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 → Chemical Bonding and Molecular Structure ↓


Kossel–Lewis approach to chemical bonding


The Kossel-Lewis approach to chemical bonding is one of the fundamental concepts in understanding how atoms combine to form molecules. This theory highlights the role of electrons in making or breaking bonds and aims to explain how atoms achieve stable electronic configurations, often called octet configurations. Let us discuss this approach in detail.

Historical context

In the early 20th century, chemists Walter Kossel and Gilbert N. Lewis independently presented theories about chemical bonding centered around electrons. They observed that the most stable compounds form when atoms achieve a stable electronic configuration similar to that of the noble gases. This observation was the birth of the octet rule.

Octet rule

The octet rule requires atoms to form bonds in such a way that they have eight electrons in their valence shell, resulting in a stable electronic configuration similar to that of the noble gases. This can be achieved through sharing, gaining, or losing electrons.

        General representation of the Octet Rule: Noble Gas Configuration: ns 2 np 6 Where n represents the principal quantum number indicating the outermost shell.

Nature of chemical bonds

According to the Kossel-Lewis theory, different types of chemical bonds arise due to the interaction of electrons between atoms.

Ionic bond

Walter Kossel focused primarily on the formation of ionic bonds. Ionic bonds are formed when there is a complete transfer of electrons from one atom to another, resulting in the formation of charged ions. For example, when a sodium (Na) atom transfers electrons to a chlorine (Cl) atom, they form a sodium ion (Na +) and a chloride ion (Cl -), resulting in sodium chloride (NaCl).

        Example of Ionic Bond Formation: Na (1s 2 2s 2 2p 6 3s 1 ) + Cl (1s 2 2s 2 2p 6 3s 2 3p 5 ) ➡ Na + + Cl - ➡ NaCl

This transfer results in electrostatic attraction between the oppositely charged ions, holding them together in an ionic lattice.

Covalent bonds

Gilbert N. Lewis primarily discussed covalent bonds, where atoms share pairs of electrons to satisfy the octet rule. In covalent bonding, the shared electrons hold the atoms together. A classic example of this is the hydrogen molecule (H 2), where two hydrogen atoms form a stable molecule by sharing their lone electrons.

        Example of Covalent Bond Formation: H (1s 1 ) + H (1s 1 ) ➡ H 2

For a visual representation, consider the water molecule (H 2O), where one oxygen atom shares electrons with two hydrogen atoms:

        [ O ]--[ H ] | [ H ]

Lewis structures

Lewis structures are diagrams that show the bonds between the atoms of a molecule and the lone pairs of electrons that are present in the molecule. They are simple ways of showing where and how electrons are being shared or transferred, and they make understanding chemical bonding much easier.

Follow these steps to draw a Lewis structure:

  1. Identify the total number of valence electrons in the molecule.
  2. Arrange the atoms, determine the central atom (usually the least electronegative).
  3. Use pairs of electrons to form bonds between atoms, and try to adjust all the atoms to satisfy the octet rule.
  4. If any electron is left then place it as a lone pair on the central atom or elsewhere if required.
  5. Double or triple bonds may be necessary to obtain an octet.

Example: For carbon dioxide (CO 2):

        Total valence electrons = 4 (Carbon) + 6*2 (Oxygen) = 16 O=C=O Oxygen forms double bonds with carbon to satisfy the octet.
    
         .. O=C=O ..

Limitations and exceptions of the octet rule

Although the octet rule is widely applied, there are some notable exceptions, including:

Incomplete octet

Some elements are stable with fewer than eight electrons in their valence shell. Boron in BF 3 is a common example.

        BF 3 Lewis Structure: F | B--F | F Boron is stable with six valence electrons.

Extended octet

Elements in the third period and beyond can expand their valence shells to have more than eight electrons. Consider PF 5:

        PF 5 Lewis Structure: F | F--P--F | FF Phosphorus can hold 10 valence electrons.

Free radicals

Molecules with an odd number of electrons have unpaired electrons and are called free radicals. For example, nitric oxide (NO):

        NO Lewis Structure: .. N--O Nitrogen has seven valence electrons leading to an unpaired electron.

Conclusion

The Kossel-Lewis approach to chemical bonding provides important information about the nature of chemical bonds. By considering the transfer and sharing of electrons, it provides a framework for understanding molecular structures and properties. Despite its limitations and exceptions, the octet rule serves as a useful tool for predicting how atoms bond, leading to the formation of molecules with diverse properties.

Overall, Kossel and Lewis's contributions mark an important milestone in the field of chemical bonding, establishing foundational concepts that continue to inform modern chemistry.


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Kossel–Lewis approach to chemical bonding

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