The solutions which resist change in pH on dilution or with the addition of small amount of acid or alkali are called buffer solutions. It is a unique kind of solution that doesn’t react considerably to minor additions of basic or acid in terms of pH. Because of this characteristic, buffer solutions are vital for both industrial and natural processes.
What is a Buffer Solution?
A buffer solution is made by combining:
A weak acid and its conjugate base (acidic buffer).
Example: Acetic acid (CH₃COOH) and sodium acetate (CH₃COONa).
A weak base and its conjugate acid (basic buffer).
Example: Ammonium hydroxide (NH₄OH) and ammonium chloride (NH₄Cl).
The main characteristic of a buffer is its capacity to keep the pH constant even when substances from outside sources that could change the pH are added. This occurs as a result of reversible reactions between the buffer solution’s constituents that neutralise additional acids or bases.
How Do Buffer Solutions Work?
Le Chatelier’s Principle, which disagree that a system at equilibrium adapts to counteract changes in conditions, is the basis for buffer solutions.
In an acidic buffer:
If a strong acid (like HCl) is added, the weak base in the buffer neutralises the added hydrogen ions (H⁺).
If a strong base (like NaOH) is added, the weak acid in the buffer neutralises the added hydroxide ions (OH⁻).
Example:
Think about a buffer made of acetic acid (CH₃COOH) and sodium acetate (CH₃COONa):
When H⁺ ions are added, acetate ions (CH₃COO⁻) combine with them to form acetic acid.
When OH⁻ ions are added, acetic acid reacts with them to form water and acetate ions.
In a basic buffer:
If a strong acid is added, the weak base neutralises the H⁺
If a strong base is added, the weak acid neutralises the OH⁻ ions
Example:
Think a buffer made of ammonium hydroxide (NH₄OH) and ammonium chloride (NH₄Cl):
When H⁺ ions are added, they react with NH₄OH to form NH₄⁺.
When OH⁻ ions are added, NH₄⁺ reacts to produce NH₄OH and water.
Types of Buffer Solutions
Acidic Buffers:
pH < 7
Made from a weak acid and its salt with a strong base.
Example: CH₃COOH and CH₃
Basic Buffers:
pH > 7
Made from a weak base and its salt with a strong acid.
Example: NH₄OH and NH₄
Importance of Buffer Solutions
Biological Systems:
In our body, blood acts as a buffer to maintain a pH around 7.4, which is vital for proper enzyme function and metabolic processes.
Buffering systems like the bicarbonate buffer (H₂CO₃/HCO₃⁻) stabilise the blood’s pH.
2.Industrial Applications:
Used in the fermentation process, where a stable pH is vital.
Essential in the manufacture of pharmaceuticals and dyes.
3. Analytical Chemistry:
Buffers are used to calibrate pH meters and maintain a stable pH in chemical reactions.
4. Agriculture:
Soil buffers help maintain the pH required for the best possible plant growth.
pH of a Buffer Solution
The pH of a buffer solution can be calculated using the Henderson-Hasselbalch equation:
For acidic buffers:
pH = pKa + log([Salt] / [Acid])
For basic buffers:
pOH = pKb + log([Salt] / [Base])
Key Features of Buffers
Effective buffering range: Buffers work the best within ±1 pH unit of the pKa or pKb of the weak acid or base.
Buffer capacity: The ability of a buffer to resist pH changes depends on the concentration of the acid / base pair. Higher concentrations offer better buffering.
Note :-
Maintaining the delicate pH equilibrium in a variety of systems requires buffer solutions. Buffers guarantee stability and uniformity in everything from the operation of human cells to the effectiveness of industrial processes.
Buffer capacity = No. of moles of acid or base added per litre of buffer / Change in pH
The range of pH over which the buffer solutions remain effective is called buffer range.
Buffer range in pH
Acidic pKb + – 1
Basic (pKa – pKb) + – 1
A buffer solution is one that, even when modest amounts of an acid or a base are added, does not notably alter its pH. Usually, it is made up of a weak base and its conjugate acid or a weak acid and its conjugate base.
There are two main types of buffer solutions:
Acidic Buffers: Made from a weak acid and its salt with a strong base (e.g., acetic acid and sodium acetate). These have a pH less than 7.
Basic Buffers: Made from a weak base and its salt with a strong acid (e.g., ammonium hydroxide and ammonium chloride). These have a pH greater
than 7.
Buffer solutions resist pH changes through neutralisation:
In acidic buffers, the weak acid neutralises added bases, and the conjugate base neutralises added acids.
In basic buffers, the weak base neutralises added acids, and the conjugate acid neutralises added bases.
The Henderson-Hasselbalch equation helps calculate the pH of a buffer solution:
For acidic buffers: pH = pKa + log([Salt] / [Acid])
For basic buffers: pOH = pKb + log([Salt] / [Base])
Buffer solutions are vital for maintaining the pH required for biological processes. For example:
The bicarbonate buffer system in blood maintains a pH of around 7.4, critical for enzyme activity and metabolic functions.
Cellular buffers stabilise the pH inside cells, ensuring proper biochemical reactions.
The effectiveness of a buffer solution depends on:
Concentration of buffer components: Higher concentrations offer better buffering capacity.
pKa or pKb of the buffer: Buffers work best when the pH is within ±1 unit of their pKa or pKb.
Ratio of acid/base to salt: The pH is influenced by this ratio, as seen in the Henderson-Hasselbalch equation.