Buffers: Definition, Applications, & Industry Uses

Buffers: Definition, Applications, & Industry Uses

Buffers: Overview

Buffers are aqueous solutions that can resist pH changes when acidic or basic components are added to a given solution. It is essential to maintain the pH (or hydrogen ion concentration) of the solution to carry out reactions or processes that demand specific pH ranges — pH is the negative logarithm of the hydrogen ion concentration.

By definition, buffers are a mixture of a weak base and a conjugate acid (or a solution of a weak acid and a conjugate base — a conjugate acid-base pair). The type of acid/base mixture used in the preparation of a buffer solution depends on the desired pH value of the buffer in the workflow.

The amount of acid/base that can be neutralized by a buffer solution is determined by its working pH range and capacity. Some examples of buffer solutions used in lab processes include:

  • Acetic acid (CH3COOH — weak acid) and its conjugate base, the acetate anion (CH3COO-), e.g. sodium acetate (CH3COONa).
  • Pyridine (C5H5N — weak base) and its conjugate acid, the pyridinium cation (C5H5NH+), such as Pyridinium Chloride.

There are also many natural buffering systems for pH regulation. A good example is the bicarbonate buffering system that regulates the pH of the blood and the ocean.

However, it must be noted that adding too much of a strong base or acid to the buffer solution can alter the pH to an unwanted degree. The number of moles of acid or base added to the 1-liter buffer solution to change the pH of a solution by 1 unit is known as buffer capacity.

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Types of Buffer and How Do They Work?

Buffers are of two main types:

Acidic Buffer

An acidic buffer has a pH of less than 7. It contains a weak acid and its conjugate base. The pH of the buffer solution can be changed by altering the ratio of acid and salt used. An example of an acidic buffer is acetic acid and sodium acetate.

CH3COOH(aq) ⇄ CH3COO–(aq) + H+(aq)

In the above reaction, the equilibrium of the reaction is going left, according to Le Chatelier’s principle.

Adding Acid to Acidic Buffer

In such a condition, the acetate ion will combine with the hydrogen ion of the added acid (such as HCl), creating acetic acid. As acetic acid is a weak acid, hydrogen ions can be removed faster than the disassociation of the acetic acid, which prevents any alteration in the pH of the solution.

Adding Base to Acidic Buffer

While adding base to the buffer, the acetic acid will interact with the hydroxide ion of the base, forming an acetate ion with a water molecule. This is how hydroxide ions will be removed from the buffer solution, maintaining its pH.

Basic Buffer

A basic buffer solution has a pH greater than 7. It contains a weak base with its conjugate acid as a salt. For instance, ammonia and ammonium chloride are examples of basic buffers.

NH3 + H2O →NH4+ + OH-

According to Le Chatelier’s principle, the equilibria of the above reaction are heading left.

Adding Acid to Basic Buffer

The added acid will release hydrogen ions in the buffer solution that will react with ammonia, forming ammonium ions. The hydroxide ion in the solution (formed as a result of the reaction between a water molecule and ammonia) reacts with hydrogen ions to form a water molecule. In this way, the pH of the buffer solution remains unaltered.

Adding Base to Basic Buffer

When bases are added to the buffer solution, the hydroxide ion will react with ammonium ions in the solution. Thus, maintaining their pH.

How Does Buffer Work?

The presence of conjugate acid and conjugate base in the buffer resists any changes in the pH of the solution. They neutralize the small amount of acid or base added in the form of protons or hydronium ions (H3O+) or hydroxide ions (OH).

Let’s understand this through the following example:

Consider a buffer made of ammonia (NH3) and its conjugate acid (NH4+). Now, when a strong acid (such as HCl) is added, NH3 reacts with the H+ ion forming a weak base NH4+, thus maintaining the pH of the solution. Similarly, when a strong base (such as NaOH) is added, the ammonium donates its one proton to the base, forming ammonia and water. This keeps the pH of the solution unaltered.

Figure: A titration curve representing the buffering action of a buffer solution.

How to Calculate The pH of A Buffer Solution

The pH of the buffer solution can be calculated using Henderson-Hasselbalch Equation.

  • For acidic buffer solution (containing a weak acid (HA) and conjugate base (KA))

Ka = [ (H+)(A-)] / [(H2O)]

pKa= pH – log[( salt)/(acid)]

pH= pK a + log[(salt)/(acid)]

the pH of acid buffer solution= pKa+[(salt)/(acid)]

Ka = Dissociation constant

  • For basic buffer solution (containing a weak base and conjugate acid)

pOH=pKb + log[(salt)/(acid)]

pOH of basic buffer solution=pKb+ log[(salt)/(acid)]

pOH = pKa – log[(salt)/(acid)]

What are Buffers Used For?

Buffers are aqueous solutions made of a weak acid and its conjugate base or weak base and its conjugate acid. Some of its examples include:

  • Bicarbonate buffer
  • Tris buffer
  • Phosphate buffer

These buffering systems have extensive applications in maintaining the pH of a system.

Maintaining pH

The chemistry of a biological system greatly depends on its pH as well. It has many roles ranging from impacting the electrical conductivity of the solution, and solubility of compounds, to the structure and activity of biological macromolecules. Thus, it prevents havoc caused by pH changes in the system.

For example, a protein-digesting enzyme in human intestine works with 2X efficiency at a pH of 8 versus a pH of 7 or 9. In the brewing industry, it’s used to prevent the product from becoming too acidic during the fermentation process.

What Industries Use Buffers?

Buffers have a range of applications in the Life science biochemical, and manufacturing industries to maintain the pH of the systems. For example, it has applications in food preservatives, electroplating, fermentation, drug delivery, printing, and in the activity of enzymes that require specific hydrogen ion concentrations.

Life Sciences

In biological systems, whether naturally or in labs, buffers are required to maintain enzymatic activities. For instance, a mixture of carbonic acid (H2CO3) and bicarbonate (HCO−3) in plasma maintains the pH in human blood between 7.35 and 7.45.

When pH levels are outside this narrow range (7.40 + 0.05 pH unit), metabolic conditions, such as acidosis or alkalosis develop rapidly, leading to death due to the inability to restore buffering capacity.

In labs, buffers have extensive applications in performing enzymatic reactions or cellular studies to maintain their pH and activity during the reaction.  Further, they also have applications in calibrating pH meters and chemical analysis.


Buffering agent is used to set the right conditions for dyes having application in coloring fabrics and fermentation processes. Additionally, buffers are used in manufacturing cosmetics to maintain their pH and prevent any reactions to the product, such as rashes.

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Buffers are aqueous solutions used to maintain the pH of a system. They consist of a weak acid and its conjugated base or a weak base and its conjugated acid. They help neutralize the effects of small amounts of acid or base when added to a solution to stabilize its pH.

It’s of two types: acidic buffer and basic buffer. A few examples of buffers include phosphate buffer, citrate buffer, bicarbonate buffer, and Tris Buffer. The preparation of such a solution requires one to use high-quality chemicals for effective experimental results.

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