Chemical Bonding and Molecular Structure Class 11 Full Chapter | NCERT Chemistry One Shot Revision

Chemical Bonding and Molecular Structure | Chemistry | Tapu Roamer

Key Concepts: Chemical Bond, Attractive Force, Causal Lewis Approach, Valence Shell, Kernel, Octet Rule, Duplet Rule, Lewis Symbols, Sharing of Electrons, Transfer of Electrons, Covalent Bond, Electrovalent Bond (Ionic Bond), Coordinate Bond (Dative Bond), Polar Covalent Bond, Non-Polar Covalent Bond, Electronegativity, Lone Pair.

Introduction to Chemical Bonding

• Chemical Bond: An attractive force that holds constituent particles (atoms, ions, molecules) together in different chemical species. It's not a physical line but a force of attraction.
• Theories of Chemical Bonding: Various theories attempt to explain the nature of chemical bonds, including:
  • Causal Lewis Approach
  • Valence Bond Theory (VBT)
  • Valence Shell Electron Pair Repulsion Theory (VSEPR)
  • Molecular Orbital Theory (MOT)
  • Hybridization

Causal Lewis Approach

• Focus: Explains which electrons participate in bonding.
• Electronic Concept of an Atom: Only certain electrons in an atom participate in bond formation.
• Valence Shell: The outermost shell of an atom, whose electrons participate in bonding.
• Kernel: The inner portion of the atom, excluding the valence shell, which does not participate in bonding. It contains the nucleus and inner electrons and has a positive charge equal to the number of valence electrons (z = p = e).
• Electronic Theory of Valence: Another name for the Causal Lewis Approach.

Octet Rule

• Postulate: Atoms combine to achieve a stable configuration with eight electrons in their outermost (valence) shell.
• Exceptions: Hydrogen and Helium achieve stability with two electrons (Duplet Rule).
• Noble Gas Configuration: Atoms strive to attain the electronic configuration of the nearest noble gas.
• Visual Representation: The outermost shell can be visualized as a cube with eight corners, each capable of holding one electron. A filled cube (octet) represents stability.
• Noble Gas Examples: Helium (2 electrons), Neon (8 electrons), Argon (8 electrons), Krypton (8 electrons), Xenon (8 electrons), Radon (8 electrons).

Lewis Symbols

• Purpose: To represent valence electrons in an atom using simple notations.
• Method: Write the element's symbol and represent valence electrons as dots around the symbol.
• Exclusion: Inner shell electrons (kernel) are not considered.
• Examples:
  • Hydrogen (H): H• (one valence electron)
  • Helium (He): He: (two valence electrons)
  • Lithium (Li): Li• (one valence electron)
  • Beryllium (Be): Be: (two valence electrons)

Achieving Stability: Sharing vs. Transfer of Electrons

• Motivation: Atoms strive to achieve a stable noble gas configuration (octet or duplet).
• Conditions:
  • Less than 8 electrons in the valence shell: Atom needs to gain electrons.
  • More than 8 electrons (considering the next shell): Atom needs to lose electrons.
• Methods to Achieve Stability:
  • Sharing of Electrons: Occurs between two non-metals.
  • Transfer of Electrons: Occurs between a metal and a non-metal.

Types of Chemical Bonds

• Covalent Bond: Formed by the sharing of electrons between two atoms (typically non-metals).

  • Mutual Sharing: Both atoms contribute electrons to the bond.
  • Single Bond: Sharing of one pair of electrons.
  • Double Bond: Sharing of two pairs of electrons.
  • Triple Bond: Sharing of three pairs of electrons.
  • Polar vs. Non-Polar Covalent Bonds:
    • Non-Polar: Equal sharing of electrons (e.g., Cl-Cl in Cl2). Forms molecules.
    • Polar: Unequal sharing of electrons due to electronegativity differences (e.g., H-Cl). Results in partial charges (δ+ and δ-).
    • Electronegativity: The tendency of an atom to attract electrons towards itself in a chemical bond.
    • Example: HCl: Chlorine is more electronegative than hydrogen, so it attracts the shared electrons, resulting in a partial negative charge (δ-) on chlorine and a partial positive charge (δ+) on hydrogen.

• Electrovalent Bond (Ionic Bond): Formed by the transfer of electrons from a metal to a non-metal.

• Coordinate Bond (Dative Bond): A special type of covalent bond where one atom provides both electrons for the shared pair.

  • Example: H3O+: A water molecule (H2O) has two lone pairs of electrons on the oxygen atom. A proton (H+) with no electrons can accept a lone pair from oxygen, forming a coordinate bond and creating the hydronium ion (H3O+).
  • Lone Pair: A pair of valence electrons that is not involved in bonding.
  • Example: NH4+: Ammonia (NH3) has one lone pair on Nitrogen. It can donate this lone pair to H+ to form Ammonium ion (NH4+).

Drawing Lewis Structures: Examples

• H2O (Water):
  1. Central Atom: Oxygen (less in number).
  2. Skeleton: H-O-H (single bonds).
  3. Valence Electrons: Hydrogen (1 each), Oxygen (6). Total = 8.
  4. Shared Pairs: One pair per single bond.
  5. Lone Pairs: Oxygen has two lone pairs.
  6. Octet/Duplet: Oxygen has 8 electrons, each hydrogen has 2.
• CH4 (Methane):
  1. Central Atom: Carbon.
  2. Skeleton: Carbon bonded to four Hydrogens.
  3. Valence Electrons: Carbon (4), Hydrogen (1 each). Total = 8.
  4. Shared Pairs: One pair per single bond.
  5. Octet/Duplet: Carbon has 8 electrons, each hydrogen has 2.
• N2 (Nitrogen):
  1. Two Nitrogen atoms.
  2. Each Nitrogen has 5 valence electrons.
  3. Each Nitrogen needs 3 more electrons to complete its octet.
  4. Therefore, they share 3 pairs of electrons, forming a triple bond.

Rules for Drawing Lewis Structures

1. Select the Central Atom: Generally the least electronegative atom or the atom present in the least number. Hydrogen and Fluorine are rarely central atoms.
2. Draw the Skeleton: Connect the central atom to the surrounding (corner) atoms with single bonds.
3. Calculate Total Valence Electrons: Sum the valence electrons of all atoms in the molecule.
4. Place Shared Electron Pairs: Place one pair of electrons between atoms connected by a single bond. Use two pairs for double bonds and three pairs for triple bonds.
5. Distribute Remaining Electrons: Distribute the remaining electrons as lone pairs to complete the octets of the atoms.
6. Verify Octets: Ensure that each atom (except hydrogen) has a complete octet of electrons.

Conclusion

The video provides a foundational understanding of chemical bonding, starting with the basic definition of a chemical bond as an attractive force between atoms, ions, or molecules. It then delves into the Causal Lewis approach, emphasizing the role of valence electrons and the concept of the kernel. The Octet Rule is explained as the driving force behind bond formation, with atoms striving to achieve a stable noble gas configuration. The video differentiates between sharing and transfer of electrons, leading to the discussion of covalent, electrovalent (ionic), and coordinate bonds. Polar and non-polar covalent bonds are distinguished based on electronegativity differences. Finally, the video outlines the rules for drawing Lewis structures, providing a step-by-step approach to represent molecular structures accurately. The key takeaway is that understanding these fundamental concepts is crucial for comprehending the properties and behavior of chemical compounds.