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 ↓


Ionic equilibrium in solutions


In chemistry, the concept of equilibrium is important for understanding reactions that occur in solutions. When we talk about ionic equilibrium, we are specifically dealing with reactions involving ions in water. Many chemical substances, when dissolved in water, dissociate into ions. Ionic equilibrium occurs when the rate of formation of ions is equal to the rate of recombination of the ions into their uncombined form. This equilibrium point is where the properties of the solution, such as pH, reach a stable value.

Basic concepts of ionic equilibrium

What are ions?

Ions are charged particles that form when atoms gain or lose electrons. There are two types of ions:

  • Cations: Positively charged ions formed by the loss of electrons. For example, N a^+ and C a^{2+}.
  • Anions: Negatively charged ions formed by the gain of electrons. For example, C l^− and S O_4^{2−}.

Electrolytes

Electrolytes are substances that turn into ions when dissolved in water. They can be broadly classified into two categories:

  • Strong electrolytes: completely dissociate into ions. Examples include strong acids such as HC l, strong bases such as N a OH, and salts such as N aC l.
  • Weak electrolytes: partially dissociate into ions. Examples include weak acids such as C H_3 COOH (acetic acid) and weak bases such as NH_3 (ammonia).

Understanding equilibrium in ionic solutions

The concept of balance

Equilibrium in the chemical sense involves a situation in which the concentrations of reactants and products remain constant over time. This occurs because the forward and reverse reaction rates are the same. In the case of ionic equilibrium, we are mainly dealing with the dissociation and recombination of ions.

Dynamic nature of equilibrium

It is important to remember that equilibrium is a dynamic state. Even though the macroscopic properties of a system do not change at equilibrium, processes of dissociation and recombination continue to occur at the microscopic level.

Equilibrium constant

For a general reaction at equilibrium:

    AA + BB ⇌ CC + DD

The expression for the equilibrium constant is written as follows:

    Ke = [C]^c [D]^d / [A]^a [B]^b

Here, [C], [D], [A], and [B] represent the molar concentrations of the respective species. The exponents correspond to their coefficients in the balanced chemical equation. This constant helps chemists understand the extent of the reaction at equilibrium.

Dissociation of weak electrolytes

Acids and bases

Weak acids and bases do not ionize completely in solution. The degree of ionization is an important aspect of ionic equilibrium. Consider acetic acid C H_3 COOH dissociating in water:

    CH_3 COOH (aq) ⇌ H^+ (aq) + CH_3 COO^- (aq)

The equilibrium constant for this reaction, known as the acid dissociation constant, is represented by K_a.

Similarly, for a weak base such as ammonia in water:

    N H_3 (aq) + H_2 O (l) ⇌ N H_4^+ (aq) + OH^- (aq)

The equilibrium constant for this reaction is known as the base dissociation constant, K_b.

Calculating pH in ionic equilibrium

pH scale

pH is a measure of the hydrogen ion concentration [H^+] in a solution. The pH scale ranges from 0 to 14, with lower values being more acidic, higher values being more alkaline, and 7 being neutral.

pH is calculated using the following formula:

    pH = -log[H^+]

Example of pH calculation

Consider a weak acid, such as acetic acid, whose concentration and K_a value are known. To find the pH of a solution containing 0.1 M acetic acid:

phase:

  1. Establish the equilibrium expression using K_a value.
  2. Assume that [H^+] is equal to [C H_3 COO^-]. Also, [C H_3 COOH] is about 0.1 - x, where x is the extent of ionization.
  3. Solve for [H^+] using K_a expression.
  4. Calculate pH using the formula pH = -log[H^+].

Common-ion effect

The common-ion effect refers to the decrease in the solubility of a substance due to the presence of a common ion. This phenomenon can alter the equilibrium position, affecting the degree of ionization.

Example:

Consider a solution of acetic acid with a salt such as sodium acetate. According to Le Chatelier's principle, the presence of C H_3 COO^- from sodium acetate suppresses the ionization of acetic acid.

Effectively, the equilibrium expression is:

    CH_3 COOH (aq) ⇌ H^+ (aq) + CH_3 COO^- (aq)

Buffer solution

Buffer solutions are an essential application of ionic equilibrium. They are able to maintain a relatively constant pH when small amounts of acid or base are added.

Components of the buffer

  • A weak acid and its conjugate base, such as acetic acid and sodium acetate.
  • A weak base and its conjugate acid, such as ammonia and ammonium chloride.

Function of buffer solution

Consider a buffer composed of acetic acid C H_3 COOH and sodium acetate C H_3 COON a:

  • If hydrogen ions (H^+) are added, they react with acetate ions (C H_3 COO^-) to form acetic acid, minimizing the pH change.
  • If hydroxide ions (OH^-) are added, they react with acetic acid (C H_3 COOH) to form acetate and water, minimizing the pH change.

Hydrolysis of salts

Hydrolysis means the reaction of an ion with water to form a solution whose pH is different from the expected pH based on the original acid and base. This process is necessary in ionic equilibrium for salt solutions.

Types of salts

  • Salts from strong acids and strong bases: These do not hydrolyse and the solution remains neutral. Example: N a C l.
  • Salts from strong acids and weak bases: These form acidic solutions. Example: N H_4 C l.
  • Salts from weak acids and strong bases: These form alkaline solutions. Example: C H_3 COON a.
  • Salts from weak acids and weak bases: pH depends on the relative strengths of the acid and base. Example: (N H_4)(C H_3 COO).

Hydrolysis constant

The equilibrium constant for hydrolysis is known as the hydrolysis constant K_h. Consider the hydrolysis of N H_4^+:

    N H_4^+ (aq) + H_2 O (l) ⇌ N H_3 (aq) + H_3 O^+ (aq)

The hydrolysis constants K_h of conjugate acids and bases can be expressed in terms of K_w (ion product of water), K_a, and K_b.

Solubility product

Definition

The solubility product, denoted as K_{sp}, is a type of equilibrium constant applied to the solubility of ionic compounds. It is defined for sparingly soluble salts.

Example of solubility product

Consider the dissolution of silver chloride, AgCl, in water:

    AgCl (s) ⇌ Ag^+ (aq) + Cl^- (aq)

K_{sp} expression for silver chloride is:

    K_{sp} = [Ag^+][Cl^-]

This expression helps in predicting whether a precipitate will form when two solutions containing ions are mixed.

Visual example

Dynamic equilibrium of aqueous solutions

H2O H + + OH - (balance)

Buffer action visualization

CH 3 COO - H + Oh -

Applications in everyday life

Ionic balance has many applications in daily life. They play an important role in maintaining the pH of the human body, industries and environmental chemistry.

Biological significance

Buffers in the blood, such as carbonic acid and bicarbonate, help maintain blood pH around 7.4. This regulation is important for bodily processes.

Industrial applications

Many industrial processes require precise pH control. For example, in the manufacture of pharmaceuticals, food processing, and water treatment facilities.

Environmental chemistry

Understanding ionic balance is essential in analyzing natural waters and predicting the effects of pollutants. Efforts to reduce acid rain or treat wastewater rely heavily on the principles of ionic balance.

Summary

Ionic equilibrium in solutions is a fundamental concept in chemistry that explains the balance of ions in aqueous solutions. By understanding the equilibrium constant, the principles of acid-base chemistry, and the behavior of buffers, one can predict and manipulate the chemical properties of solutions in a variety of scientific and industrial contexts.


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Ionic equilibrium in solutions

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