UNZA Pharmacy Training Lecture notes

2. 2
Ionization of Solutes:
Electrolytes
Are solutes which dissociate into ions if the dielectric constant of the
solvent is high enough to cause sufficient separation of the attractive
forces between the oppositely charged ions.
Ionization (dissociation) of electrolytes has several consequences e.g.
Hydrogen ion concentration and pH:
The dissociation of water can be represented by:
H2O H+ + OH-
In pure water the concentrations of H+ and OH- ions are equal and at
25°C both have the values of 1 x 10-7 mol/l.
Acid
a substance which donates a proton (or hydrogen ion)
the addition of an acid to water will increase hydrogen ion concentration
(more than 10-7 mol/l)
Base
a substance that accepts protons
the addition of a base will decrease the concentration of hydrogen ions.
[H3O+]

3. 3
The hydrogen ion concentration range is from 1 mol/l for
a strong acid down to 1 x 10-14 mol/l for a strong base.
To avoid the use of such low values  pH has been introduced
as a more convenient measure of hydrogen ion concentration.
pH is defined as the negative logarithm of the hydrogen ion
concentration [H+]
pH = -log10 [H+]
 pH of a neutral solution like pure water is 7, why?
because the conc. of H +ions (and OH -) ions = 1 x 10-7 mol/l
 pHs of acidic solutions will be less than 7
pHs of alkaline solutions will be greater than 7
pH has several important applications in pharmaceutical practice.
- Affect the solubilities of drugs that are weak acids or bases
- Affect the stabilities of many drugs
- Affect the ease of absorption of drugs from the GIT

4. 4
Dissociation (or ionization) constants and pKa:
In solutions of weak acids or weak bases equilibria exist between
undissociated molecules and their ions.
For a weakly acidic drug HA: The ionization constant
(dissociation constant) Ka of a weak acid can be obtained by
applying the Law of Mass Action: HA H+ + A-
[H+] [A-] [HA]
Ka= pKa = pH + log
[HA] [A-]
pKa = the negative logarithm of Ka
Henderson-Hasselbalch equation:
A general equation that is applicable to any acidic drug with one
ionizable gp : Cu = conc. of the unionized Ci = conc. of the ionized species
Cu
pKa = pH + log
Ci

5. 5
The ionization constant (dissociation constant) Ka of a protonated
weak base is given by B + H+ BH+
[H+] [B]
Ka=
[BH+]
Taking the negative log of this equation:
[BH+]
pKa = pH + log
[B]
Henderson-Hasselbalch equation:
A general equation that is applicable to any weak basic drug with
one ionizable group where:
Ci = conc. of protonated ; Cu = conc. of the unionized species
Ci
pKa = pH + log
Cu

6. 6
Buffer Capacity
A buffer counteracts the change in pH of a solution upon the
addition of a strong acid, a strong base, or other agents that tend
to alter the hydrogen ion concentration.
Buffer capacity β: buffer efficiency, buffer index or buffer value
Is the resistance of a buffer to pH changes
upon the addition of a strong acid or base.
Definition
The ratio of amount added of strong base (or acid) to small
change in pH brought about by this addition.
β = Δ B
Δ pH
ΔB = the small addition in gram equiv./liter of strong base
added to the buffer solution to produce a pH change
Δ pH = pH change

7. 7
The buffer capacity of the solution has a value of 1:
of strong base (or acid) togram equiv.1when the addition of
pH unit.1of the buffer solution results in a change ofliter1
Acetate bufferExample:
acetic acid & sodium acetate
0.1 mole each in 1 liter of solution.
a) 0.01 mole portions of NaOH is added
HAc + NaOH NaAc + H2O
(0.1 – 0.01) (0.01) (0.1 + 0.01)
b) The conc. of Na acetate (the [salt] in buffer equation)  by 0.01 mol/l
& the conc. of acetic acid [acid]  by 0.01 mol/l
because each addition of base converts 0.01 mole of acetic acid into
0.01 mole of sodium acetate according to the reaction.

8. 8
Before the addition of the first portion of NaOH,
the pH of the buffer solution is:
pKa = pH + log pH = pKa - log
pH = pKa + log
pH = 4.76 + log (0.1) = 4.76
(0.1)
pH = pKa
The changes in concentration of the salt and the acid by the
addition of a base are represented by
pH = pKa + log pH = 4.76 + log[salt ] + [base]
[acid] - [base]
Cu [acid]
Ci [salt]
[acid]
[salt]
[salt ]
[acid]
(0.1) + 0.01
(0.1) – 0.01

9. 9
pH of
solution
Moles of
NaOH added
4.760
4.850.01
4.940.02
5.030.03
5.130.04
5.240.05
5.360.06
Buffer
Capacity, β
0.11
0.11
0.11
0.10
0.09
0.08
0.07

10. 10
The buffer capacity is not a fixed value for a given buffer
system, but depends on the amount of base added.
With the addition of more NaOH, the buffer capacity
decreases rapidly, and, when sufficient base has been
added the acid convert completely into sodium ions and
acetate ions
The buffer has it’s greatest capacity before any base is
added where [salt] / [acid] = 1, and according to equation,
pH = pKa.
The buffer capacity is influenced by an increase in the total
conc. of the buffer constituents since a greater conc. of salt
and acid provides a greater alkaline and acid reserve.

11. 11
formore exact equationVan Slyke developed a
calculation of buffer capacity β
]+O3Ka [HC3.2=β
[H3O+])2+Ka)
C = The total buffer concentration (the sum of the molar
concentrations of the acid and the salt).
Ka = dissociation constant
H3O+ = hydrogen ion concentration
The equation permits the calculation of the buffer capacity
at any hydrogen ion concentration, i.e. when no acid or base
has been added to the buffer
[H+]

12. 12
Example:
If hydrogen ion concentration is 1.75 x 10-5, pH = 4.76
what is the capacity of the buffer containing 0.10 mole of
each of acetic acid and sodium acetate per liter of solution ?
The total concentration , C = [acid] + [salt], is 0.20 mol/l and
the dissociation constant Ka is 1.75 x 10-5
]+O3Ka [HC3.2=β
(Ka + [H3O+])2
115.0=)5-10X75.1) x (5-10x75.1x (20.0x3.2β =
[(1.75x10-5) +(1.75 X 10-5)]2

13. 13
Maximum buffer capacity .
The maximum buffer capacity occurs when pH = pKa or
when (H3O+) = Ka
C576.0=2)+O3H(C303.2max =β
2)+O3H2(
β max = 0.576 C
Where C is the total buffer concentration
Example:
What is the maximum buffer capacity of an acetate buffer
with a total concentration of 0.20 mol/l?
β max = 0.576 C
= 0.01152 = 0.01

14. 14
Pharmaceutical Significance
Buffer in biological & pharmaceutical systems
In vivo biological buffer systemsI.
Blooda)
 Blood is maintained at a pH of about 7.4 by:
the 1° buffers in the plasma &
the 2° buffers in the erythrocytes.
The buffer capacity of blood = 0.039 gram equiv. per liter/pH
unit for whole blood of which: 0.031 by the cells
0.008 by the plasma
* When the pH of the blood goes below 7.0 or
above 7.8, life is in serious danger.
* The pH of the blood in diabetic coma is dropped to about
6.8

15. 15
Lacrimal fluidb)
Tears have a great degree of buffer capacity, allowing
a dilution of 1:15 with neutral distilled water before an
alteration of pH is noticed.
The pH of tears is about 7.4 with a range of 7 to 8

16. 16
Pharmaceutical BuffersII.
Buffer solutions are used in pharmaceutical formulation
particularly in ophthalmic preparations
Gifford suggested two stock solutions of:
- boric acid and monohydrated sodium carbonate
- mixed in various proportions to yield buffer solutions
of pH values from about 5 - 9.
Sorensen proposed a mixture of the salts of:
- sodium phosphate for buffer solutions of pH 6 to 8.
The Clark-Lubs mixtures and their pH ranges
a. pH 1.2 to 2.2: HCI and KCI
b. pH 2.2 to 4.0: HCI and potassium hydrogen phthalate
c. pH 4.0 to 6.2: NaOH and potassium hydrogen phthalate
d. pH 5.8 to 8.0: NaOH and KH2PO4
e. pH 7.8 to 10 : H3BO3, NaOH and KCl
Sodium chloride is added to buffer mixture to make it isotonic with body.

17. 17
of Pharmaceutical buffer solutionsPreparation
Factors of some importance in the choice of pharmaceutical
buffer include:
 Availability and cost of chemicals
 Sterility of the final solution.
 Stability of the drug and buffer on aging.
 Freedom from toxicity.
For example, a borate buffer, because of its toxic effects,
cannot be used for a solution to be administrated orally or
parenterally.

18. 18
The following steps should be used in preparing buffer systems
a. Select a weak acid having a pKa approximately equal to
the pH wanted to insure maximum buffer capacity.
b. From the buffer equation, calculate the ratio of salt and weak
acid required to obtain the desired pH.
log Cu = pKa - pH
Ci
c. Consider the individual concentrations of the buffer salt
and acid needed to obtain a suitable buffer capacity.
β = 2.3 C Ka [H3O+]
)Ka + [H3O+])2
A concentration of 0.05 to 0.5 molar is sufficient and
a buffer capacity of 0.01 to 0.1 is sufficient.
d. Finally, determine the pH and buffer capacity of the completed
buffered solution using a pH meter.

19. 19
III. Influence of Buffer Capacity and pH
on Tissue Irritation
Solutions to be applied to tissues or administered
parenterally are liable to cause irritation, if their pH is
greatly away from the normal pH of the body fluid.
 must be considered when formulating:
- ophthalmic solutions
- parenteral products
- fluids to be applied to abraded surfaces.
Factors affecting: (i &ii are of greater significance)
i) The buffer capacity of the solution
ii) The volume to be used in relation to that of body fluid
with which the buffered solution will come in contact
iii) Actual pH of the solution
iv) The buffer capacity of the body fluid

20. 20
Tissue irritation due to large pH differences between:
the solution administered & the physiological fluid
is minimized:
(a) The lower the buffer capacity of the solution
(b) The smaller the volume used for a given concentration.
(c) The larger the volume and buffer capacity of the
physiological fluid
 The pH of solutions for introduction into the eye may vary
from 4.5 to 11.5 without marked pain or damage.
This is true only if the buffer capacity was kept low.
 Sorensen phosphate buffer produced irritation in the eyes of
a number of subjects when used outside the pH range of 6.5
to 8
 Boric acid solution of pH 5 produced no discomfort in the
eyes of the same subjects. Why?
Because of the very low buffer capacity of boric acid
compared to that of the phosphate buffer.

21. 21
Parenteral solutions for injection into the blood stream are
usually not buffered or they are buffered to a very low
capacity so that the buffers of the blood may bring them
within the physiological pH range.

22. 22
IV. Influence of Buffer Capacity and pH
on Optimum Therapeutic Response
• The undissociated form of a weakly acidic or basic drug has
a higher therapeutic activity than the dissociated salt form
WHY?
• Because: the undissociated form is lipid soluble and can
penetrate body membranes, whereas the ionic form is not
lipid-soluble and can only penetrate membranes with great
difficulty.
• Thus, the therapeutic response of weakly basic alkaloids
(used as ophthalmic drugs) increases as the pH of the
solution increases, and hence-concentration of the
undissociated base, was increased.

23. 23
Mandelic acid, benzoic acid and salicylic acid have
pronounced antibacterial activity in non ionized form but
have no activity in ionized form. Accordingly, these
substances require an acidic pH to function as
antibacterial agents. Thus sodium benzoate is effective
as a preservative:
4% concentration at pH 7
0.06 to 0.1% concentration at pH 3.5 to 4
0.02 to 0.03% concentration at pH 2.3 to 2.4

24. 24
V. Influence of Buffer Capacity and pH on
Drug Stability
Buffer is used to prevent changes in pH due to the
alkalinity of the glass or acidity of CO2 from dissolved
air
Solutions as Thiamine hydrochloride may be
sterilized by autoclaving without decomposition if the
pH is below 5 above this pH thiamine hydrochloride is
unstable.
The stability of emulsions is pH dependent.

25. 25
VI. PH and solubility.
The influence of buffering on the solubility of the alkaloidal
base:
 At a low pH a base is predominantly in the ionic form which
is usually very soluble in aqueous media
as the pH is raised more undissociated base is formed.
 Therefore, solution should be buffered at sufficiently low
pH.
 Yet, when the solution is instilled in the eye, the tears
participate in the gradual neutralization of the solution and
the conversion of the drug from the physiologically inactive
form to the undissociated base the base can then readily
penetrate the lipoidal membrane.
 As the base is absorbed at the pH of the eye, more of the
salt is converted into base to preserve the equilibrium,
hence the alkaloidal drug is gradually absorbed.

26. 26
Any Questions

27. 27
Study Questions
Define the following terms:
[Ionization, buffer capacity, in-vivo, etc]
Respond to the following questions:
 Considering a practical process, illustrate the procedural significance of buffer systems in moderation of the
reactions of a solution system
 What steps should be adopted to prepare a buffer system
Group work discussional questions:
 Discuss the variations in a solution that may constitute the buffering effects of such
pharmaceutical solutions
 What are the main key points to consider when a buffer system is being pharmaceutically
processed