PHARMACEUTICAL BUFFERS: Buffers are used to establish and maintain ion activity within narrow limits. The most common buffered systems in pharmacy are used:
- To establish hydrogen-ion activity for the calibration of pH meters.
- To prepare isotonic dosage form formulations.
- To adjust pH of system in analytical procedures.
- To maintain stability of dosage forms.
- To study the pH dependence of drug substance solubility.
- To study the pH-stability profile of drug substances.
Selection of Buffer System- The selection of a buffer system for use in a pharmaceutical dosage form is relatively clear-cut. It is evident from the preceding discussion that the most important condition for a buffer is the approximate equality of the pKa value of the buffer with the proposed optimal pH value for the formulation. Knowledge of the pH-stability profile of a drug substance enables one to deduce the pH range in the desired formulation. The basis for the most appropriate buffer system would be the weak acid or base whose pKa or pKb value is numerically equal to the midpoint of the pH range of stability. Other considerations that need to be monitored include compatibility with the drug substance. Boylan has provided the following summary of the selection criteria for buffering agents:
- The buffer must have adequate buffer capacity in the desired pH range.
- It must be biologically safe for the intended use.
- It should have little or no deleterious effect on the stability of the final product.
- It should permit acceptable flavouring and colouring of the final product
BUFFER PREPARATION: The following steps are involved in the preparation of a new buffer:
- Decide for what pH the buffer is needed.
- Select a weak acid whose pKa is approximately equal to the pH at which the buffer is to be used. This ensures maximum buffer capacity.
- Apply the Henderson–Hasselbalch equation and calculate the ratio of quantities of salt and weak acid required to obtain the desired pH. The buffer equation is suitable for approximate calculations within the pH range 4–10.
- Consider the individual concentrations of the buffer salt and acid needed to obtain a suitable buffer capacity. A concentration in the range 0.05–0.5 M is usually sufficient and a buffer capacity of 0.01–0.1 is generally satisfactory.
- Consider other important factors when deciding a pharmaceutical buffer: availability of chemical sterility of the final solution, stability of the drug and buffer upon aging, cost of materials and freedom from toxicity. For example, a borate buffer, because of its potential toxic effects, cannot be used to stabilize a solution to be administered orally or parenterally.
- One should determine the pH and buffer capacity of the buffered solution thus obtained using a reliable pH meter or pH papers (for rough estimate).
STANDARD BUFFER SOLUTION- Standard solutions of a definite pH are readily available in buffer solutions prepared from appropriate reagents. For preparing these solutions, the crystalline reagents should be dried, except boric acid, at 110–120°C for 1 h and carbon dioxide-free water should be used for making solution or for dilution purposes. The prepared standard buffer solutions should be stored in chemically resistant tight containers such as Type I glass bottles and the solution should be used within 3 months. Standard buffer solutions for various pH ranges from 1.2 to 10.0 may be prepared by appropriate combinations of the solutions.
Stabilization of Drug Substances in Formulations by Buffers- The stability of many active pharmaceutical greatly depends on the degree of acidity or basicity to which they are exposed, and any change in pH can cause considerable changes in the rate of degradation reactions. For such compounds, a buffer system is included to ensure the stability of the drug substance either during the shelf life of the product or during the period associated with its administration.
BUFFERED ISOTONIC SOLUTIONS: In addition to carrying out pH adjustment, pharmaceutical solutions that are meant for application to delicate membranes of the body should be adjusted to approximately the same osmotic pressure as that of the body.
When a solution is placed in contact with a membrane which is permeable to the solvent molecules, but not to that of solute (semipermeable membrane), the movement of solvent molecules from region of lower solute concentration to higher solute concentration, the phenomenon is called as osmosis. Consider two solutions on either side of a semipermeable membrane, which have different concentration of solute. There is a tendency of movement of solvent molecules from region of lower solute concentration to higher solute concentration until equilibrium is reached. The pressure required to prevent this movement is known as osmotic pressure. Osmotic pressure is a colligative property dependent on the number of particles of solute in solution, its degree of ionization and aggregation. Body fluids (blood and lacrimal fluid) have an osmotic pressure corresponding to that of 0.9% (w/v) sodium chloride. Thus, 0.9% (w/v) solution of sodium chloride is iso-osmotic (same osmotic pressure) with physiological fluids. In medicine, the term isotonic (of equal tone) is used interchangeably with iso-osmotic. Physiological solutions that have an osmotic pressure lower than that of body fluids, or of 0.9% (w/v) sodium chloride, are termed as hypotonic and physiological solutions with a higher osmotic pressure are known as hypertonic.
Osmolality and osmolarity: Osmolality and osmolarity are expressions of concentration reflecting the osmoticity of solutions.
Osmolality is the expression of osmolal concentration. A solution has an osmolal concentration of one when it contains 1 Osm of solute per kilogram of water and it has an osmolality concentration of n when it contains n osmol per kilogram of water. It reflects a weight-to-weight relationship between a solute and a solvent and is a counterpart of molal solutions.
Osmolarity is the expression of osmolar concentration. A solution possesses an osmolar concentration of one when it contains 1 Osm of solute per litre of solution and it has an osmolarity of n when it has n Osm per litre of solution. It represents a weight-to-volume relationship between solute and final solution and is a counterpart of molar solutions.
Multiple choice questions (MCQs)
1.The pH of the buffer solution depends upon the concentration of
a)Acid (H+) only
b)Conjugate base (OH–) only
d)Acid (H+) and Conjugate base (OH–)
2.The buffers present in the blood contain
d)All of the above
3.Range of pH scale is
a)7 to 10
b)0 to 10
c)0 to 14
d)7 to 14
4.Level of pH found in antacid solution
5.pH of neutral salt is
6.When more and more water is diluted with acids its H+ion concentration will
c)remains the same
d)depends on the type of acids
7.Which of the following is not a simple buffer?
c) H3PO4 + NaH2PO4
d) (NH4)2 CO3
8.Which of the following is not a type of Acidic buffer mixture?
a) Na2HPO4+ Na3PO4
b) CH3COOH+ CH3COONa
9.Which of the following is not a type of Basic buffer mixture?
d) Glycine + Glycine hydrochloride
10.Which one of the following is equal to the pKaof a weak acid?
a)its relative molecular mass
b)the pKbof its conjugate base
c)the pH of a solution containing equal amounts of the acid and its conjugate base
d)The equilibrium concentration of its conjugate base
11.Which of the following relationships is true for an acidic solution at 25ºC?
a)[H+] > [OH–]
b)pH > 7.00
c)Kw> 1 10-14
d)the solution is negatively-charged
12.Which one of the following relationships is true in water at 25ºC?
a)[H+] = [H2O]
b)[OH–] = [H2O–]
c)Kw> 1 10-14
d)[H+] = [OH–]
13.If the ratio of base to acid in a buffer changes by a factor of 10, the pH of the buffer
a) increases by 2.
b) increases by 1.
c) decreases by 1.
d) decreases by 2.
14.Identify a good buffer
a) significant amounts of both a strong acid and a strong base
b) significant amounts of both a weak acid and a strong acid
c) small amounts of both a weak acid and its conjugate base
d) significant amounts of both a weak acid and its conjugate base
15.Which of the following is TRUE?
a) An effective buffer has very small absolute concentrations of acid and conjugate base.
b) A buffer can not be destroyed by adding too much strong base. It can only be destroyed by adding too much strong acid.
c) A buffer is most resistant to pH change when [acid] = [conjugate base]
d) An effective buffer has a [base]/[acid] ratio in the range of 10 – 100.
- d) Acid (H+) and Conjugate base (OH–)
- d) all of the above
- c) 0 to 14
- b) ≥ 7.0
- a) 7
- b) decrease
- c) H3PO4 + NaH2PO4
- a) Na2HPO4+ Na3PO4
- c) H2CO3+Na2CO3
- c) the pH of a solution containing equal amounts of the acid and its conjugate base
- a) [H+] > [OH–]
- d) [H+] = [OH–]
- b) increases by 1.
- d) significant amounts of both a weak acid and its conjugate base
- c) A buffer is most resistant to pH change when [acid] = [conjugate base]
- GAURAV KUMAR JAIN – THEORY & PRACTICE OF PHYSICAL PHARMACY, 1st edition 2012 Elsevier, page no. 147-151.
- Martins Physical Pharmacy, 6th edition 2011, page no. 315-320.