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 ↓


Lewis concept


Chemistry often offers different approaches to understanding the behaviour of substances. The study of acids and bases is one area where several theories provide complementary insights. The Lewis concept is a sophisticated approach that extends our understanding of acids and bases beyond the traditional hydrogen-based model.

Introduction to the Lewis concept

The Lewis concept, named after Gilbert N. Lewis, was developed in the early 20th century. This approach considers the behavior of electron pairs rather than hydrogen ions, which allows it to be more widely applied to a wider range of chemical reactions. In the Lewis model, substances are classified based on their ability to accept or donate electron pairs.

Definitions

According to Lewis concept:

  • A Lewis acid is a chemical compound that can accept an electron pair.
  • A Lewis base is a chemical compound that can donate an electron pair.

Understanding the mechanism

The essence of Lewis acid-base interactions can be understood by considering electron pair exchange. Let us consider some examples to make this concept clear:

Example: ammonia and boron trifluoride

The reaction between ammonia (NH 3) and boron trifluoride (BF 3) is a classic example:

        NH 3 + BF 3 → NH 3 → BF 3

In this example, ammonia acts as a Lewis base because it donates its electron pair, while boron trifluoride acts as a Lewis acid because it accepts that electron pair.

Example: water and hydrogen ions

The interaction of water with a hydrogen ion is another great example:

        H 2 O + H + → H 3 O +

Here, water (H 2 O) donates a pair of electrons to form a coordinate covalent bond with the hydrogen ion (H +), making water a Lewis base and the hydrogen ion a Lewis acid.

Visualization of electron pair interactions

To better understand these interactions, consider the following diagram showing the reaction of ammonia with a hydrogen ion:

, NH 3 H + NH4 +

In this visual illustration, the arrows show the flow and transfer of the electron pair from the ammonia to the hydrogen ion.

Textual examples of Lewis acids and bases

Boron compounds

Many boron compounds act as Lewis acids due to the electron-deficient nature of boron. In the example of boron trifluoride:

        BF 3 + :NH 3 → BF 3 NH 3

Boron trifluoride lacks enough electrons to achieve a stable octet, making it a powerful Lewis acid, accepting electrons from a donor.

Metal ions

Metal ions such as Cu 2+, Al 3+, and Fe 3+ are also excellent examples of Lewis acids because they can easily accept electron pairs from ligands (Lewis bases) to form coordination complexes.

        Cu 2+ + 4NH 3 → [Cu(NH 3) 4 ] 2+

Comparison with other acid–base theories

Compared to the Arrhenius and Brønsted–Lowry theories, the Lewis concept provides a broader definition:

  • The Arrhenius theory focuses on the production of H + ions and OH - ions in water.
  • The Bronsted-Lowry theory defines acids as proton donors and bases as proton acceptors.
  • Lewis theory extends this concept to include reactions that do not involve protons, and instead focuses on the transfer of electron pairs.

For example, consider the reaction of sulfur trioxide and oxygen:

        SO 3 + O 2- → SO 4 2-

Such reactions do not fit well into the Arrhenius or Bronsted-Lowry definitions, but are elegantly explained by the Lewis concept through electron pair acceptance and donation.

Importance of Lewis concept in chemistry

The Lewis concept is fundamental in understanding a wide variety of chemical reactions, including complex formation and catalysis. It also provides a unifying approach to understanding diverse chemical phenomena such as acid rain, industrial synthesis, and biological processes. For example, enzymes can often be explained as Lewis bases that interact with metal ion cofactors, which act as Lewis acids.

Conclusion

The Lewis concept significantly enriches our understanding of acid-base chemistry, providing explanations beyond the limitations of earlier models. With its focus on electron pair interactions, this theory helps explain complex chemical processes, making it a cornerstone in the study of chemistry. Whether in the laboratory, in industry, or in the natural world, Lewis theory provides a powerful framework for explaining a wide range of chemical contingencies.


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Lewis concept

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