Grade 11
  1. Basic concepts of chemistry  1.1. Importance of Chemistry  1.2. Nature and scope of chemistry  1.3. laws of chemical combination  1.3.1. Law of conservation of mass  1.3.2. Law of definite proportions  1.3.3. Law of multiple proportions  1.3.4. Law of gaseous volume  1.3.5. Avogadro's Law  1.4. Dalton's atomic theory  1.5. Mole concept and molar mass  1.6. Percent composition  1.7. Empirical and molecular formula  1.8. Stoichiometry and Stoichiometric Calculations  1.9. Limiting reagent concept
  2. Structure of the atom  2.1. Discovery of the electron, proton, and neutron  2.2. Atomic Model  2.2.1. Thomson's model  2.2.2. Rutherford's model  2.2.3. Bohr's model  2.3. Quantum mechanical model of the atom  2.4. Dual nature of matter and radiation  2.5. Heisenberg's uncertainty principle  2.6. Quantum numbers  2.7. Atomic orbitals and their shapes  2.8. Electronic configuration of elements  2.9. Aufbau principle  2.10. Pauli's exclusion principle  2.11. Hund's maximum multiplicity rule  2.12. de Broglie's hypothesis  2.13. Photoelectric effect
  3. Classification of elements and periodicity in properties  3.1. Historical development of the periodic table  3.2. Modern Periodic Law and Periodic Table  3.3. Periodic trends in properties  3.3.1. Atomic and Ionic Radius  3.3.2. Ionization enthalpy  3.3.3. Electron gain enthalpy  3.3.4. Electronegativity  3.3.5. Valency  3.3.6. Metallic and Non-Metallic Properties  3.3.7. Oxidation states  3.4. Unusual Properties of the First Elements
  4. Chemical Bonding and Molecular Structure  4.1. Kossel–Lewis approach to chemical bonding  4.2. Ionic or electrovalent bond  4.3. Covalent bond and its features  4.4. Bond parameters  4.5. VSEPR Theory  4.6. Valence bond theory  4.7. Hybridization and its types  4.8. Molecular orbital theory  4.8.1. Bond order and stability  4.9. Hydrogen bonding
  5. States of matter  5.1. Intermolecular forces  5.2. Gas Laws  5.2.1. Boyle's law  5.2.2. Charles's Law  5.2.3. Avogadro's Law  5.2.4. Dalton's law of partial pressure  5.2.5. Graham's law of spread  5.3. Ideal Gas Equation and Applications  5.4. Real gases and deviations from ideal behaviour  5.5. Kinetic molecular theory of gases  5.6. Liquefaction of gases  5.7. Properties of Liquids  5.8. Surface tension and viscosity
  6. Thermodynamics  6.1. System and environment  6.2. Types of systems in thermodynamics  6.3. Types of procedures  6.4. First law of thermodynamics  6.5. Internal energy and enthalpy  6.6. Heat capacity and specific heat  6.7. Hess's law of constant heat summation  6.8. Bond dissociation enthalpy  6.9. Enthalpy of formation and combustion  6.10. Enthalpy of solution and neutralization  6.11. Spontaneity and the second law of thermodynamics  6.12. Gibbs free energy and its applications
  7. Balance  7.1. The concept of balance  7.2. Equilibrium constant and its applications  7.3. Le Chatelier's principle  7.4. Ionic equilibrium in solutions  7.5. Acid and Base Theory  7.5.1. Arrhenius concept  7.5.2. Bronsted-Lowry concept  7.5.3. Lewis concept  7.6. pH Scale and pOH  7.7. Common ion effect  7.8. Hydrolysis of salts  7.9. Buffer solutions and their mechanism of action  7.10. Solubility equilibrium of sparingly soluble salts
  8. Redox reactions  8.1. Oxidation and reduction concepts  8.2. Oxidation number and its applications  8.3. Redox reactions in terms of electron transfer  8.4. Balancing redox reactions (oxidation number and ion-electron methods)  8.5. Types of redox reactions  8.6. Electrochemical Cells and Redox Reactions
  9. Hydrogen  9.1. Position of Hydrogen in the Periodic Table  9.2. Isotopes of hydrogen  9.3. Preparation of Hydrogen  9.4. Properties of Hydrogen  9.5. Hydrides and their classification  9.6. Water and heavy water  9.7. Hydrogen peroxide and its properties  9.8. Uses of Hydrogen and Its Compounds
  10. S-block elements (alkali and alkaline earth metals)  10.1. Group 1 Elements - Properties and Trends  10.2. Group 2 Elements - Properties and Trends  10.3. Unusual properties of lithium and beryllium  10.4. Important compounds of sodium and their uses  10.5. Important Compounds of Magnesium and Calcium
  11. Some p-block elements  11.1. General Introduction to p-Block Elements  11.2. Group 13 Elements  11.2.1. Boron and its compounds  11.3. Group 14 Elements  11.3.1. Carbon and its allotropes  11.3.2. Important Compounds of Carbon and Silicon
  12. Organic Chemistry - Some Basic Principles and Techniques  12.1. Classification and Nomenclature of Organic Compounds  12.2. Structural representation of organic compounds  12.3. Classification of organic reactions  12.4. Electronic Effects in Organic Chemistry  12.4.1. Inductive effect  12.4.2. Resonance effect  12.4.3. Hyperconjugation  12.5. Reaction mechanism and intermediate species  12.6. Purification and qualitative analysis of organic compounds
  13. Hydrocarbons  13.1. Hydrocarbons  13.1.1. Preparation and Properties of Alkenes  13.2. alkene  13.2.1. Preparation and Properties of Alkenes  13.2.2. Electrophilic addition reactions  13.3. Alkynes  13.3.1. Preparation and Properties of Alkynes  13.4. Aromatic hydrocarbons  13.4.1. Structure and properties of benzene  13.4.2. Electrophilic substitution reactions of benzene
  14. Environmental Chemistry  14.1. Environmental Pollution  14.2. Atmospheric pollution and the greenhouse effect  14.3. Water pollution and treatment methods  14.4. Soil pollution and its effects  14.5. Industrial waste and its treatment  14.6. Strategies for pollution control

Grade 11 → Balance ↓


Common ion effect


The common ion effect is an important concept in chemistry. It refers to the shift in equilibrium that occurs when a compound that has a common ion with the dissolved substance is added to a solution. It is a specific application of Le Chatelier's principle, which states that if a system in equilibrium is disturbed, it will adjust itself to counteract the disturbance and restore a new equilibrium.

Understanding the concept

The common ion effect can be explained through various examples and applications in chemistry. It is especially important when dealing with poorly soluble salts and weak acids or bases. Before diving into the examples, let's consider the equilibrium concepts involved.

Chemical equilibrium basics

Chemical equilibrium occurs when the forward and reverse reactions occur at the same rate. For a general reaction:

A + B ⇌ C + D

The equilibrium constant (K_c) for the reaction is defined as:

K_c = [C][D] / [A][B]

Now, when the common ion is introduced to the components of this equilibrium state, it affects the position of the equilibrium state by changing the concentration of the products or reactants.

Role of the common ion effect

When a salt, which shares a common ion with a solute already present in the solution, is added, the solubility of the solute decreases. This happens because the added common ion increases the ion concentration on one side of the equilibrium equation. As a result, the equilibrium will shift in favor of the formation of the reactant side to achieve a new equilibrium.

Example

Example 1: Common ion effect in solubility

Consider a saturated solution of calcium sulphate (CaSO_4):

CaSO_4 (s) ⇌ Ca^{2+} (aq) + SO_4^{2-} (aq)

Now, let's add another compound, such as sodium sulfate (Na_2SO_4), which provides an additional source of sulfate ions (SO_4^{2-}) in the solution. The equilibrium will shift to the left according to Le Chatelier's principle, resulting in a decreased solubility of CaSO_4.

CaSO₄ Na₂SO₄ Ca2+SO4

Example 2: Common ion effect with weak acids

Consider acetic acid (CH_3COOH) in water. The equilibrium expression is:

CH_3COOH (aq) ⇌ CH_3COO^- (aq) + H^+ (aq)

If sodium acetate (CH_3COONa) is added to the solution, it will dissociate and provide additional acetate ions (CH_3COO^-). This additional source of acetate ions will push the equilibrium to the left, resulting in fewer hydrogen ions being produced, which will decrease the acidity of the solution.

CH₃COOH CH₃COONa CH₃COO⁻ + H⁺

Mathematical representation

For a general equilibrium reaction involving a common ion, consider a solution of a weak electrolyte in water:

AB(s) ⇌ A^+ (aq) + B^- (aq)

Adding an electrolyte that shares a common ion, say B^−, shifts the equilibrium:

[B^-]_{initial} rightarrow [B^-]_{new} = [B^-]_{initial} + [common  ion]

With an increase in [B^-], the reaction quotient changes, and the system responds by shifting the equilibrium position to maintain K_c constant according to the equilibrium constant expression. This indicates a decrease in the solubility of the solute.

Applications of common ion effect

The common ion effect has many applications in chemistry, including:

  • Buffer solutions: Buffer solutions are able to resist changes in pH when small amounts of acid or base are added. This is because they contain a weak acid (or base) and its conjugate base (or acid), which produce a common ion effect.
  • Precipitation of salts: It is used to precipitate salts by adding a common ion to the solution, thereby decreasing the solubility of the salt in a controlled manner.
  • Industrial processes: The common ion effect is used in a variety of industrial processes where the control of solubility and precipitation is important, such as water treatment and purification.

Conclusion

The common ion effect is an interesting and highly relevant concept in the field of chemistry, especially for those who deal with equilibrium and solution chemistry. Understanding this effect provides insight into how the presence of a common ion affects the concentration of species in a solution, which ultimately affects solubility, pH, and the ability to control chemical processes.


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Common ion effect

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