TITRATION CURVES
The way the pH changes during an acid-base titration depends on the relative
strengths of the acid and base. Four characteristic graphs can be produced
for the various combinations of strong and weak acids and bases.
A]
Strong acid-strong base.
50.0 mL of 0.100 mol L-1 HCl with 0.100 mol L-1 NaOH
B]
Weak acid-strong base
50.0 mL 0.100 mol L-1 ethanoic acid with 0.100 mol L-1 NaOH
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C]
Strong acid-weak base
50.0 mL 0f 0.100 mol –1 ammonia with 0.100 mol L-1 HCl
D]
Weak acid-weak base
50.0 mL of 0.100 mol L-1 ethanoic acid with 0.100 mol L-1 ammonia
solution.
Equivalence point:

The equivalence point is the point at which the amount of acid
and base are equal.
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
Close to the equivalence point a small addition from the burette
will cause a large change in pH.

You can read the equivalence point of the graph by finding the
centre of the vertical (steepest slope) section.

For strong acid-strong base titration the equivalence point is at a
pH of 7, indicating that [H+] = [OH-].

For weak acid-strong base titration the equivalence point is
about a pH of 9, indicating that [H+] < [OH-]. The excess OH- is
produced by hydrolysis of the conjugate base of the weak acid
i.e. CH3COO-.

The reverse is true for strong acid-weak base titration. The equivalence
point is acidic due to the hydrolysis of the conjugate acid of the weak
base (NH4+).

In the case of weak acid-weak base the equivalence point is approx.
pH 7 but it is not clearly defined. This titration is avoided whenever
possible.
Buffer zone:

Buffers are formed when we mix weak acids (or bases) with their
conjugate bases (or acids).

Whenever a weak acid (or base) is involved in a titration we are going
to make a buffer.

The buffer zone on the pH graphs is centred halfway to the equivalence
point, i.e. the point at which the concentration of the acid (or base) and
its salt are equal.

At that point the pH of the solution is equal to the pKa of the acid.
Indicators:

Indicators are weak acids whose conjugate bases are a different colour
from their acids.

The colour change occurs when [acid] = [base]. When the pH of the
solution < pKa of the indicator, the indicator is one colour. When the
pH > than pKa of the indicator, the indicator is a second colour.

The colour change generally occurs over 2 pH units.

When the indicator changes colour pH = pKa of the indicator.

The correct indicator for a titration can be chosen by getting one whose
pKa is in the region of the steepest part of the titration curve.
Indicator
Methyl red
Methyl orange
Phenolphthalein
Bromothymol blue
Cresyl blue
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pKa value
5
4
9
7
11
pH range
4–6
3–5
8 – 10
6–8
10 -12
Example:
Use the graph below to answer the following questions.
The titration is for 25 mL 0.1 mol L-1 ammonia with dilute HCl
Titration HCl versus ammonia
14
13
12
11
10
9
pH
8
7
6
5
4
3
2
1
0
0
5
10
15
20
25
30
35
40
45
50
Volume HCl
a.
b.
c.
d.
e.
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What is the pH at the beginning of the titration?
What volume of acid has been added at the equivalence point?
Label the buffer zone.
At what point in the titration is the pH of the solution equal to the pKa?
Reading from the curve, what is the magnitude of pka(NH4+)?