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Buffering agents are substances that adjust the pH of a solution. Buffering agents can be either weak acids or weak bases that make a buffer solution. These substances are usually added to water to form buffer solutions and are responsible for the buffering action seen in these solutions. The objective of a buffer is to keep the pH of a solution within a narrow range. The function of a buffering agent is to drive an acidic or basic solution to a certain pH state and prevent a change in pH. For example, buffered aspirin has a buffering agent magnesium oxide that maintains the pH of the aspirin as it passes through the stomach of the patient. The monopotassium phosphate also is an example of a buffering agent. Buffering agents are primarily used to lower the acidity of the stomach, for example, antacid tablets. These agents have variable properties that have wide differences in solubility and acidity characteristics. As pH controllers, they are important in medicine. The buffering agents work similarly to buffer solutions. As we know to avoid the little change in the concentration of the acid and base the solution is buffered. A buffering agent upon addition by providing the corresponding conjugate acid or base sets up such a concentration ratio that stabilizes the pH of that solution. The resulting pH of this combination can be calculated using the Henderson-Hasselbalch equation. Buffering agents are the main and active components of buffer solutions. They both regulate the pH of a solution as well as resist changes in pH.
A buffer solution maintains the pH for the whole system in which it is placed, whereas a buffering agent is added to an already acidic or basic solution, which is modified to maintain a new pH. Buffering agents and buffer solutions are similar except for a few differences that buffer solution maintains the pH of a system by preventing large changes in it, whereas agents modify the pH of what they are placed into. Buffering agents in humans, functioning in acid-base homeostasis, are extracellular agents, for example, bicarbonate, ammonia as well as intracellular agents including proteins and phosphate.
Buffering agents (Buffer salts) and buffer solutions (buffer systems) have different applications that improve stability (for example, aspartame), control gelling (for example, pectin-based products), reduce the rate of reaction (for example, sucrose inversion), and reduce variation in pH. Therefore, the color, flavor (for example, foods and beverages), and texture (for example, gelled products) are maintained. A buffer can be made by partially neutralizing a weak acid like citric or malic acid with sodium hydroxide. However, sodium hydroxide, or caustic soda, is both hygroscopic and hazardous. Instead of using sodium hydroxide, salts of weak acids such as trisodium citrate, sodium lactate, trisodium phosphate, or sodium acetate are used to partially neutralize the acid. Since they contribute to the buffer capacity themselves, these salts are buffer salts.
As shown in Fig.1, the variation in pH from lot to lot is reduced after the addition of a buffer salt. The buffer salts increase the buffer capacity of the buffer system and stabilize pH.
Preparing Buffer Solutions
The simplest way of preparing a buffer solution is dissolving a known quantity of the salt of the weak acid (or base) in a solution of a weak acid (or base) of known concentration. Another way is to neutralize an excess of a weak acid (or weak base) with some strong base (or strong acid). The neutralization produces the salt of the weak acid (or base) ‘in situ’. As the weak acid is in excess, there will still be some weak acid in the mixture. The resultant mixture contains both the salt of the weak acid and the weak acid itself.
Weak acid and a salt of acids conjugate base in sufficient amounts are required to maintain the ability of buffer. Citric acid-phosphate buffer, for example, is prepared by adding 0.1 M citric acid to 0.2 M disodium phosphate (Na2HPO4) solution followed by mixing to make 100 mL solution. The total amounts of these solutions of specific strength required to make buffer solution of particular pH are given in Table.1.
General considerations for preparing buffers:
1. Determine the optimal pH for the product, based on physical and chemical stability, therapeutic activity, and patient comfort and safety (must consider chemical and physical nature of the active and other ingredients and the route of administration).
2. Select a weak acid with a pKa near the desired pH (must be non-toxic and physically/chemically compatible with other solution components).
3. Calculate the ratio of salt to acid required to produce the desired pH (use Henderson-Hasselbalch equation).
4. Determine desired buffer capacity of the product (consider the stability of the product, route of administration, the volume of dose, and chemical nature of product).
5. Calculate the total buffer concentration required to produce desired buffer capacity (Van Slyke equation).
6. Determine the pH and the buffer capacity of the prepared buffer solution by using a suitable method.
There are four commonly used methods to prepare buffer solutions:
1. The Slow and Stupid Method:
A buffer composed of an acid and its salt is prepared by dissolving the buffering agent (acid form) in about 60% of the water required for the final solution volume. The pH is adjusted using a strong base, such as NaOH. To prepare a buffer composed of a base and its salt, start with the base form and adjust the pH with strong acids, such as HCl. When the pH is correct, dilute the solution to just under the final volume of the solution. Check the pH and correct if necessary, then add water to make the final volume. This method is easy to understand but is slow and may require lots of bases (or acid). If the base (or acid) is concentrated, it is easy to increase the pH. If the base (or acid) is dilute, it is easy to increase the volume. Adding a strong acid or base can result in temperature changes, which make pH readings inaccurate (due to its temperature dependence) unless the solution is brought back to its initial temperature.
2. The Mentally Taxing Method:
In this method using buffer pKa, the amounts (in moles) of acid/salt or base/salt present in the buffer at the desired pH are calculated. If both forms (i.e., the acid and the salt) are available, the amount required is converted from moles to grams, using the molecular weight of that component. The correct amounts of both forms are weighed and used. If only one form is available then the buffer is prepared by adding the entire buffer as one form, and then acid or base is added to convert some of the added buffers to the other form. Once the total concentration of buffer in the solution is decided, it is converted to amount (in moles) using the volume of solution and then to grams, using the molecular weight of the buffer.
The amounts (in moles) of each form that will be present in the final solution are calculated using the buffer pKa and the desired pH. Then the amount of strong acid or base that must be added to give the correct amounts of each form at the pH of the final solution is calculated. The buffer and strong acid or base are dissolved in slightly less water than is required for the final solution volume. The pH is checked and corrected if necessary. Water is added to make up the final volume. It is a fast method and easy to prepare. This method requires the buffer pKa value. Additional pH adjustment is rarely necessary, and when needed, the adjustment is small.
3. The Two Solution Method:
The separate solutions of the acid form and base form of the buffer are prepared from solutions having the same buffer concentration. To obtain the desired pH, one solution is added to the other with continuous monitoring of the pH. This method is easy to do but requires both forms of the buffer. The required solution volumes are proportional to the ratio of buffer components in the final solution at the desired final pH.
4. The Completely Mindless Method:
The correct amounts of acid or its salt or base or its salt required for different pH values are selected from the standard data value tables and the same amounts of the components are dissolved in the slightly less water than is required to make the final volume of solution. The pH is checked and corrected if required followed by adjusting the final volume by adding water. This method is easy to do because of the use of a suitable reference table. It is a convenient method for frequently prepared buffers but it may be difficult to find such a table. This method requires both forms of the buffer. Components amounts from the table need to be adjusted to produce the required buffer concentration and volume.
5. Alternative Method for Preparing Buffer Solutions:
This method is used rarely. In this method rather than mixing the weak acid with its salt, a buffer solution is prepared by adding a limited amount of strong base to the weak acid to solve the weak acid (or base) and its conjugate base (or acid) which results in the weak acid and the salt of the weak acid.
Standard Buffer Solutions:
The standard buffer solution of pH ranging from 1.2 to 10 is possible to prepare by appropriate combinations of 0.2 N HCl or 0.2 N NaOH or/and 0.2 M solutions of potassium hydrogen phthalate, potassium dihydrogen phosphate, boric acid-potassium chloride as given in pharmacopeia. The pH range and the quantities of the ingredients used to make the respective standard buffer at 25 °C are given in Table.2. Buffers have great use in biological research. Various criteria that can be applied while making buffers for this application are listed below.
1. Buffers must possess enough buffer capacity in the required pH range.
2. It must be available in highly purified form.
3. It must be highly water-soluble and impermeable to biological membranes.
4. It must be stable especially concerning hydrolysis and enzymatic action.
5. It must maintain pH which is influenced to a very small value by their concentration, temperature, and ionic strength as well as the salting-out effect of the medium.
6. It must be non-toxic with no biological inhibition activity.
7. Buffers must not form complexes.
8. It must not absorb light in the visible or ultraviolet regions.
9. It must not precipitate in redox reactions.
10. It must not alter the solubility of active ingredients.
11. It must be safe to use in biological systems and do not alter the pharmacological responses of the active ingredients.
Generally, buffers are used in pharmaceutical products for two purposes viz. to adjust the pH of product for maximum stability and to maintain the pH within the optimum physiological pH range. Pharmaceutical solutions generally have a low buffer capacity to prevent overwhelming the body’s buffer systems and significantly changing the pH of the body fluids. Buffers have concentrations in the range of 0.05 to 0.5 M and buffer capacities in the range of 0.01 to 0.1 which are usually sufficient for pharmaceutical solutions. Table.3 gives some of the buffer systems used in the pharmaceutical formulations along with their pKa values. Most pharmaceutical buffers are composed of ingredients that are found in the body (for example, acetate, phosphate, citrate, and borate). While selecting, the right pharmaceutical buffers choose a weak acid with pH > pKa. Carry out calculations using buffer equation to the determination of acid/base needed to give required pH. Also, choose the proper concentration needed to give suitable buffer capacity. The ingredients are selected from available ones considering their sterility, stability, cost, toxicity, etc.
Buffers in Pharmaceuticals:
1. Solid dosage forms: Buffers have been used widely in solid dosage forms such as tablets, capsules, and powders for controlling the pH of the environment around the solid particles. This has practical application for the drugs that have dissolution rate-limited absorption from unbuffered solutions. One of the special applications of buffers is to reduce the gastric irritation caused by acidic drugs. For example, sodium bicarbonate, magnesium carbonate, and sodium citrate antacids, used for reducing acidity.
2. Semisolid formulations: Semisolid preparations such as creams and ointments undergo pH changes upon storage for a long time resulting in their reduced stability. Hence buffers such as citric acid and sodium citrate or phosphoric acid/sodium phosphate are included in these preparations to maintain their stability.
3. Parenteral products: The use of buffers is common in parenteral products. Since the pH of blood is 7.4 these products are required to be adjusted to this pH. Change in pH to the higher side (more than 10) may cause tissue necrosis while on the lower side (below 3) it may cause pain at the site of action. As blood, itself functions as a buffer, adjustment of pH for small volume parenteral preparations is not required. Commonly used buffers include citrate, glutamate, phthalate, and acetate. The pH optimization is generally carried out to have better solubility, stability, and reducing irritancy of the product.
4. Ophthalmic products: Many drugs, such as alkaloidal salts, are most effective at pH levels that favor the undissociated free bases. However, at such pH levels, the drug may be unstable Therefore such pH levels must be obtained by use of buffers. The purpose of buffering some ophthalmic solutions is to prevent an increase in pH caused by the slow release of hydroxyl ions by the glass. Such a rise in pH can affect both the solubility and the stability of the drug. The decision of whether buffering agents should be added in preparing an ophthalmic solution must be based on several considerations. Normal tears have a pH of about 7.4 and possess some buffer capacity.
The application of a solution to the eye stimulates the flow of tears and the rapid neutralization of any excess hydrogen or hydroxyl ions within the buffer capacity of the tears. Many ophthalmic drugs are weakly acidic and have only weak buffer capacity. Where only 1 or 2 drops of a solution containing them are added to the eye, the buffering action of the tears is usually adequate to raise the pH and prevent marked discomfort. In some cases, pH may vary between 3.5 and 8.5. Some drugs, notably pilocarpine hydrochloride and epinephrine bitartrate, are more acid and overtax the buffer capacity of the lachrymal fluid. Ideally, an ophthalmic solution should have the same pH, as well as the same isotonicity value, as lachrymal fluid. This is not usually possible since, at pH 7.4, many drugs are not appreciably soluble in water.
Most alkaloidal salts precipitate as these exist as free alkaloids at this pH. Additionally, many drugs are chemically unstable at pH levels approaching 7.4. This instability is more marked at the high temperatures employed in heat sterilization. For this reason, the buffer system should be selected that is nearest to the physiological pH of 7.4 and does not cause precipitation of the drug or its rapid deterioration.
An ophthalmic preparation with a buffer system approaching the physiological pH can be obtained by mixing a sterile solution of the drug with a sterile buffer solution using an aseptic technique. Even so, the possibility of a shorter shelf-life at the higher pH must be taken into consideration, and attention must be directed toward the attainment and maintenance of sterility throughout the manipulations. Boric acid is often used to adjust isotonicity in ophthalmic solutions because of its buffering and anti-infective properties.
Many drugs, when buffered to a therapeutically acceptable pH, would not be stable in solution for long periods. Hence these products are lyophilized and are intended for reconstitution immediately before use, for example, Acetylcholine Chloride Ophthalmic Solution.
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