Buffer Solutions - Allied Guru

Table of contents

Introduction

Buffer solutions are special solutions that resist changes in pH when small amounts of acid or base are added.
They are composed of a weak acid and its conjugate base, or a weak base and its conjugate acid.
The main function of a buffer is to maintain a constant hydrogen ion concentration (pH) in a solution, even when external factors tend to change it.
Buffer solutions play a crucial role in biological and chemical systems, where enzymes and reactions require a specific pH to function optimally.
Common examples include the phosphate buffer, acetate buffer, and bicarbonate buffer, which help maintain pH stability in blood, cells, and laboratory experiments.
The capacity of a buffer depends on the concentration of its components and the ratio between acid and base forms.

Types of Buffer Solutions

Acidic Buffer Solution:
- Maintains an acidic pH (below 7).
- Formed by mixing a weak acid with the salt of that acid and a strong base.
- Example: Acetic acid (CH₃COOH) and sodium acetate (CH₃COONa) maintain a pH of around 4.75.
Basic Buffer Solution:
- Maintains a basic or alkaline pH (above 7).
- Prepared by combining a weak base with its salt formed with a strong acid.
- Example: Ammonium hydroxide (NH₄OH) and ammonium chloride (NH₄Cl) — maintains pH around 9.25.
Neutral Buffer Solution:
- Maintains a neutral pH (around 7).
- Usually made from neutral salts of weak acids and weak bases.
- Example: A potassium dihydrogen phosphate (KH₂PO₄) and disodium hydrogen phosphate (Na₂HPO₄) mixture maintains pH near 7.0.
Biological Buffer Solution:
- Found naturally in living systems to maintain constant physiological pH.
- Example: Bicarbonate buffer system (H₂CO₃/HCO₃⁻) in blood maintains pH around 7.4.
Mixed Buffer System:
- Contains more than one buffer pair and can resist pH changes over a wider range.
- Example: Phosphate buffer (NaH₂PO₄ and Na₂HPO₄) used in laboratory and biological experiments.

Mechanism

Buffer mechanism works on the principle of the common ion effect — the presence of a common ion suppresses the ionization of a weak acid or base, helping the solution resist pH changes.
When a small amount of acid (H⁺) is added to a buffer, the conjugate base present in the buffer combines with these hydrogen ions to form the weak acid, preventing a large increase in H⁺ concentration.

Example: In acetic acid–sodium acetate buffer:
H⁺ + CH₃COO⁻ → CH₃COOH
When a small amount of base (OH⁻) is added, the weak acid in the buffer reacts with it to form water and its conjugate base, minimizing any rise in pH.

Example: CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O
In a basic buffer like NH₄OH/NH₄Cl, the mechanism is similar:
- Added acid (H⁺) reacts with NH₄OH → NH₄⁺ + H₂O
- Added base (OH⁻) reacts with NH₄⁺ → NH₃ + H₂O
In this way, a buffer maintains nearly constant pH by neutralizing small additions of acid or base, keeping the ratio of weak acid to conjugate base nearly constant.

Preparation of Buffer Solutions

Acidic Buffer Solution (Example: Acetic acid–Sodium acetate buffer)
- Take 50 mL of 0.1 M acetic acid (CH₃COOH) and 50 mL of 0.1 M sodium acetate (CH₃COONa).
- Mix them thoroughly in a volumetric flask.
- This combination gives a buffer of pH ≈ 4.75.
- The pH can be calculated using the Henderson–Hasselbalch equation:
  
  For acetic acid, pKa = 4.75, so when [Salt] = [Acid], pH = 4.75.
Basic Buffer Solution (Example: Ammonium hydroxide–Ammonium chloride buffer)
- Take 25 mL of 0.1 M ammonium hydroxide (NH₄OH) and 25 mL of 0.1 M ammonium chloride (NH₄Cl).
- Mix both solutions in a clean flask.
- This gives a buffer of pH ≈ 9.25.
- The pH can be calculated by the relation:
  
  For NH₄OH, pKb = 4.75 → pH = 14 – 9.25 = 9.25.
Neutral Buffer Solution (Example: Phosphate buffer)
- Mix 50 mL of 0.1 M NaH₂PO₄ (acidic component) with 50 mL of 0.1 M Na₂HPO₄ (basic component).
- The resulting buffer maintains a pH around 7.0.
- This buffer is widely used in biological and biochemical experiments because it mimics physiological pH.
Biological Buffer (Example: Bicarbonate buffer system)
- The combination of carbonic acid (H₂CO₃) and sodium bicarbonate (NaHCO₃) forms a buffer that maintains blood pH around 7.4.
- Reaction:
  H₂CO₃ ⇌ H⁺ + HCO₃⁻
- This buffer system helps regulate pH in the human body and other biological fluids.

Steps for Preparation

Identify the Required pH Range:
Decide whether an acidic or basic buffer is needed based on the desired pH of the solution.
- For pH < 7 → Acidic buffer
- For pH > 7 → Basic buffer
Select Appropriate Components:
- For an acidic buffer, choose a weak acid and its salt with a strong base (e.g., CH₃COOH and CH₃COONa).
- For a basic buffer, choose a weak base and its salt with a strong acid (e.g., NH₄OH and NH₄Cl).
Calculate Required Molar Ratios:
Use the Henderson–Hasselbalch equation to determine the correct ratio of acid to salt (or base to salt) for the desired pH:

or
Prepare Stock Solutions:
Prepare separate standard (0.1 M or 1 M) solutions of the weak acid/base and its salt using distilled water.
Mix in Correct Proportion:
Measure the calculated volumes of both components using a pipette or burette and mix them in a clean volumetric flask.
Check and Adjust pH (if necessary):
Use a pH meter to measure the solution’s pH. Adjust slightly by adding small amounts of acid or base until the desired pH is achieved.
Make Up to Final Volume:
Dilute the mixture to the required final volume with distilled water and mix thoroughly to ensure uniformity.
Label and Store:
Transfer the prepared buffer into a clean, airtight bottle, label it with composition and pH, and store it properly to avoid contamination or CO₂ absorption.

Example:

Preparation of 0.1 M Phosphate Buffer (pH 7.4):

Select the pKa:
- Phosphoric acid has multiple pKa values, but for this buffer, the relevant pKa is 7.2, close to the desired pH of 7.4.
Henderson-Hasselbalch Equation:

Ph = pKa + log ([Base]/[Acid])

Since the desired pH (7.4) is slightly higher than the pKa (7.2), the base (Na₂HPO₄) concentration will be slightly higher than the acid (NaH₂PO₄).

Calculation of Molar Ratio:
- The Henderson-Hasselbalch equation gives:

7.4 = 7.2 + log ([Base]/[Acid])

Solving this gives a base-to-acid ratio of approximately 1.6:1.

Mixing:
- Prepare 0.1 M solutions of Na₂HPO₄ and NaH₂PO₄.
- Mix 61.5 mL of 0.1 M Na₂HPO₄ with 38.5 mL of 0.1 M NaH₂PO₄.
Adjust Volume:
- Add distilled water to bring the total volume to 100 ml.
Check pH:
- Use a pH meter to check the buffer’s pH. Adjust slightly with small amounts of NaOH or HCl if necessary to keep the pH to 7.4.

Applications

In Biological Systems:
Buffer solutions help maintain a constant pH in body fluids like blood and intracellular fluids, which is vital for enzyme activity and metabolic processes.
Example: The bicarbonate buffer system maintains blood pH around 7.4.
In Pharmaceutical Preparations:
Buffers are used in medicines and injections to maintain the required pH for stability and proper absorption of drugs.
Example: Phosphate buffers are used in ophthalmic and intravenous preparations.
In Chemical and Biochemical Laboratories:
Buffers are essential in titrations, electrophoresis, chromatography, and enzyme assays where precise pH control is required.
Example: Tris buffer and phosphate buffer are used in enzyme kinetics experiments.
In Industrial Processes:
Used in fermentation, dyeing, tanning, and electroplating industries where controlled pH is necessary for optimal reactions.
In Food Industry:
Buffers maintain the taste, texture, and shelf life of foods by preventing unwanted pH changes.
Example: Citric acid and sodium citrate buffer used in soft drinks and dairy products.
In Agriculture:
Buffer solutions are used to calibrate pH meters and to maintain soil pH for plant growth studies.
In Analytical Chemistry:
Buffers provide a stable pH environment for accurate chemical analysis, especially in colorimetric and spectrophotometric assays.
In Molecular Biology and Biotechnology:
Commonly used in DNA extraction, PCR reactions, and protein purification to maintain enzyme activity and structural integrity of biomolecules.

MCQs

A buffer solution resists changes in:
A. Temperature
B. pH
C. Concentration
D. Pressure
A buffer is generally made of:
A. Strong acid and strong base
B. Weak acid and its salt
C. Strong acid and its salt
D. Neutral compounds
Which of the following is an acidic buffer?
A. NH₄OH + NH₄Cl
B. CH₃COOH + CH₃COONa
C. Na₂HPO₄ + NaH₂PO₄
D. HCl + NaCl
Which of the following is a basic buffer?
A. NH₄OH + NH₄Cl
B. H₂CO₃ + NaHCO₃
C. CH₃COOH + CH₃COONa
D. H₃PO₄ + NaH₂PO₄
The pH of an acidic buffer is:
A. Greater than 7
B. Equal to 7
C. Less than 7
D. Cannot be determined
The Henderson–Hasselbalch equation is used to:
A. Calculate pH of a buffer
B. Calculate temperature of a reaction
C. Measure molar concentration
D. Determine ionic strength
The Henderson–Hasselbalch equation for an acidic buffer is:
A. pH = pKa + log [acid]/[salt]
B. pH = pKa + log [salt]/[acid]
C. pH = pKb + log [base]/[salt]
D. pOH = pKa + log [salt]/[acid]
The buffer capacity depends on:
A. Volume of buffer
B. Ratio of acid to base
C. Total concentration of buffer components
D. Both B and C
The pKa of acetic acid is 4.75. When [acid] = [salt], the pH of buffer is:
A. 2.5
B. 4.75
C. 7.0
D. 9.0
Which of the following maintains blood pH?
A. Phosphate buffer
B. Acetate buffer
C. Bicarbonate buffer
D. Ammonium buffer
The bicarbonate buffer system maintains pH around:
A. 6.8
B. 7.0
C. 7.4
D. 8.2
In a buffer, if acid concentration is doubled, pH will:
A. Increase
B. Decrease
C. Remain same
D. Be zero
Which of the following pairs is not a buffer system?
A. H₂CO₃ / NaHCO₃
B. NH₄OH / NH₄Cl
C. HCl / NaCl
D. CH₃COOH / CH₃COONa
What happens when a small amount of acid is added to a buffer?
A. pH increases sharply
B. pH decreases sharply
C. pH remains nearly constant
D. Solution becomes neutral
In a basic buffer, the weak base reacts with:
A. Added acid (H⁺)
B. Added base (OH⁻)
C. Salt ions
D. None of these
In an acidic buffer, the weak acid reacts with:
A. Added OH⁻ ions
B. Added H⁺ ions
C. Both
D. Neither
Which of the following is a neutral buffer?
A. NaH₂PO₄ + Na₂HPO₄
B. HCl + NaCl
C. NH₄OH + NH₄Cl
D. CH₃COOH + CH₃COONa
Which of the following maintains the pH of intracellular fluids?
A. Bicarbonate buffer
B. Phosphate buffer
C. Acetate buffer
D. Tris buffer
Buffers are important in biochemical reactions because:
A. They speed up reactions
B. They maintain optimal pH for enzyme activity
C. They reduce ionic strength
D. They act as catalysts
A buffer solution has maximum capacity when:
A. pH = pKa
B. pH = 7
C. pH > pKa
D. pH < pKa
Which of the following combinations gives a buffer of pH ≈ 4.75?
A. CH₃COOH + CH₃COONa
B. NH₄OH + NH₄Cl
C. NaH₂PO₄ + Na₂HPO₄
D. H₂CO₃ + NaHCO₃
Buffer solutions are used in:
A. Enzyme assays
B. Blood pH control
C. Pharmaceutical formulations
D. All of these
The buffering action fails when:
A. Large amounts of acid/base are added
B. pKa changes
C. Concentration becomes zero
D. All of these
Which of the following buffers is commonly used in biological experiments?
A. Tris buffer
B. Borate buffer
C. Glycine buffer
D. All of these
pH of a buffer solution mainly depends on:
A. Temperature
B. Ratio of salt to acid/base
C. Total volume
D. Color of solution
The common ion effect plays a key role in:
A. Buffer mechanism
B. Redox reaction
C. Electrolysis
D. Combustion
Which buffer system is used in the human kidney for acid-base balance?
A. Ammonium buffer
B. Phosphate buffer
C. Bicarbonate buffer
D. Both A and B
The term “buffer capacity” refers to:
A. Ability to change color
B. Ability to resist pH change
C. Amount of acid/base neutralized
D. Both B and C
When equal moles of acid and salt are mixed, the buffer shows:
A. Maximum buffering capacity
B. Minimum buffering capacity
C. No buffering action
D. Unstable pH
Buffers are used in calibration of:
A. Colorimeter
B. pH meter
C. Spectrophotometer
D. Centrifuge

Answer Key