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 ↓


Solubility equilibrium of sparingly soluble salts


Introduction to solubility equilibrium

Sparingly soluble salts are those that dissolve in water to a minimal extent. Understanding their solubility is important in many fields, including chemistry, environmental science, and engineering. Solubility equilibrium helps us determine how these salts behave in aqueous systems and calculate their solubility product.

Basic definitions and concepts

Before we discuss solubility equilibrium, let us discuss some fundamental concepts.

Solubility

Solubility is the maximum amount of solute that can dissolve in a solvent at a given temperature, resulting in a saturated solution. For less soluble salts, this amount is much less, often expressed as moles per liter (mol/L).

Saturated solution

A saturated solution contains the maximum concentration of solute that can dissolve at a particular temperature. Any more solute added will not dissolve and result in the formation of additional solid.

Solubility product constant (K sp)

The solubility product constant, K sp, is an equilibrium constant that applies to the solubility of sparingly soluble salts. It shows the extent to which the compound dissociates in water.

Expression for K sp

To split a common salt AB into its ions:

AB(s) ⇋ A + (aq) + B - (aq)

K sp expression is given as:

K sp = [A + ][B - ]

Visual examples of solubility equilibrium

Consider the solubility equilibrium of silver chloride (AgCl):

AgCl(s) ⇋ Ag + (aq) + Cl - (aq)
AgCl2 undissolved Ag+ Cl-

Here, K sp = [Ag + ][Cl - ]

General calculations related to K sp

Let us see how the solubility product constant is calculated for different situations.

Determining solubility from K sp

For a salt AB where K sp = s^2 :

s = √(K sp )

For example, if K sp is given for AgCl, then consider AgCl ⇋ Ag + (aq) + Cl - (aq) :

K sp = s^2

Calculating ion concentration

If the solubility s is given then:

[A + ] = [B - ] = s

Illustrative examples

Consider calcium sulfate (CaSO₄):

CaSO₄(s) ⇋ Ca 2+ (aq) + SO₄ 2- (aq)

Given K sp = 2.4 × 10 -5, solve for s:

s = √(2.4 × 10 -5 ) = 4.9 × 10 -3 mol/L

Factors affecting solubility

To apply these concepts in practical scenarios it is necessary to understand what factors affect solubility.

Common-ion effect

When a common ion is already present in a solution, the solubility of the salt decreases.

For example, in a solution containing NaCl, the solubility of AgCl is reduced due to the Cl - ion common to both salts.

pH of the solution

Solubility can be increased or decreased depending on the pH of the solution. In some cases, the addition of an acid or base affects the equilibrium.

Illustrative example using pH

For lead carbonate (PbCO₃), the addition of acid shifts the equilibrium:

PbCO₃(s) ⇋ Pb 2+ (aq) + CO₃ 2- (aq)

The H + from the acid reacts with CO₃ 2-, increasing the solubility.

Le Chatelier's principle and solubility

Le Chatelier's principle states that if an external condition changes, causing a shift in equilibrium, the system will adjust to counteract this change.

Applications to solubility

When the concentration of an ion increases, the equilibrium shifts, causing more of the solid to form, which decreases solubility, and vice versa.

Real-world applications

The solubility equilibrium of sparingly soluble salts has many practical applications.

Water treatment

Understanding solubility helps in removing unwanted ions from water by precipitating them as insoluble salts.

Medicines

Sparingly soluble compounds are used to control drug delivery, where solubility determines the release rate.

Practice problems

Practicing calculations related to solubility equilibrium will deepen your understanding. Here are some problems you can solve:

Problem 1

K sp = 1.1 × 10 -10. Calculate the solubility of BaSO₄ in mol/L.

Problem 2

If the solubility of PbI₂ is 1.3 × 10 -3 mol/L, find K sp.

Problem 3

Find the new solubility of AgCl in the presence of 0.1 M NaCl solution.

Conclusion

Understanding the solubility equilibrium of sparingly soluble salts is essential for predicting how different salts will behave under different conditions. Along with practical applications, recognizing the factors that affect solubility provides deep insight into chemical reactions in natural and industrial processes.


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Solubility equilibrium of sparingly soluble salts

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