A buffer is an aqueous solution that consists of a mixture of a weak acid and its salt (acid buffer) or a weak base with its salt (basic buffer). Its pH changes very little when a small amount of strong acid or base is added to it and is thus used to prevent a solution’s pH change.
Buffer solutions are used for a wide range of chemical applications. Blood is one example of a buffer solution found in nature. Human blood has a natural pH of 7.4. Many people experience severe anxiety and suffer from alkalosis. Alkalosis is a disease in which blood pH is excessively high. The reverse condition is called acidosis-a blood, pH greater than 7.4
Some chemical reactions only occur at a certain pH. Other households and consumer items need to monitor their pH values, such as shampoo to combat the soap’s alkalinity to avoid inflammation, baby lotion to retain a pH of around 6 to discourage multiplication of bacteria, washing powder, eye drops, fizzy lemonade etc.
Buffer Solution Definition
Solutions with the stable concentration of hydrogen ions and thus typically with no change in pH which is almost independent of dilution and which change very little with small additions of a strong acid or alkali are called buffers. It can also be described in simple terms as a solution that prevents any pH change when a small amount of a strong acid or a strong base is applied to it, which is called a buffer solution or simply as a buffer. Both buffers have acidity and alkalinity balance.
Any compounds, such as ammonium acetate, tend to resist any change in their concentration of hydronium ions or pH, whenever a small amount of a strong acid or a strong base is applied to it.
Buffer solutions usually consist of a mixture of a weak acid and salt with a strong base like CH3COOH and CH3COONa, or a weak base with a strong acid like NH4OH and NH4Cl and salt.
Mechanism of Buffering Action
Consider the example of a buffer solution made by dissolving sodium acetate into acetic acid, to consider how a buffer functions. As you can see from the name, acetate acid is an acid: CH3COOH, while sodium acetate dissociates in solution to yield the conjugate base, CH3COO-acetate ions. The reaction equation is:
CH3COOH (aq) + OH-(aq) 🡪CH3COO-(aq) + H2O (aq)
If this solution is combined with a strong acid, the acetate ion can neutralise.
CH3COO-(aq) + H+(aq) 🡪CH3COOH (aq)
It changes the original buffer reaction equilibrium, thereby holding the pH steady.
Preparation of Buffer Solution
There are a few methods to prepare a buffer solution with a different pH. Prepare a solution with acid and its conjugate base in the first approach by dissolving the acid component of the buffer in around 60 per cent of the amount of water used to produce the final volume of solution.
Instead, use a pH detector to test the pH of the solution. Using a strong base like NaOH the pH can be changed to the desired value. If a base and its conjugate acid are used to make the buffer, the pH can be modified using a strong acid, like HCl. Dilute the solution to the final desired volume, once the pH is right.
Additionally, you should prepare solutions for both the solution’s acid type and base form. Both solutions must have the same quantity of buffer as in the final solution. Add one solution to the other while tracking the pH to get the final buffer.
In a third method, using the Henderson-Hasselbach equation, you can determine the exact amount of acid and conjugate base required to make a buffer of a certain pH:
 pH = pKa + log|A−||HA|
Types of Buffer Solution
There are two buffer forms, acid buffer, and base buffer.
Acid Buffer
A buffer solution that contains large quantities of a weak acid, and its salt with a strong base, is called an acid buffer. On the acidic side, such buffer solutions have pH, i.e.pH is below 7 at 298 K. The equation gives the pH of an acid buffer. CH3COOH, with CH3COONa.
pH = pKa + ln(Salt)Acid
Where Ka —–acid dissociation constant of the weak acid
Basic Buffer
A buffer solution that contains relatively large quantities of a weak base and its salt with a strong acid is called a simple buffer. On the alkaline side, these buffers have pH, i.e., pH is higher than 7 at 298 K. For example, NH4OH and NH4Cl.
The pH of an appropriate buffer is determined by the equation
pOH = pKb + ln(Salt)Acid
Where, Kb ——base dissociation constant.
These equations are called Henderson Hasselbalch equations
Buffer Solution Examples
Blood – contains a bicarbonate buffer system
Tris buffer
Phosphate buffer
As mentioned, buffers are beneficial over specific pH ranges. For example, here is the pH range of common buffering agents:
While making a buffer solution, the pH of the solution is changed to get it within the right effective range. A strong acid, such as hydrochloric acid (HCl), is usually added to reduce the pH of acidic buffers. A strong base such as sodium hydroxide (NaOH) solution is added to increase the pH of the alkaline buffers.
Importance of Buffers
The acidity of the solution in which they occur affects a lot of chemical reactions. The pH of the reaction medium must be controlled for a given reaction to occur or to occur at a suitable rate. This control is provided by buffer solutions, which are solutions that preserve a certain pH. Biochemical reactions are particularly sensitive to pH. Most biological molecules contain groups of atoms that can be charged or neutral based on pH, and whether these groups are charged or neutral has a significant effect on the molecule’s biological activity.
The fluid within the cell and the fluids around the cells have a characteristic and almost constant pH in all multicellular organisms. This pH is preserved in several ways, and one of the most important is through buffer systems.