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 ↓


Acid and Base Theory


Understanding acids and bases is a central part of chemistry. They are two classes of compounds that have important effects on chemical reactions, biological systems, and the environment. Several theories help explain the behavior of acids and bases, including the Arrhenius theory, the Bronsted-Lowry theory, and the Lewis theory. These theories highlight different aspects of acid-base behavior and expand the scope of what is considered acidic or alkaline.

Arrhenius theory

The Arrhenius theory, named after Swedish scientist Svante Arrhenius, is one of the earliest models developed to explain acids and bases. According to this theory:

  • An acid is a substance that increases the concentration of hydrogen ions, H +, in an aqueous solution.
  • A base is a substance that increases the concentration of hydroxide ions, OH - in an aqueous solution.

For example, when hydrochloric acid (HCl) dissolves in water, it dissociates into hydrogen ions and chloride ions:

HCl (aq) → H + (aq) + Cl - (aq)

Sodium hydroxide (NaOH), a base, dissociates in water to form sodium ions and hydroxide ions:

NaOH (aq) → Na + (aq) + OH - (aq)

The Arrhenius theory is straightforward but is limited to reactions that occur in aqueous solutions. It does not take into account acid-base reactions that occur in other solvents or in the absence of a solvent.

Bronsted–Lowry theory

Developed by Johannes Bronsted and Thomas Lowry in 1923, the Bronsted-Lowry theory extends the concept of acids and bases beyond aqueous solutions. According to this theory:

  • An acid is a substance that can donate protons (H +).
  • A base is a substance that can accept a proton.

In this context, water can act as both an acid and a base depending on the reaction. This theory also introduces the concept of conjugate acid-base pair. When an acid donates a proton, it becomes a conjugate base. When a base accepts a proton, it becomes a conjugate acid.

For example, consider the reaction between ammonia and water:

NH 3 (aq) + H 2 O (l) ⇌ NH 4 + (aq) + OH - (aq)

In this scenario, NH 3 acts as a base by accepting a proton from water, which acts as an acid. After donating its proton, water becomes hydroxide (OH -), which is the conjugate base, while ammonia becomes NH 4 +, which is the conjugate acid.

Lewis theory

The Lewis theory, proposed by Gilbert N. Lewis in 1923, provides a comprehensive definition of acids and bases:

  • Lewis acid is an electron pair acceptor.
  • Lewis base is an electron pair donor.

This theory focuses on the transfer of electron pairs rather than protons. Because of this, Lewis theory can explain reactions that do not involve hydrogen ions. It also introduces concepts such as coordinate covalent bonds, where the two electrons come from the same atom.

Consider the reaction between boron trifluoride (BF 3) and ammonia (NH 3):

BF 3 + NH 3 → F 3 BNH 3

In this reaction, BF 3 acts as a Lewis acid by accepting an electron pair from NH 3, which is a Lewis base. The electron pair donation allows for the formation of a covalent bond between the two molecules.

Equilibrium and acid-base reactions

Many acid-base reactions reach a state of equilibrium, meaning that the rate of the forward reaction equals the rate of the reverse reaction. At equilibrium, the concentrations of reactants and products remain constant over time.

The equilibrium constant (K_a for acids and K_b for bases) measures the strength of acids and bases in solution:

  • For acids:
K_a = [H + ][A - ] / [HA]

where [HA] is the concentration of the acid, [A - ] is the concentration of the conjugate base, and [H + ] is the concentration of hydrogen ions.

  • For the bases:
K_b = [BH + ][OH - ] / [B]

where [B] is the concentration of the base, [BH + ] is the concentration of the conjugate acid, and [OH - ] is the concentration of hydroxide ions.

The larger K_a or K_b value, the stronger the acid or base. Strong acids and bases dissociate completely in water, while weak acids and bases partially dissociate, establishing an equilibrium between the uncombined and dissociated forms.

Visual example

Consider the Bronsted-Lowry reaction between acetic acid and water:

CH 3 COOH (aq) + H 2 O (l) ⇌ CH 3 COO - (aq) + H 3 O + (aq)
CH 3 COOH H2O CH 3 COO - H 3 O +

In this reaction, acetic acid (CH 3 COOH) donates a proton to water, forming an acetate ion (CH 3 COO -) and a hydronium ion (H 3 O +). Both acetic acid and acetate form a conjugate acid-base pair. Similarly, water and hydronium form another conjugate pair.

Applications and significance

Understanding acids and bases is essential to many scientific fields. For example, in biology, maintaining a balance between acids and bases (pH) is important for enzyme activity and cellular function. In industry, acid-base reactions create fertilizers, pharmaceuticals, and chemicals. Environmental science focuses on the effects of acid rain and ocean acidification, both of which are affected by acid-base chemistry.

Water treatment processes, food chemistry, and even our everyday experiences, such as the sour taste of lemon or the bitter taste of baking soda, involve acid-base reactions.

Conclusion

The theories of acids and bases – Arrhenius, Bronsted-Lowry and Lewis – provide several approaches to understanding these important chemical compounds. Each theory expands on the concept of what an acid or base is, allowing scientists to explain and predict a wide range of chemical reactions in a variety of environments. This multifaceted understanding is important for applications in industry, biology, environmental science and beyond. Recognizing the equilibrium that exists in acid-base reactions helps chemists manipulate the reactions to achieve desired results, making the study of acids and bases fundamental to both theoretical and applied chemistry.


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Acid and Base Theory

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