This document discusses acid-base titrations. It explains that titrations can be used to determine the concentration of acids and bases in samples. The titrant is typically a strong acid or base while the sample species can be either a strong or weak acid or base. The document also discusses titration curves, the effects of concentration on the curves, and indicators used to detect the endpoint in titrations.

1. 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

2. 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.
Plot of ml titrant (or
pH
% titration) vs pH will
result in a typical
titration curve.
mmoles acid - mmoles base
% titration
total volume

3. mmoles excess
total volume
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
(100 ml)(0.10 M) - (10 ml)(0.10 M) 60 160 0.025 1.60
100 ml + 10 ml 70 170 0.018 1.74
80 180 0.011 1.96
90 190 0.0053 2.28
2.5
2
1.5
pH
1
0.5
0
0 30 60 90
ml NaOH

4. 7
6
5
pH
4
3
2 10 ml
210 ml
1
0
0 10 20 30 40 50 60 70 80 90 100
ml NaOH
ml titrant total V [OH-] pH 14
110 210 0.0048 11.68 12
120 220 0.0091 11.96 10
130 230 0.013 12.11 8
140 240 0.017 12.23
pH
150 250 0.020 12.30 6
160 260 0.023 12.36 4
170 270 0.026 12.41 2
180 280 0.029 12.46
0
190 290 0.031 12.49
0
20
40
60
80
100
120
140
160
180
200
200 300 0.033 12.52
ml NaOH
basic sample
pH
acidic sample
ml titrant

5. The concentration of either our sample or titrant We get a similar effect as the concentration of
can affect the shape of our titration curve. the titrant is reduced.
different acid
sample This is one of the
As the [acid] decreases, concentrations
we get a less distinct
reasons that most
jump in pH. strong acid-base
titrations are done
in the 0.5-0.1 M
range.
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
acid base involved and then work through an example.
[H3O+][A-] [OH-][HA]
[HA] [A-]

6. [A-]
[HA]
[HA]
[A-]
% titration
100 - % titration

7. [H3O+][A-]
[HA]
x2
0.10
% titration
100 - % titration
% titration pH 6
0 2.60 Note:
5
At 50% titration,
10 3.24
pH = pKA
20 3.60 4
pH
30 3.83 Also, the was only 3
40 4.02 a change of 1.91 2
50 4.20 pH units as we
60 4.38 went from 10 to 1
70 4.57 90 % titration.
0
80 4.80 0 20 40 60 80 100
90 5.15 % titration

8. [OH-][HA]
[A-]
9
8
x2 7
0.050 6
pH
5
4
3
2
1
0
0 20 40 60 80 100
% titration
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

9. 14
12
10
8
pH
6
4
2
0
0 50 100 150
% titration
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. eq. pt.
pH
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.
% titration
Normal
titration
curve
pH
Too much
indicator is
present
% titration

10. 1
10
10
1
pH transition
Indicator range color Phenophthalein
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
Methyl Red Bromothymol blue

11. [A-]
[HA]
Initially, each solution is at pH 7.00.
After adding 10 ml of 1.0 M HCl we have:
Pure water
(10 ml)(1.0 M)
[H3O+] = = 0.091
(110 ml)
pH = 1.04
This is a pretty big jump!

12. Addition of 10 ml 1.0 M HCl to our buffered
system.
We started with 0.10 moles of both the acid [A-]
and conjugate base forms. [HA]
The addition of our first 10 ml can be 0.09 mmol
expected to react the the conjugate base, 1.1 mmol
converting it to the acid.
After addition, there are 0.09 moles of the
base form and 0.11 moles of the acid form.
7
ml HCl pH
added unbuffered buffered 6
0 7.00 7.00 5
10 1.04 6.91
4
buffered
pH
20 0.78 6.82
30 0.64 6.73 3
unbuffered
40 0.54 6.63 2
50 0.48 6.52
1
60 0.43 6.40
70 0.39 6.25 0
0 20 40 60 80 100
80 0.35 6.05
90 0.32 5.72 ml HCl added
For the addition of 100 ml of HCl, we have
converted virtually all of A- to HA so the
calculation is different.
[H+][A-]
KA = 1.00 x 10-7 =
[HA]
When we account for dilution,
1.0 M = [HA] + [A-]
where [A -] is negligible.
This is a standard weak acid calculation.

13. 7 Obviously, a buffer only has a limit ability to
6 reduce pH changes.
5
buffered
4 The pKA determines the range where a buffer
pH
3
unbuffered is useful.
2
1 The concentration of our buffer system
0 determines how much acid or base it can deal
0 20 40 60 80 100 with.
ml HCl added
[A-]
[HA]
[A-]
[HA]
Assume that [ HA ] = [ A- ] = C, where C is the Concentration, M Buffer capacity (mol)
initial concentration of either the acid or base 1.0 0.22
Buffer capacity in general is then 0.50 0.11
0.10 0.022
1 = log [A-]
[HA] We’ll only worry 0.050 0.011
[ A- ] = 10 [ HA ] about addition 0.010 0.0022
of a base.
[A-] = 2 C - [ HA ]
The capacity may be smaller if you don’t start
[ HA ] = 0.22 C with a 1:1 mixture.