What is Buffer solution physical pharmaceutics1 notes

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Buffer solution definition:-A buffer solution in which resists changes in pH when small amount of an acid and or an alkali are added to it.

Buffer Solution Definition, notes by phrmacypdfnotes

Buffer solution

buffer solution (more precisely, pH buffer or hydrogen ion buffer) is an aqueous solution that combines a weak acid mixture with its conjugate base, or vice versa. Its pH changes very little when a small amount of solid acid or foundation is added to it. Buffer solutions are used as a way to keep the pH at almost equal value in various chemical applications. In nature, there are many systems that use bath storage to control pH. For example, the bicarbonate buffering system is used to control blood pH, and bicarbonate also acts as a barrier at sea.

The principles of the buffers

Buffer solutions are resistant to changes in pH due to the balance between the weak acid HA and its base conjugate A-:

HA ⇌ H + + A−

When a certain solid acid was added to an equal mixture of weak acid and its base conjugate, hydrogen ions (H +) were added, and the balance was shifted to the left, according to Le Châtelier's principle. As a result, the concentration of hydrogen ion increases below the expected value in the amount of solid acid added. Similarly, when solid alkali is added to a mixture, the concentration of hydrogen ion drops below the expected value of the added alkali value. The result is indicated by the comparable titration of weak acid with pKa = 4.7. The associated concentration of unmixed acid is indicated in blue, with its conjugate base in red. The pH changes slightly compared to the database, pH = pKa ± 1, which focuses on pH = 4.7, where [HA] = [A-]. The concentration of hydrogen ion decreases below the expected value because most of the extra hydroxide ions are used in the reaction.

OH− + HA → H2O + A−

and only a small amount was used in the neutral response (which is a reaction that leads to an increase in pH)

OH− + H + → H2O.

When the acid is over 95% deprotonated, the pH rises rapidly because most of the extra alkali is consumed in a neutral reaction.

The capacity of the buffer

Buffer volume is a measure of the resistance to change of pH of a solution containing an agent that maintains interference with respect to acid changes or alkaline concentrations. It can be described as follows:

where  is an infinitesimal amount of added base, or

where  is an infinitesimal amount of added acid. pH is defined as −log₁₀[H⁺], and d(pH)

 it is a constant amount of extra acid. The pH is defined as −log10 [H +], and d (pH) is an infinite pH change.

By any definition the buffer capacity of weak acid HA with a consistent dissociation Ka can be expressed as

where [H +] is the concentration of hydrogen ions, the total amount of acid added. Kw is a constant consistency of artificial ionization of water, equivalent to 1.0 × 10−14. Note that in solution H + exists as a hydronium ion H3O +, and the further aquation of hydronium ion has a negative effect on the separation ratio, without the concentration of very high acid.

In the central region of the curve (green in the structure), the second word is dominant, 

$${\displaystyle \beta \approx 2.303{\frac {T_{{\ce {HA}}}K_{a}[{\ce {H+}}]}{(K_{a}+[{\ce {H+}}])^{2}}}.}$$

The buffer volume increases with local magnitude to pH = pKa. The height of this peak depends on the pKa value. The buffer volume is ignored when the [HA] concentration of the buffer agent is very low and increases with the increasing concentration of the buffer agent. Some authors show this region only in bar volume graphs.

The volume of the buffer drops to 33% of the maximum pH = pKa ± 1, to 10% to pH = pKa ± 1.5 and to 1% to pH = pKa ± 2. For this reason the most widely used range is almost is pKa ± 1. When selecting a database to be used for a particular pH, it should have a pKa value that is very close to that pH.

With strong acid solutions, pH less than 2 (red in color), the first term in the equation dominates, and the buffer volume increases significantly with a decrease in pH:

$${\displaystyle \beta \approx 10^{-\mathrm {pH} }.}$$

This results in the second and third terms becoming insignificant with very low pH. This term is independent of the presence or absence of an agent that limits volume.

With strong alkaline solutions, the pH is above 12 (blue in the structure), the third term in the equation dominates, and the buffer volume increases significantly with increasing pH:

$${\displaystyle \beta \approx 10^{\mathrm {pH} -\mathrm {p} K_{\text{w}}}.}$$

This is because the first and second terms become insignificant with very high pH. The term also stands for the presence or absence of a buffer agent.

Buffer applications

The pH of the solution containing the bath agent can only vary within a small distance, regardless of what else may be present in the solution. In biological systems this is an important condition for the enzymes to function properly. For example, in human blood a mixture of carbonic acid (H2CO3)

3) and bicarbonate (HCO3-)

3) is present in the plasma fraction; this creates a great way to maintain blood pH between 7.35 and 7.45. Without this subdivision (7.40 ± 0.05 pH unit), acidosis and alkalosis metabolic conditions develop rapidly, eventually leading to death if the proper bath dose is not restored immediately.

When the pH value of the solution rises or decreases significantly, enzyme efficiency decreases in a process, known as denaturation, which is usually irreversible. Most of the biological samples used in the study are stored in a buffer solution, usually phosphate buffered saline (PBS) at pH 7.4.