Buffer Solutions-Acidic and Basic Buffers, Preparation, And Applications

What are Buffer Solutions?

A buffer solution is a solution that can resist changes in pH when small amounts of acid or base are added to it. This is because buffer solutions contain a mixture of a weak acid and its conjugate base, or a weak base and its conjugate acid. Some examples of buffer solutions are

• Acetic acid and sodium acetate
• Citric acid and sodium citrate
• Phosphoric acid and sodium phosphate
• Boric acid and sodium borate

Buffer solutions can be prepared by mixing a weak acid with its conjugate base, or a weak base with its conjugate acid. The ratio of the weak acid to its conjugate base, or the weak base to its conjugate acid, determines the pH of the buffer solution.

Types of Buffer Solutions

There are two main types of buffer solutions: acidic buffers and basic buffers.

Acidic buffers

Acidic buffers are solutions that have a pH below 7 and contain a weak acid and one of its salts. For example, a mixture of acetic acid and sodium acetate acts as a buffer solution with a pH of about 4.75.

When a small amount of strong acid is added to an acidic buffer solution, the conjugate base of the weak acid reacts with the strong acid to form the weak acid and a salt. This reaction consumes the strong acid and prevents it from lowering the pH of the solution.

Basic buffers

Basic buffers are solutions that have a pH above 7 and contain a weak base and one of its salts. For example, a mixture of ammonium chloride and ammonium hydroxide acts as a buffer solution with a pH of about 9.25.

When a small amount of strong base is added to a basic buffer solution, the conjugate acid of the weak base reacts with the strong base to form the weak base and a salt. This reaction consumes the strong base and prevents it from raising the pH of the solution.

Properties of Buffer Solutions

Buffer solutions have the following properties:

• They have a definite pH. The pH of a buffer solution is determined by the ratio of the weak acid to its conjugate base or the weak base to its conjugate acid.
• Their pH does not change easily when small amounts of acid or base are added. 
• Their pH does not change significantly when diluted. This is because the ratio of the weak acid to its conjugate base, or the weak base to its conjugate acid, remains the same even when the solution is diluted.
• They have a limited capacity to buffer acids and bases. The amount of acid or base that a buffer solution can absorb before its pH changes significantly is called its buffer capacity. The buffer capacity of a buffer solution is determined by the concentration of the weak acid and its conjugate base, or the weak base and its conjugate acid.
• Ita pH does not change on keeping for a long time.

Preparation of Buffer Solutions

Buffer solutions are created to maintain a stable pH. This is accomplished by combining a weak acid with its corresponding conjugate base, or a weak base with its corresponding conjugate acid.

The selection of the acid and base is crucial because it depends on their known dissociation constants, specifically the pKa for the weak acid and the pKb for the weak base.

For instance, consider the preparation of a phosphate buffer solution, which maintains a pH of 7.4. In this case, you’re mixing two components: HPO₄^2- (the conjugate base) and H₂PO₄^– (the weak acid). The pKa values for these components are known, allowing you to control the ratio of salt (conjugate base) to acid (weak acid) to achieve the desired pH.

Why do we need buffer solutions?

We need buffer solutions for a variety of reasons, including:

• To maintain the pH of biological fluids. The pH of many biological fluids, such as blood and urine, must be maintained within a narrow range for optimal function. Buffer solutions help to keep the pH of these fluids stable, even when there are changes in the concentration of acids or bases.
• To control the pH of food and beverages. The pH of food and beverages can affect their taste, texture, and shelf life. Buffer solutions are used to control the pH of food and beverages to ensure that they are safe and palatable.
• To stabilize the pH of industrial processes. Many industrial processes require a specific pH in order to proceed efficiently and safely. Buffer solutions are used to maintain the pH of these processes within a narrow range.
• To calibrate pH meters. pH meters are used to measure the pH of solutions. Buffer solutions are used to calibrate pH meters so that they provide accurate readings.

Uses of Buffer Solutions

Here are some specific examples of how buffer solutions are used:

• In the human body: Buffer solutions help to maintain the pH of blood, urine, and other bodily fluids. This is important for the proper functioning of enzymes and other proteins in the body.
• In the food industry: Buffer solutions are used to control the pH of food and beverages. This is important for taste, texture, and shelf life. For example, buffer solutions are used to make soft drinks and wine.
• In the pharmaceutical industry: Buffer solutions are used to control the pH of pharmaceutical products. This is important for stability and efficacy.
• In the environmental industry: Buffer solutions are used to neutralize acidic or basic wastewater. This is important to protect the environment and human health.

Buffer solutions are a valuable tool for controlling and maintaining pH in a wide variety of applications.

Limitations of Buffer Solutions

1. Limited pH Range: Buffer solutions can only effectively control pH within a certain range. Once the pH of a solution moves too far outside this range, the buffering capacity diminishes. This range depends on the specific buffer system in use.
2. Capacity: Buffer solutions have a finite capacity to resist changes in pH. If the amount of acid or base added to the solution exceeds the buffer’s capacity, the pH will change significantly.
3. Choice of Buffer: Selecting the right buffer for a particular application is crucial. Not all buffers are suitable for every situation, and the wrong choice can result in ineffective pH control.
4. Temperature Sensitivity: The pH of a buffer solution can be sensitive to changes in temperature. Some buffer systems are more temperature-stable than others, and adjustments may be needed if the temperature changes significantly.
5. Chemical Compatibility: Buffers may not be compatible with certain chemical reactions or processes. They can interfere with specific reactions or react with other components in the system, leading to unintended consequences.
6. Dilution Effect: As a buffer solution is diluted, its buffering capacity decreases. This is important to consider when making serial dilutions in laboratory experiments.
7. Maintenance: Buffers need to be maintained and monitored regularly. Over time, the buffer capacity may decrease due to contamination or chemical reactions, and adjustments may be necessary.
8. Limited Use in Extreme Conditions: Buffer solutions may not be suitable for very high or very low pH environments. Extreme conditions can overwhelm the buffering capacity of most standard buffers.
9. Precise Preparation: Preparing a buffer solution with the exact desired pH can be challenging. Small errors in the preparation process can lead to pH variations.
10. Shelf Life: Buffer solutions can degrade or become contaminated over time, reducing their effectiveness. Proper storage and regular replacement are important to maintain their reliability.

Buffer Solutions Solved Problems

You have a buffered solution containing 0.3 M acetic acid (CH₃COOH) and 0.6 M acetate ions (CH₃COO^–) in a total volume of 500 mL. You then add 0.2 moles of hydrochloric acid (HCl) to the solution. What is the final pH of the buffer solution?

Solution:

First, calculate the initial pH of the buffer solution using the Henderson-Hasselbalch equation before adding HCl: The pKa for acetic acid is 4.76

pH=pKa+log([CH₃COO^–][CH3COOH]​)=4.76+log(0.30.6​)=4.76+0.3010=5.0610

So, the initial pH of the buffer solution is approximately 5.06.

Now, you add 0.2 moles of HCl to the buffer solution. This will react with the acetate ions (CH3COO^–) to form acetic acid (CH₃COOH).

The moles of acetate ions before adding HCl is 0.6 M x 0.5 L = 0.3 moles.

The moles of acetate ions after the reaction will be 0.3 moles – 0.2 moles = 0.1 moles.

The moles of acetic acid (CH₃COOH) after the reaction will be 0.3 moles (initial) + 0.2 moles (formed) = 0.5 moles.

Now, calculate the new pH after the reaction using the Henderson-Hasselbalch equation:

pH=pKa+log([CH₃COO^–][CH₃COOH]​)=4.76+log(0.50.1​)=4.76+0.6020=5.3620

So, the final pH of the buffer solution after adding HCl is approximately 5.36.

This example demonstrates how a buffer solution can resist changes in pH even when an acid is added, thanks to the ability of the buffer to neutralize the added acid and maintain a relatively stable pH.