Buffer Solutions: A steady \({\rm{pH}}\) is required for the proper functioning of many chemical and biological systems, including our blood, for reactions to take place. Buffers are required in such systems in order to maintain a consistent \({\rm{pH}}.\) An aqueous solution comprising a weak acid and its conjugate base, or a weak base and its conjugate acid, is known as a buffer solution.

A buffer solution is used to prevent the drastic change in the \({\rm{pH}}\) value of a solution as it causes a minimal change in the \({\rm{pH}}\) on the addition of strong acid or a strong base in a solution. Due to its unique property, the buffer is widely used in different experiments carried out in laboratories. Buffers are broadly classified into two types – Acidic and Alkaline buffer solutions. Solutions having \({\rm{pH}}\) below \(7\) are known as “Acidic buffers.” While solutions having \({\rm{pH}}\) above \(7\) are known as “Alkaline buffers” or “Basic buffers.” In this article, let’s learn everything about Buffer Solutions in detail. Read on to find more.

Buffer Solution Definition

A solution that can resist any change in \({\rm{pH}}\) upon adding an acidic or alkaline component in a solution is defined as a buffer solution. It maintains a stable \({\rm{pH}}\) range of a particular reaction or process as a buffer solution can neutralize the solution when a small amount of acid or base is added to it. 

A buffer solution should meet these two criteria:

1. Weak acid or weak base – A buffer should compose of a weak acid or a weak base. For example, acetic acid, nitrous acid, ammonia, etc.
2. Conjugate salt – Another essential criterion of a buffer is that it should have a conjugate salt of that weak acid or weak base. A conjugate salt differs from its original acid or base by one \({{\rm{H}}^{\rm{ + }}}\) ion or one \(\mathrm{OH}^{-}\) ion.

Thus, examples of buffer solutions are \({\rm{HF}}\) and \({\rm{KF}}\), Acetic acid and Sodium acetate, Ammonia and Ammonium chloride, etc.

Types of Buffer Solution

Buffer solutions are broadly characterized into two types:

1. Acidic Buffer – The buffer solution that maintains the acidic environment is known as an acidic buffer. An acidic buffer is a mixture of a weak acid and its salt with a strong base. The \({\rm{pH}}\) value of the acidic buffer is below \(7\) at room temperature. For example, acetic acid \(\left(\mathrm{CH}_{3} \mathrm{COOH}\right)\) and sodium acetate \(\left(\mathrm{CH}_{3} \mathrm{COONa}\right)\) solutions are mixed to form an acidic buffer with \({\rm{pH}}\) \(4.74.\) 

2. Alkaline or Basic Buffer – The buffer solution that maintains the basic environment of a solution is known as an alkaline or basic buffer. An alkaline buffer is a mixture of a weak base and its salt with a strong acid. The \({\rm{pH}}\) value of acidic buffer is above \(7\) at room temperature. For example, ammonium hydroxide \(\left(\mathrm{NH}_{4} \mathrm{OH}\right)\) and ammonium chloride \(\left(\mathrm{NH}_{4} \mathrm{Cl}\right)\) solutions are mixed to form alkaline or basic buffer with a \({\rm{pH}}\) of around \(10.\) 

Buffer Action Mechanism

The property of the solution to resist the changes in its \({\rm{pH}}\) value on the addition of a minimum amount of strong acid or strong base is known as “Buffer Action.”

Mechanism of Acidic Buffer Solution

To understand the acidic buffer action, let us take an example of a weak acid (acetic acid) and its conjugate base (sodium acetate). In the solution, the weak acid and its conjugate base will ionize as shown below:

\({\rm{C}}{{\rm{H}}_3}{\rm{COOH}}({\rm{aq}}) \leftrightarrow {\rm{C}}{{\rm{H}}_3}{\rm{CO}}{{\rm{O}}^ – }({\rm{aq}}) + {{\rm{H}}^ + }({\rm{aq}})\)

\(\mathrm{CH}_{3} \mathrm{COONa}(\mathrm{aq}) \leftrightarrow \mathrm{CH}_{3} \mathrm{COO}^{-}(\mathrm{aq})+\mathrm{Na}^{+}(\mathrm{aq})\)

Thus, the acidic buffer contains both acid \(\left(\mathrm{CH}_{3} \mathrm{COOH}\right)\) and its conjugate base \(\left(\mathrm{CH}_{3} \mathrm{COO}^{-}\right) .\)

When a small quantity of a strong acid is added to the buffer solution, the \({{\rm{H}}^{\rm{ + }}}\) ions of the acid are removed by the conjugate base \(\left(\mathrm{CH}_{3} \mathrm{COO}^{-}\right)\) by forming acetic acid that is already present in the solution. Thus, the resultant \({\rm{pH}}\) change of the solution will be almost constant.

\({\rm{C}}{{\rm{H}}_3}{\rm{CO}}{{\rm{O}}^ – }({\rm{aq}}) + {{\rm{H}}^ + }({\rm{aq}}) \leftrightarrow {\rm{C}}{{\rm{H}}_3}{\rm{COOH}}({\rm{aq}})\)

If, in contrast, a strong base is added to the buffer solution, the added \(\mathrm{OH}^{-}\)  ion is neutralized by the reaction with the acid in the buffer because, \(\mathrm{OH}^{-}\) ion reacts with the \({{\rm{H}}^{\rm{ + }}}\) ions and produces water. Again, the resultant \({\rm{pH}}\) change of the solution will remain almost constant.

\(\mathrm{CH}_{3} \mathrm{COOH}(\mathrm{aq})+\mathrm{OH}^{-}(\mathrm{aq}) \leftrightarrow \mathrm{CH}_{3} \mathrm{COO}^{-}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})\)

Mechanism of Basic Buffer Solution

To understand the basic buffer action, let us take an example of an Ammonia buffer solution. 

In addition of acid, \({{\rm{H}}^{\rm{ + }}}\) combines with the base \(\mathrm{NH}_{3}\) to form \({\rm{NH}}_4^ + \). Thus, removing most of the added hydrogen ions. 

\({\rm{N}}{{\rm{H}}_{\rm{3}}}({\rm{aq}}){\rm{ + }}{{\rm{H}}^{\rm{ + }}}({\rm{aq}}) \leftrightarrow {\rm{NH}}_4^ + ({\rm{aq}})\)

On the other hand, on the addition of base, \(\mathrm{OH}^{-}\) combines with the acid \(\left(\mathrm{NH}_{4}{ }^{+}\right)\) to form ammonia, and water. As a result, most of the added hydroxyl ions are removed. So, in this manner, the buffer solution keeps the \({\rm{pH}}\) of the buffered solution relatively constant.

\({\rm{NH}}_4^ + ({\rm{aq}}) + {\rm{O}}{{\rm{H}}^ – }({\rm{aq}}) \leftrightarrow {\rm{N}}{{\rm{H}}_3}({\rm{aq}}) + {{\rm{H}}_2}{\rm{O}}({\rm{l}})\)

The pH of Buffer Solution

Buffer solutions contain a weak acid, and it’s a conjugate base. Therefore, they can absorb excess \({{\rm{H}}^{\rm{ + }}}\) ions or \(\mathrm{OH}^{-}\) ions, as a result, they can maintain an overall steady \({\rm{pH}}\) in the solution. 

There are two factors that determine the pH of a buffer:

1. The equilibrium constant \({{\rm{K}}_{\rm{a}}}\) of the weak acid
2. The ratio of a weak base \(\left[A^{-}\right]\) to weak acid \({\rm{[HA]}}\) in solution. 

The \({\rm{pH}}\) of a buffer solution is calculated by using the Henderson-Hasselbalch Equation with a known amount of acid and its conjugate base. 

Henderson-Hasselbalch Equation Derivation

When an aqueous solution of acid \({\rm{HA’}}\) is dissociated to form \({{\rm{H}}^{\rm{ + }}}\) ion and its conjugate base, i.e., \({{\rm{A}}^{\rm{ – }}}\)

The equilibrium constant for such ionization reaction is denoted as \({{\rm{K}}_{\rm{a}}}\)

\({{\rm{K}}_{\rm{a}}}{\rm{ = }}\frac{{\left[ {{{\rm{H}}^{\rm{ + }}}} \right]\left[ {{{\rm{A}}^{\rm{ – }}}} \right]}}{{{\rm{[HA]}}}}\)

On rearranging the above equation and taking the negative logarithm of LHS and RHS, we get:

\({\rm{ – log}}\left( {\frac{{\left[ {{\rm{HA}}} \right]}}{{\left[ {{{\rm{A}}^{\rm{ – }}}} \right]}}} \right){\rm{ – log}}{{\rm{K}}_{\rm{a}}}{\rm{ = – log}}\left[ {{{\rm{H}}^{\rm{ + }}}} \right]\,…\left( {\rm{i}} \right)\) 

We know that \({\rm{pH =\, – log}}\left[ {{{\rm{H}}^{\rm{ + }}}} \right]\) and \({\rm{p}}{{\rm{K}}_{\rm{a}}}{\rm{ =\,  – log}}{{\rm{K}}_{\rm{a}}}\)

Substituting \({\rm{pH}}\) and \({\rm{p}}{{\rm{K}}_{\rm{a}}}\) in equation (i), we get:

\({\rm{ – log}}\left( {\frac{{{\rm{[HA]}}}}{{\left[ {{{\rm{A}}^{\rm{ – }}}} \right]}}} \right){\rm{ + p}}{{\rm{K}}_{\rm{a}}}{\rm{ = pH}}\)

Thus, the Henderson-Hasselbalch Equation is:

\({\rm{pH = p}}{{\rm{K}}_{\rm{a}}}{\rm{ + log}}\left( {\frac{{\left[ {{{\rm{A}}^{\rm{ – }}}} \right]}}{{{\rm{[H,A]}}}}} \right)\)

It can also be written as:

\({\rm{pH = p}}{{\rm{K}}_{\rm{a}}}{\rm{ + log}}\left( {\frac{{\left[ {{\rm{Salt}}} \right]}}{{{\rm{[Acid]}}}}} \right)\)

By using the above equation, the pH of a buffer solution can be calculated.

Buffer Isotonic Solution

Buffer solutions needed for biological or pharmaceutical purposes must be isotonic also because these buffer solutions in the form of drugs should match with the concentration of the fluids inside our body. 

Two solutions having the same osmotic pressure across a semipermeable membrane are defined as isotonic solutions. Thus, the addition of a small quantity of a buffer solution to an isotonic solution will change its isotonicity since it is the property of the number of particles in a solution. For example, the concentration of our blood is \({\rm{0}}{\rm{.9\% }}\frac{{\rm{w}}}{{\rm{v}}}\) of \({\rm{NaCl}}.\) So, if we need a solution to be injected into the blood, it should match the tonicity of blood, i.e., \({\rm{0}}{\rm{.9\% }}\frac{{\rm{w}}}{{\rm{v}}}\) of \({\rm{NaCl}}.\) Thus, buffer isotonic solutions are essential for pharmaceutical purposes.

Uses of Buffer Solution

1. Buffer solutions are used to maintain a constant \({\rm{pH}}\) value of a solution under chemical reactions. This is widely used in many chemical laboratories for carrying out numerous chemical research analyses.
2. It plays a major role in the biological system. As we know, Biochemical reactions, including enzymatic activities, are very sensitive to \({\rm{pH}}\). Thus, these biochemical reactions need to be carried out in a constant \({\rm{pH}}\) environment.
3. Human blood contains a buffer of carbonic acid \(\left(\mathrm{H}_{2} \mathrm{CO}_{3}\right)\) and bicarbonate anion \(\left.\mathrm{HCO}_{3}^{-}\right)\) to maintain the \({\rm{pH}}\) of the blood. Phosphate buffer is very useful in maintaining the \({\rm{pH}}\) of blood around \(7.4.\)
4. The rate of a particular reaction also depends upon the \({\rm{pH}}\) of the reaction medium. Therefore, this control is provided by buffer solutions that preserve the pH level.
5. Buffered isotonic solutions are widely used in pharmaceutical industries to avoid any sudden \({\rm{pH}}\) change inside the body due to drug administration.
6. In cosmetic products such as shampoos, body lotions, detergents, etc., \({\rm{pH}}\) balance is very important, and this is done by the buffer solution.

Summary

In short, it can be said that buffer solutions are the solutions that resist pH variation in a particular solution. A buffer must contain a weak acid or a weak base along with its conjugate salt. For example, Acetic acid and Sodium acetate, Ammonia and Ammonium chloride, etc., make buffer solutions. Buffers are broadly divided into two types – Acidic buffers and Alkaline buffers.

Solutions having pH below \(7\) are known as “Acidic buffers,” such as a solution of acetic acid \({\rm{(C}}{{\rm{H}}_3}{\rm{COOH}})\) and sodium acetate \(\left(\mathrm{CH}_{3} \mathrm{COONa}\right) .\) While solutions having \({\rm{pH}}\) above \(7\) are known as “Alkaline buffers” or “Basic buffers” such as ammonium hydroxide \(\left(\mathrm{NH}_{4} \mathrm{OH}\right)\) and ammonium chloride \(\left(\mathrm{NH}_{4} \mathrm{Cl}\right)\) solution. 

The \({\rm{pH}}\) of a buffer solution is calculated by using Henderson-Hasselbalch Equation. \({\rm{pH}}\) plays a vital role in the biological system; blood in our body is a good example of a basic buffer. The \({\rm{pH}}\) of blood in the human body is \(7.4.\) Both higher and lower \({\rm{pH}}\) values of blood may be fatal. Human blood contains a buffer of carbonic acid \(\left(\mathrm{H}_{2} \mathrm{CO}_{3}\right)\) and bicarbonate anion \(\left( {{\rm{HCO}}_{\rm{3}}^{\rm{ – }}} \right)\) to maintain blood \({\rm{pH}}\) In this buffer system, hydronium and bicarbonate anion are in equilibrium with carbonic acid.

FAQs on Buffer Solutions

Q.1. Among \(\mathrm{CH}_{3} \mathrm{COOH}+\mathrm{CH}_{3} \mathrm{COONa}\) and \(\mathrm{HCl}+\mathrm{NaCl}\) Which of the following is a buffer solution?
Ans: Among \(\mathrm{CH}_{3} \mathrm{COOH}+\mathrm{CH}_{3} \mathrm{COONa}\) and \(\mathrm{HCl}+\mathrm{NaCl}\) the mixture of \(\mathrm{CH}_{3} \mathrm{COOH}+\mathrm{CH}_{3} \mathrm{COONa}\) is a buffer solution because a buffer solution must contain a weak acid or a weak base along with its conjugate salt. Thus, acetic acid and sodium acetate are a weak acid and its conjugate salt respectively. On the other hand, \(\mathrm{HCl}\) is a strong acid, so it cannot be a buffer solution.

Q.2. How to prepare a buffer solution?
Ans: A buffer solution is prepared by mixing a weak acid like acetic acid or a weak base with the conjugate salt of that weak acid or weak base. A conjugate salt should differ from its original acid or base by one \({{\rm{H}}^{\rm{ + }}}\) ion or by one \(\mathrm{OH}^{-}\) ion.

Q.3. What is meant by buffer solution?
Ans: A buffer solution means the solution that can resist any change in \({\rm{pH}}\) upon the addition of acidic or alkaline components in a solution. It maintains a stable \({\rm{pH}}\) range of a particular reaction or process because a buffer solution can neutralize the solution when a small amount of acid or base is added to it during a particular chemical reaction. 

Q.4. How to calculate the pH of the buffer solution?
Ans: The \({\rm{pH}}\) of a buffer solution is calculated by using Henderson-Hasselbalch Equation with a known amount of acid and its conjugate base as \({\rm{pH = p}}{{\rm{K}}_{\rm{a}}}{\rm{ + log}}\left( {\frac{{\left[ {{\rm{Salt}}} \right]}}{{\left[ {{\rm{Acid}}} \right]}}} \right)\)

Q.5. Among acidic buffer and basic buffer, which buffer solution has maximum pH?
Ans: A basic buffer solution has maximum \({\rm{pH}}\) among acidic buffer and basic buffer because acidic buffer solutions have pH below \(7\) and basic buffer solutions have \({\rm{pH}}\) above \(7.\)

Q.6. How to prepare a buffer solution in the laboratory?
Ans: In the laboratory, buffer solutions can be prepared by using a conjugate acid-base pair with a \({\rm{p}}{{\rm{K}}_{\rm{a}}}\) value close to the desired \({\rm{pH}}\) and then adding a strong acid or base to get the exact \({\rm{pH}}\) of the solution that is required.

Q.7.How to determine if two solutions will make a buffer?
Ans: To determine if the two solutions will make a buffer or not, it should fulfill the following two criteria: Firstly, a buffer should compose of a weak acid or a weak base, for example, acetic acid, nitrous acid, ammonia, etc. Secondly, it should be a conjugate salt of that weak acid or weak base. A conjugate salt always differs from its original acid or base by one \({{\rm{H}}^{\rm{ + }}}\) ion or by one \(\mathrm{OH}^{-}\) ion.