Buffers are substances or a combination of substances that, by their presence in solution, resist changes in pH upon the addition of small quantities of acid or alkali. Buffered solutions are necessary in many experiments conducted in pharmaceutical research. Drug stability, partitioning, diffusion and dissolution studies are few of the applications where buffers are used to mimic biological fluids. A buffer acts by neutralizing hydrogen ions or hydroxyl ions added to it. Buffers can function as such because they are either weak acids or bases and have their roots in their respective ionic equilibria. The term buffer implies protection or shielding. Buffer protects the formulation from a sudden change in pH. The resistance to this change is known as buffer action.
For example, a mixture of acetic acid and sodium acetate is an acidic buffer and a mixture of ammonium hydroxide and ammonium chloride is an example of basic buffer.
Auto ionization of Water-
Water contains low concentrations of ions, which is a result of the transfer of a proton from one water molecule to another.
H2O + H2O ↔ H3O+ + OH– eq (1)
In Eq. (1), H3O+ is the hydronium ion and OH− is the hydroxyl ion. The equilibrium constant for this reaction can be written as
Kc = [H3O+] [OH–] / [H2O]2 eq (2)
In aqueous solutions, the concentration of water is effectively a constant (55.55 M), and thus Eq. (2) simplifies to
Kw = [H3O+] [OH–] eq (3)
where, Kw is the autoionization constant of water, also known as the ionic product of water. The value of Kw is very small, being equal to 1.007 × 10–14 at 25°C.
For convenience, Sørensen proposed the p scale, where numbers such as Kw would be expressed as the negative of their base 10 logarithms. The value of pKw can then be calculated as:
pKw = -logKw eq (4)
and has a value equal to 13.997 at 25°C.
pH is defined as
pH = -log [H3O+] eq (5)
pOH = -log [OH–] eq (6)
Hence, Eq. (3) can then be expressed as
pKw = pH + pOH eq (7)
Buffer Equation for Weak Acid and Its Salt
The pH of a buffer solution and the change in pH upon the addition of an acid or base may be calculated using the buffer equation. This expression is developed by considering the effect of a salt on the ionization of a weak acid when the salt and acid have an ion in common.
According to the Brønsted–Lowry model, an acid is a substance capable of donating a proton to another substance, such as water:
HA + H2O ↔ H3O+ + A–
The dissociation constant for the weak acid, as per the above equation, can be written as
Ka = [H3O+] [A–] / [HA]
Henderson–Hasselbalch equation, for a weak acid and its salt is
pH = pKa + log [salt] / [acid]
where pH = – log [H3O+] [from Eq. (5)], pKa = – log Ka, known as the dissociation exponent.
Buffer Equation for Weak Base and Its Salt
A base is a substance capable of accepting a proton donated by another substance, such as water:
B + H2O ↔ BH+ + OH–
In case of a weak base, the ionization constant can be written as:
Kc = [BH+] [OH–] / [B] [H2O]
Kb = [BH+] [OH–] / [B]
This is the Henderson–Hasselbalch equation for a weak base and its salt
pH = pKw – pKb – log [base] / [salt]
BUFFER CAPACITY: Buffers are able to protect preparations from large swings in pH. However, every buffer will reach a point where it no longer can protect the preparation from pH changes. The magnitude of resistance of a buffer to pH changes is referred to as its buffer capacity, E. It is also called buffer efficiency, buffer index and buffer value, and it designates the effectiveness of a buffer in minimizing pH change. It is defined as the ratio of the increment of strong base (or acid) (in gram equivalents per litre) to the small change in pH brought about by this addition. The formula for calculating an average buffer capacity is as follows:
ß = ∆B / ∆pH
where ∆ depicts a finite change and ∆B denotes a small increment in gram equivalents per litre of strong base (or acid) added to the buffer solution to produce a pH change of ∆pH. According to this equation, the buffer capacity of a solution has a value of 1 when an addition of 1 g Eq of strong base (or acid) to 1 L of the buffer solution results in a change in pH by unity.
Koppel, Spiro and Van Slyke developed a more exact equation:
ß = 2.303C Ka [H3O+] / (Ka + [H3O+])2
where C is the total buffer concentration (i.e. the sum of the molar concentrations of the acid and the salt).
The buffer capacity is affected not only by the [salt]/[acid] ratio but also by the total concentrations of acid and salt. An increase in the concentration of the buffer components results in a greater buffer capacity or efficiency.
Maximum buffer capacity: The maximum buffer capacity is achieved when pH = pKa, or in equivalent terms, where [H3O+] = Ka.
Multiple choice questions (MCQs)
1.A buffer solution comprises which of the following
a)A weak acid in solution
b)A strong acid in solution
c)A weak base in solution
d)A weak acid and its conjugate base in solution
2. The concept of pH was introduced by
3. Which of the following type of electrode used in pH meter?
d)All of the above
4.The term ‘Buffer capacity’ first introduced by
5.Buffer contains acetic acid-sodium acetate having pH range ____ used in the pharmaceutical preparation
a)1.2 to 6.6
b)5.0 to 8.0
c)3.8 to 5.6
d)7.8 to 10.6
6. pH of tears has the value of
7.As the pKa of an acid increases, the acid will be
c)Converted to neutral solution
d)Converted to basic solution
8.Buffers are mixtures of
a)Strong acid and strong base
b)Strong acid and weak base
c)Weak acid and their conjugate base
d)Weak base and their conjugate acid
9.If a solution has to be buffer, its pH should be
a)At its pKa value
b)At its Ka value
10.Normal pH of blood is
11.Buffer Capacity is the maximum at
a)pKa = pH
b) pKa < pH
c) pka = Concentration
d) pka > pH
12.The pKw at 25°C is
13.Isotonic solutions have the same
a) Vapour pressure
b) Osmotic pressure
c) Atmospheric pressure
d) Internal pressure
14.pH of 0.1M NaOH is
15.pH can be kept constant with the help of
d)super saturated solution
- d) A weak acid and its conjugate base in solution
- c) Sorensen
- d) A weak acid and its conjugate base in solution
- a) Van Slyke
- c) 3.8 to 5.6
- d) 7.4
- a) more weaker
- c) Weak acid and their conjugate base
- a) At its pKa value
- d) 7.4
- a) pKa = pH
- a) 14.0
- b) Osmotic pressure
- c) 13
- c) buffer solution
1. GAURAV KUMAR JAIN – THEORY & PRACTICE OF PHYSICAL PHARMACY, 1st edition 2012 Elsevier, page no. 142-147.
2. Martins Physical Pharmacy, 6th edition 2011, page no. 302-312.