate KA or KB relationship.
! We’ll start by reviewing the type of calculations involved and then work through an example.
[H3O+][A-] [HA]
[OH-][HA]
[A-]
[HA]
[A-] [A-] [HA]
% titration 100 - % titration
[H3O+][A-] [HA]
x 2 0.10
% titration 100 - % titration
% titration pH
0 2.60
10 3.24
20 3.60
30 3.83
40 4.02
50 4.20
60 4.38
70 4.57
80 4.80
90 5.15
Note:
At 50% titration, pH = pKA
Also, the was only a change of 1.91 pH units as we went from 10 to 90 % titration.
0 1 2 3 4 5 6
0 20 40 60 80 100
% titration
pH
[OH-][HA]
[A-]
x 2 0.050
0 1 2 3 4 5 6 7 8 9
0 20 40 60 80 100
% titration
pH
ml total
titrant volume [OH-] pH 110 210 0.0048 11.68 120 220 0.0091 11.96 130 230 0.013 12.11 140 240 0.017 12.23 150 250 0.020 12.30 This is identical to what we obtained for our
strong acid/strong base example
0 2 4 6 8 10 12 14
0 50 100 150
% titration
pH
During an acid-base titration, the indicator acts as an additional weak acid or base.
It must be weaker than the species being determined - titrated after analyte.
It must be present at relatively low
concentrations so as not to interfere with the normal titration curve and equivalence point.
It must give a sharp and distinct color change.
pH eq. pt.
% titration
pH
% titration Normal
titration curve
Too much indicator is present
1 10 10 1
pH transition
Indicator range color Bromophenol Blue 6.2 - 7.6 yellow - blue Methyl Orange 3.1 - 4.4 red - orange Methyl Red 4.2 - 6.2 red - yellow Bromothymol Blue 6.2 - 7.6 yellow - blue Cresol Purple 7.6 - 9.2 yellow - purple Phenolphthalein 8.3 – 10 colorless - red Thymolphthaleine 9.3 - 10.5 colorless - blue Alizerin Yellow GG 10 – 12 yellow - red
Phenophthalein
Methyl Red Bromothymol blue
[A-] [HA]
Initially, each solution is at pH 7.00.
After adding 10 ml of 1.0 M HCl we have:
Pure water
[H3O+] = = 0.091
pH = 1.04
This is a pretty big jump!
(10 ml)(1.0 M) (110 ml)
Addition of 10 ml 1.0 M HCl to our buffered system.
We started with 0.10 moles of both the acid and conjugate base forms.
The addition of our first 10 ml can be expected to react the the conjugate base, converting it to the acid.
After addition, there are 0.09 moles of the base form and 0.11 moles of the acid form.
[A-] [HA]
0.09 mmol 1.1 mmol
ml HCl pH
added unbuffered buffered
0 7.00 7.00
10 1.04 6.91
20 0.78 6.82
30 0.64 6.73
40 0.54 6.63
50 0.48 6.52
60 0.43 6.40
70 0.39 6.25
80 0.35 6.05
90 0.32 5.72
0 1 2 3 4 5 6 7
0 20 40 60 80 100
ml HCl added
pH
buffered unbuffered
For the addition of 100 ml of HCl, we have converted virtually all of A- to HA so the calculation is different.
KA = 1.00 x 10-7 =
When we account for dilution, 1.0 M = [HA] + [A-] where [A-] is negligible.
This is a standard weak acid calculation.
[H+][A-] [HA]
0 1 2 3 4 5 6 7
0 20 40 60 80 100
ml HCl added
pH
buffered unbuffered
Obviously, a buffer only has a limit ability to reduce pH changes.
The pKA determines the range where a buffer is useful.
The concentration of our buffer system determines how much acid or base it can deal with.
[A-] [HA]
[A-] [HA]
Assume that [ HA ] = [ A- ] = C, where C is the initial concentration of either the acid or base Buffer capacity in general is then
1 = log
[ A- ] = 10 [ HA ] [A-] = 2 C - [ HA ] [ HA ] = 0.22 C
[A-]
[HA] We’ll only worry
about addition of a base.
Concentration, M Buffer capacity (mol)
1.0 0.22
0.50 0.11
0.10 0.022
0.050 0.011
0.010 0.0022
The capacity may be smaller if you don’t start with a 1:1 mixture.
