Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → General chemistry → Acids and bases ↓


Buffer Solution


A buffer solution is a type of chemical solution that resists significant changes in its pH level when a small amount of acid or base is added. This feature makes buffer solutions highly essential in various chemical, biological, and industrial applications. In the field of chemistry, understanding buffer solutions is important for students, especially those pursuing graduate studies. The aim of this lesson is to understand the intricacies of buffer solutions, including their structure, functioning, types, and applications in simple English terms.

Composition of buffer solutions

Buffer solutions usually contain a weak acid and its conjugate base or a weak base and its conjugate acid. The presence of these components makes the solution able to resist large changes in pH.

For example, a common buffer solution is acetic acid-sodium acetate buffer. Here, acetic acid (CH 3 COOH) is the weak acid, while sodium acetate (CH 3 COONa) provides the conjugate base (acetate ion, CH 3 COO -).

      CH3 COOH (aq) ⇌ H + (aq) + CH3 COO - (aq)

Working of buffer solution

The primary function of a buffer solution is to maintain the pH level of the solution near a selected value. This occurs basically through two mechanisms:

1. Addition of acid

When a small amount of acid (e.g., HCl) is added to a buffer solution, the hydrogen ions (H +) from the acid react with the conjugate base present in the solution, forming a weaker acid. This dampens any increase in H + ion concentration, thus counteracting changes in pH.

      H + (aq) + CH 3 COO - (aq) → CH 3 COOH (aq)

2. Sum of bases

When a small amount of a base (e.g., NaOH) is added, it removes H + ions from the solution by forming water, resulting in a greater development of conjugate base character. This again minimizes the change in pH by maintaining a balance between the acid and base forms.

      OH - (aq) + CH 3 COOH (aq) → CH 3 COO - (aq) + H 2 O (l)

Visual representation of the buffer system

buffer solution weak acid conjugate base

Henderson–Hasselbalch equation

The pH value of a buffer solution can be estimated using the Henderson-Hasselbalch equation, which is a fundamental equation in chemistry.

      pH = pK A + log([A - ]/[HA])

In this equation:

  • pH is the potential of hydrogen, which measures the acidity or alkalinity of a solution.
  • pK a is the acid dissociation constant, a measure of the strength of a weak acid in solution.
  • [A - ] is the concentration of the conjugate base.
  • [HA] is the concentration of the weak acid.

Types of buffer solutions

1. Acidic buffer

These are solutions that have a pH less than 7. They are composed of a weak acid and its salt and a strong base. An example of an acidic buffer is a solution containing acetic acid and sodium acetate.

2. Basic buffers

Basic buffers have a pH greater than 7. They are composed of a weak base and its salt and a strong acid. A common example of a basic buffer is a solution of ammonium hydroxide and ammonium chloride.

Applications of buffer solutions

Buffer solutions are applied in many areas due to their pH stabilizing properties:

1. Biological systems

In biological systems, buffers maintain the optimal pH necessary for the proper functioning of enzymes and metabolic processes. For example, human blood contains a bicarbonate buffer system that helps maintain a pH of about 7.4.

2. Chemical reactions

Buffer solutions are used in chemical experiments where maintaining a constant pH is necessary for a reaction to proceed correctly. This is especially important in analytical chemistry procedures such as titrations.

3. Industrial processes

In industries, buffer solutions are used in a variety of processes such as fermentation, where pH control is critical for the productivity and viability of microorganisms.

4. Pharmaceuticals

Buffer solutions are used in pharmaceuticals to create formulations where pH stability is critical to the efficacy and stability of the drug.

Calculation examples

Example 1: Calculating the pH of an acetic acid-sodium acetate buffer

Consider a buffer solution containing 0.1 M acetic acid and 0.1 M sodium acetate. pK a of acetic acid is 4.76. Calculate the pH of this buffer solution.

      pH = pK a + log([CH 3 COO - ]/[CH 3 COOH])
      pH = 4.76 + log(0.1/0.1)
      pH = 4.76 + log(1)
      pH = 4.76

The pH of the buffer solution is 4.76.

Example 2: Change in pH by adding acid

To the same solution, let us add a small amount of strong acid, say 0.01 M HCl. The new concentration will be:

      CH3COO- concentration : 0.1 M - 0.01 M = 0.09 M
      CH3COOH concentration: 0.1 M + 0.01 M = 0.11 M

Now, calculate the new pH:

      pH = 4.76 + log(0.09/0.11)
      pH = 4.76 + log(0.818)
      pH = 4.76 - 0.087
      pH = 4.673

After adding the strong acid the pH changes only slightly to 4.673, indicating the buffer capacity of the solution.

Buffer capacity

Buffer capacity refers to the ability of a buffer solution to resist changes in pH when an acid or base is added. It is determined by the concentration of the weak acid and its conjugate base. The higher the concentration, the greater the buffer capacity.

Limitations of buffer solutions

Buffer solutions have their limitations. Their ability to maintain a constant pH is not unlimited and is affected by:

  • Concentration of buffer components: More concentrated buffers have greater capacity.
  • Limitations of buffer capacity: Excessive acid or base can overwhelm the buffer.
  • Changes in temperature, which can affect the dissociation constant (K a or K b).

Conclusion

Buffer solutions are essential to chemistry and various scientific fields due to their stabilization properties. They provide a stable pH environment important for many chemical reactions, biological processes, and industrial operations. By understanding the components, functioning, capabilities, and limitations of buffer solutions, one can effectively apply them in both academic and practical scenarios.
Based on a simple and clear framework, with fundamental principles such as the Henderson-Hasselbalch equation, the world of buffers opens up myriad possibilities for scientific exploration and real-world application.


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Buffer Solution

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