This document discusses displacement titration, specifically titrations of anions of weak acids (Bronsted bases) with strong acids. It provides examples of titrating salts of weak acids like sodium acetate, potassium cyanide, and borax with hydrochloric acid. The pH at the endpoint is calculated based on the concentration of the weak acid produced. Indicators that can be used for different titrations are also discussed based on the pH range at the endpoint.

1. Displacement Titration
Pabitra Kumar Mani , Assoc. Prof., Ph.D. ACSS, BCKV

2. Volumetric analysis (Titrimetry)
(i) Acid-Base and (ii) Displacement titrimetry
Titrations of anions of weak acids (Bronsted Bases) with
strong acids (Displacement Titrations)
On the Bronsted theory the so –called titration of
solutions of hydrolysed salts, is merely the titration of a weak
Bronsted base(carbonate ion , borate ion, acetate ion) with a
strong acid.
B+A- + HX
⇋ HA
+ B+XSalt
strong Acid
weak Acid
Salt
B+A- + DOH ⇋
BOH
+ D + AStrong base
weak base
CH3COONa + HCl = CH3.COOH + NaCl
The weak acetic acid was apparently displaced by the strong
HCl acid, and the process was refrred to as displacement
titration

3. Salt of weak acid vs agt. a strong acid
100 ml of 0.2 (N) KCN ----- 0.2(N) HCl
Initial pH of 0.2 N KCN
KCN + H2O = KOH + HCN
for HCN, Ka= 7.2 X 10-10, pka=9.2
pH= ½ pKw + ½ pKa + ½log c
= 7 + 4.6 + ½log 0.2 =11.45
The pH at e.p is due to 0.1 (N) HCN produced (vol.is diluted so, (N)
reduces)
KCN+ HCl = HCN + KCl
pH = ½ pKa - ½log c = 4.6 + 0.5 = 5.1
When 0.1 ml acid (HCl) is in excess, pH= 4.0
Thus the inflexion region in the titration curve covers
the pH range 5-3.7 and the indicator like M.O, BCG,
M.R can be used with success for detecting the e.p.

4. Borax (Na2B4O7) can similarly be titrated with HCl
Na2B4O7 +2HCl +5H2O = 4 H3BO3+ 2NaCl
(0.2N)
0.2 N
weak monobasic acid
Ka= 6x10-10
The pH at e.p. (pH=5.1) is due to 0.1 M boric acid, i.e., 5.1.
Further addition of HCl will cause a sharp decrease of pH
and any indicator covering the pH range 3.7-5.1 may be
used, Bromo cresol green, methyl orange and methyl
red

5. Titration of a carbonate ion with strong acid
In case of Na2CO3 can be tritrated in 2 stages according to the equn.
Na2CO3 + HCl = NaHCO3 + NaCl........1
NaHCO3 + HCl = H2CO3 + NaCl...........2
At 1st e.p. , pH = ½pK1 + ½pK2 = 8.3
Here indicator PhTh becomes
colourless and can be used to detect
the e.p.
The pH at the 2nd e.p. is due to that of
H2CO3 produced in soln.
pH = ½pK1 - ½log c = 3.8
c= 0.1 M
So, indicator that can be used is M.O;
congo red and BPB

7. Salt of Weak base-Strong Base
NH4Cl (0.2 N) - NaOH (0.2 N)
NH4Cl + NaOH = NH4OH + NaCl
The pH at e.p will be the alkaline side, and it can be calculated
in the same manner as the pH of a free weak base like NH 4OH
At pH due to NH4OH (0.1N)
pH = pKw - ½pKb + ½log c = 14-2.4 + 3.8+ ½log (0.1)= 11.1

8. Criteria for using a pH indicator
There must be at least 2 units of pH change near the
stoichiometric end point for the solution of 0.1 ml of the titrant
pH on the either side of the e.p due to ± 0.1 ml addition of
the titrant can be obtained from relevant equations and the
difference will indicate whether the change is large enough
to permit a sharp end point to be determined. If the pH
change is satisfactory an indicator should be selected which
changes colour at or near the e.p.
Generally if pK In-a ±1 value for an indicator forms within the
range of pH change at the e.p. the indicator is suitable for the
determination of the e.p

9. Other types of pH indicators
(i) Mixed indicators:
Mixture of 2 or more simple indicators or mixture of an indicator
with a suitable dye are often employed for better detection of end point. in
pH titration.
The purpose of using mixed indicator is to render the colour changes
more contrasting or to obatined a sharp colour change in a narrow
range of pH. Thus methylene Blue (dye) (red-ox indicator) modifies the
yellow to red colour change of Methyl orange
(M.O.+ Methylene blue) : green- gray-violet
M.O :
Yellow to Red
Equal parts of neutral red indicator and methylene blue (both 0.1% concn)
makes the colour changes more contrasting from green to violet and both
these mixed indicators are suitable for titration of very weak base like
pyridine agt. a strong acid.
The main indicator mixed with a pre or oxillary indicator or two indicators
with overlapping pH ranges may help the colour change to take place
over a narrower pH change. Thus 1 drop of Sodium Cresol Red with 3
drops of Na-Thymol Blue (both 0.1%) enables a pH of 8.3 to be exactly
detected 8.2(pink)—8.3 –8.4 (violet)

10. (ii) Universal pH indicators :
Multiple range pH indicators are prepared by mixing a no. of
indicators in proper ratio so that final indicator can show
different colours over a very long range of pH.
Thus a mixture of 0.05 g M.O., 0.15 g M.R, 0.30 g BTB,
0.35 g PhTh in one Litre of 60% ethanol
pH
Colour
3
4
5
6
Red
Orange
Red
orange
Y
7
8
Y-G Greenish
blue
9
10
11
B
V
R-V

11. Precipitation titrations
The most important precipitation processes in titrimetric analysis utilise
silver nitrate as the reagent (argentimetric processes)

12. Detection of the end point: INDICATORS
Indicator forming a coloured compound with a titrant. The Mohr
method for determining chloride serves as an example. The
chloride is titrated with std silver nitrate soln. A soluble
chromate salt is added as the indicator. This produces a yellow
soln. When the pptn of the chloride is complete, the first excess
of Ag+ reacts with the indicator to precipitate red silver
chromate:

13. In actual practice, The indicator concentration is kept at 0.002
to 0.005 M
2CrO42- + 2H+  2HCrO4 Cr2O72- +H2O
Titration should be done in
alkaline pH,6-9, in acid soln
the rn occurs as follows

15. The action of these indicators is due to the fact at the e.p the
indicator is adsorbed by the ppt, and during the process of
adsorption a change occurs in the indicator which leads
to a substances of difft. colour, they have therefore been
termed as adsorption indicator
The indicator which is a dye, exists in soln as the ionized form, usually an
anion, In-.
Consider the titration of Cl- with Ag+. Before the equivalence point, Cl- is in
excess, and the primary adsorbed layer is Cl-. This repulses the indicator
anion, and the more loosely held secondary (counter ) layer of adsorbed
ions is cations such as Na+:
Beyond the equivalence point, Ag+ is in excess, and the surface of the
precipitate becomes positively charged, with the primary layer being Ag +.
This will now attract the indicator anion and adsorb it in the counterlayer:

16. Thus the colour is formed of Ag+In- on the ppt surface which
is more intense than In- colour in solution. It is interesting to
note that Ag+In- comp. as such is not as intensity coloured as
the adsorbed colour on the ppt.
Ag-eosinate is soluble in water, and has reddish orange
colour but adsorbed Ag-eosinate on AgX is intense reddish
violet colour.
Fajans (1923) first introduced
Fluorescein ----( for Cl-)
Eosin------------ (Br- substituted)
Erythrosin.............(I- substituted)
The adsorption indicator should not be strongly adsorbed
on the ppt and it must be reversible

23. One important group of colour indicators is derived from 1:10
phenantholine (ortho-phenanthroline) which forms a 3:1
complex with iron(II). The complex known as 'ferroin‘ undergoes
a reversible redox reaction accompanied by a distinct colour
change

26. Equivalence point is determined by using an indicator (W.Ostald)
(pH-indicator)
HIn ⇋
H+ +
Acidic
Indicator(mol.form)
InOH
In-
HIn and In- are differently coloured
ionic form
⇋ In+ + OH-
InOH and In+ are differently coloured
Basic
ionic form
Indicator(mol.form)
⇋
Tautomeric transformation
HIn
Hin*
⇋ H+ + In-
HIn and In- are differently coloured
HIn ⇋ H+ + In[ H ][ In ]
kIn-a =
+
[ HIn]
[H ]
+
= kIn-a
pH = - logkIn-a
−
[ HIn]
[ In ] [ HIn]
−
- log [ In ]
−
pH = pKIn-a + log [ In ]
−
[ HIn]
Alkaline colour intensity
acidic colour intensity

27. [ In ]
When [ HIn] = 10, pH= pKIn-a + 1
[ In ]
and When [ HIn] = 1/10, pH= pK In-a - 1
−
−
Operationally,
pH= pKIn-a ± 1
[ In ] 10 times concn of HIn i.e., when the color due to [ In ] is dominant
−
−
During titration, if pH at equivalence point, lies in the range pK In-a ± 1,
that indicator can be used to detect the e.p. In the particular titration
InOH
⇋ In+ + OH-
[ In ][ OH ]
=
[ InOH ]
+
KIn-b
−
[ InOH ]
∴ [OH ] = kIn-b [ In ]
-
[ In ]
+
-log[OH ] or, pOH = pKIn-b + log10
-
[ InOH ]
[ In ]
pH = 14 - pKIn-b -log10 [ InOH ]
+
[ In ]
+
pH = pKIn-a -log10 [ InOH ]
+
pH + pOH =pKw=14
pH -14 = - pOH
HA ⇋ H+ + ApKa+pKb= pKw