8. THE LAW OF MASS ACTION
The rate of a
chemical reaction
is proportional to
the active masses
of the reacting
substances.
9. APPLICATIONS OF THE LAW OF MASS ACTION TO SOLUTIONS
OF WEAK ELECTROLYTES
Electrolytes can be strong or weak
Their dissociation
Strong electrolytes----- no equilibrium
Weak electrolytes------ equilibrium between
undissociated molecules and ions
Reference: Beckett and Stenlake, Practical Pharmaceutical Chemistry,
Part-1, 4th Edition, page no. 84
10. DISSOCIATION OF WATER
Water extremely weak electrolyte and slightly dissociated.
Reference: Beckett and Stenlake, Practical Pharmaceutical Chemistry, Part-1,
4th Edition, page no. 84
12. STRENGTHS OF ACIDS AND BASES
Definitions of acids and bases.
A H+ + A-
Strength of acid: Concentration of H+ ions which it yields
upon ionization and will depend upon the degree of
dissociation, α, at any given concentration.
The acid dissociation constant, ka, gives a relationship
between α and the concentration.
It is measure of strength of acid
Similarly for base.
Reference: Beckett and Stenlake, Practical Pharmaceutical Chemistry, Part-1,
4th Edition, page no. 84
13. The relationship between dissociation constant of an acid
and its conjugate base.
From equation, Stronger the acid weaker the conjugate base and vice-versa.
Reference: Beckett and Stenlake, Practical Pharmaceutical Chemistry,
Part-1, 4th Edition, page no. 84
14. HYDROLYSIS OF SALTS
When salts are dissolved in water, interaction may occur with
the ions of water and the resultant solution may become acid
or alkaline according to the nature of salt.
Such interaction may be termed as hydrolysis
1. Strong acid- strong base
Ex. NaCl- Solution neutral
1. Weak base- strong acid salts
Ex. NH4Cl = NH4
+ combined with OH-, therefore acidic.
Reference: Beckett and Stenlake, Practical Pharmaceutical Chemistry,
Part-1, 4th Edition, page no. 84
15. 3. Weak acid- strong base salts
Ex. Potassium cyanide--- CN- reacts with H+, solution
remains alkaline
4. Weak acid- weak base salts
Ex. Ammonium acetate
Reference: Beckett and Stenlake, Practical Pharmaceutical Chemistry,
Part-1, 4th Edition, page no. 84
16. END POINT DETECTION
The object of carrying out an acid base titration is to
determine the equivalence point/ quantity of the other
substance required for neutralization.
This point is called end-point or the equivalence point.
If both acid and the base are strong electrolytes the resulting
solution will be neutral having pH =7.
But if either of the two is a weak electrolyte the resulting salt
formed will hydrolyze to some extent and the solution will
possess some acidic or alkaline properties.
17. END POINT DETECTION
Indicator selection
Colour change detection
Quantity of indicator (0.0001% - 0.0004%)
Presence of colloidal substances
Effect of temperature
18. NEUTRALIZATION INDICATORS
Acid - Base indicators (also known as pH indicators) are substances which change
colour with pH. They are usually weak acids or bases, which when dissolved in
water dissociate slightly and form ions.
Consider an indicator which is a weak acid, with the formula HIn.
At equilibrium, the following equilibrium equation is established with its conjugate
base:
The acid and its conjugate base have different colors.
19. EXAMPLE
Colorless (Acid) Pink (Base)
Phenolphthalein
Phenolphthalein is a colorless, weak acid which dissociates in water forming
pink anions.
Under acidic conditions, the equilibrium is to the left, and the concentration
of the anions is too low for the pink color to be observed. However, under
alkaline conditions, the equilibrium is to the right, and the concentration of the
anion becomes sufficient for the pink color to be observed.
21. RANGE OF INDICATORS
pH indicators generally change color over a range of two pH units.
At a low pH, a weak acid indicator is almost entirely in the HIn form, the color of
which predominates.
As the pH increases- the intensity of the color of HIn decreases and the equilibrium is
pushed to the right.
Therefore, the intensity of the color of In- increases.
An indicator is most effective if the color change is distinct and over a low pH range.
For most indicators the range is within ±1 of the pKln value.
Indicator
Color
pKln pH range
Acid Base
Thymol Blue - 1st change Red Yellow 1.5 1.2 - 2.8
Methyl Orange Red Yellow 3.7 3.2 - 4.4
Bromocresol Green Yellow Blue 4.7 3.8 - 5.4
Methyl Red Yellow Red 5.1 4.8 - 6.0
Bromothymol Blue Yellow Blue 7.0 6.0 - 7.6
Phenol Red Yellow Red 7.9 6.8 - 8.4
Thymol Blue - 2nd change Yellow Blue 8.9 8.0 - 9.6
Phenolphthalein Colourless Pink 9.4 8.2 - 10.0
22. Buffer solution
It is solution of substance or a mixture of substances which
helps in maintaining and establishing specific pH.
Buffer capacity
Buffer capacity is a measure of the efficiency of a buffer in
resisting changes in pH. Conventionally, the buffer capacity
is expressed as the amount of strong acid or base, in gram-
equivalents, that must be added to 1 liter of the solution to
change its pH by one unit.
23. THEORY OF INDICATORS- OSTWALD THEORY
First theory
According to this theory the un-dissociated indicator acid (HIn)
or a base (InOH) has a colour different than its ion.
For an acid indicator equivalence can be written as-
HIn H+ + In-
24.
25.
26.
27. RESONANCE THEORY / CHROMOPHORE THEORY
The difference in colour of same compound in acid and
base medium is apparently due to difference in
structure if two forms.
Colour change by the compound is associated with the
capability of the compound to absorb visible light and
this capability is related to the electronic structure.
Change in the electronic structure will result in
absorption of different colour components of light with
resultant colour change.
28. THE RED COLOR IN ALKALINE SOLUTION IS DUE TO THE QUINONOID
STRUCTURE WITH RESULTING INCREASED POSSIBILITIES FOR RESONANCE
BETWEEN VARIOUS IONIC FORMS.
29. MIXED INDICATORS
In some cases, the pH range is very narrow and the
colour change over this range must be very sharp.
This is not easily possible with ordinary acid-base
indicators.
The result may be achieved by the suitable mixture of
indicators.
Ex. A mixture of equal parts of neutral red (0.1%
solution in alcohol) and ethylene blue (0.1% solution in
alcohol) gives a sharp color change from violet-blue to
green in passing from acid to alkaline solution at pH 7.
30. UNIVERSAL OR MULTIPLE RANGE INDICATORS
By suitably mixing certain indicators the color change may be
made to extend over a considerable portion of the pH range.
Such mixtures are called Universal indicators.
They are not suitable for quantitative titrations but may be
employed for the determination of the approximate pH of
the solution by colorimetric method.
Ex. Dissolve 0.1 g phenolphthalein, 0.2 g methyl red, 0.3 g
methyl yellow, 0.4 g bromothymol blue and 0.5 g of thymol
blue in 500 ml absolute alcohol and sufficient NaOH solution
until color is yellow.
31. THE COLOR CHANGES ARE-
pH 2– Red
pH 4– orange
pH 6- Yellow
pH 8– Green
pH 10– Blue
37. Weak Acid Vs Strong Base
The neutralization of 100 mL of 0.1 M acetic acid (ethanoic acid) with 0.1 M
sodium hydroxide solution will be considered here;
40. Strong Acid Vs Weak Base
This may be illustrated by the titration of 100 mL of 0.1M aqueous
ammonia (Kb= 1.85 x with 0.1M hydrochloric acid at the ordinary
laboratory temperature.
41. Strong Acid Vs Weak Base
Fig. 10.5 Neutralisation curves of 100 mL 0.1M aqueous ammonia (Ka = 1.8 x and
of 0.1M base (Ka = 1 x with 0.1M hydrochloric acid.
42. This case is exemplified by the titration of lOOmL of 0.1M acetic acid
(Ka = 1.82 x 10-5) with 0.1 M aqueous ammonia (K, = 1.8 x
Weak Acid Vs Weak Base
43. Weak Acid Vs
Weak Base
Fig. 10.6 Titration of 50
mL of 0.1M H3PO4 with
0.1M KOH
