Applications of Buffers, Buffer Equation and Buffer Capacity

As a result of small amounts of organic acids or alkalis present within crystalline compounds, buffers prevent acid or alkali from altering the pH of a solution.

The necessity of buffers - There are times when preparation and storage of a solution of a defined pH are necessary. This process is even more difficult than making the solution itself. The CO2 in a solution will absorb from the air and cause it to become acidic. Those with alkaline impurities in glass bottles may alter the pH of the solution, thus pharmaceutical solutions are buffered so that the pH will remain relatively constant when even minute amounts of acid or base are added.

Application of buffers

pH levels are usually maintained using biological buffers in the pharmaceutical industry.

1. Stabilize the drug's components: Prevent the gastrointestinal environment from destroying or altering the pH value of essential components, such as aspirin.
2. Purify specific components, such as insulin, by separating them and purifying them.
3. Make sure the drug is solubilized: Some ingredients are only dissolved at specific pH levels.
4. Make sure ingredients are biologically active: Certain ingredients, such as pepsin, cannot maintain their activity below a specific pH level.
5. A drug's ability to be injected into the human body: Injections closely match blood pH values; eye drops closely match conditions around the eye.
6. It inhibits gastric acid production by acting as a buffer by reducing the stomach's acid content.
7. In biological systems - Two buffer systems keep the pH of the blood at 7.4. First is plasma - A primary buffer is found in plasma. There are two kinds of acids/alkalis present in plasma: carbonic acid and sodium carbonate. In erythrocytes, secondary buffers are oxyhemoglobin, hemoglobin, and potassium salts of acid and alkali phosphoric acid.
8. In pharmaceutical systems - Pharmacy utilizes buffers to ensure maximum stability for products by adjusting their pH levels. If you are going to use parenteral preparations (i.e. injections), you should be careful about pH as large deviations are potentially harmful. Parenteral products should have a pH of 7.4, which is the pH of the blood. As part of parenteral products (injections), acetate, phosphate, citrate, and glutamate are often used as buffers.

Buffer equation

Henderson-Hasselbalch Equation is also known to be a buffer equation.

For acid buffers - The pH can be calculated by comparing the dissociation constant K of the weak acid with the concentrations of acid and salt. Here is how dissociation expressions of weak acids look.

HA ↔ H+ + A-

Ka = [H+] [A-] / [HA]

Or [H+] = Ka [HA] / [A-] ------------- (1)

Weak acids are only marginally dissociated, and their dissociation is further depressed by the addition of A-ion (common ion effect), resulting in an equilibrium concentration of the unionized acid that is nearly the same as its initial concentration after unionization. We assume that the equilibrium concentration of A- is equal to the concentration of salt added at the start, due to its complete dissociation. The concentration of A- is represented by the salt concentration in equation (1) above.

[H+] = Ka. [Acid] / [Salt] --------- (2)

If you take the logs on both sides, we attain:

log[H+] = logKa + log [Acid] / [Salt]

when both sides are multiplied by –ve,

-log[H+] = -logKa - log [Acid] / [Salt]

As -log[H+] = pH & -logKa = pka pH = pka - log[Acid] / [Salt] OR pH = pka + log[Salt] / [Acid] ---------- (3)

Equation (3) is known as the Henderson-Hasselbalch equation. By knowing the acid and salt concentrations, it assists in calculating a buffer solution's pH value.

For basic buffers - Based on the same principle as for acidic buffers, we can calculate the buffer equation for basic buffers. Suppose a basic buffer is made up of a mixture of weak organic acids (BOH) and their salts (BA). We can write the dissociation constant for the base as follows:

BOH ↔ B+ + OH- Kb = [B+] [OH-] / [BOH] OR [OH-] = Kb [BOH] / [B+] ------------- (1)

In other words, the balance between equilibrium concentrations of the unionized base and its initial concentrations is close. This is because a weak base is only slightly dissociated by the salt, BA, providing the B+ ion (common ion effect). Since the salt has been completely dissociated by the time it reaches equilibrium, B+ is assumed to have the same equilibrium concentration as the initially added salt. Accordingly, in equation (1), the salt concentration is used as a measure of B+ concentration.

[OH-] = Kb. [Base] / [Salt] --------- (2)

By adding the two logs together, we get:

log[OH-] = logKb + log [Base] / [Salt]

Adding the –ve sign to both sides,

-log[OH-] = -logKb - log [Base] / [Salt]

As -log[OH-] = pOH & -logKb = pkb pOH = pkb – log [Base] / [Salt] Or pOH = pkb + log[Salt] / [Base] -----(3)

Buffer capacity

Adding acid or base to a buffer solution will change its buffer capacity, which measures its resistance to pH changes. Besides buffer capacity, buffer value, buffer efficiency, and buffer coefficient are other terms for buffer capacity. The buffer capacity for which the symbol 'β' stands may also be defined as follows:

The ratio between a small increase in acidity or alkalinity (amount added) and the small change in pH (ΔpH) that results from its addition.

β = ΔA or ΔB / ΔpH

You can change the pH of a buffer by mixing gradations of strong acid or base (in gram equivalents or liters) with it. It is necessary to add one gram of a strong acid or base to one liter of a solution to adjust its pH by one unit. Therefore, if the pH level of a solution changes less when acid or base is added, the buffer capacity is higher and vice versa.

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