Chemical Bonding

Valence Electrons and Bonding

Valence electrons are the electrons in the outermost shell of an atom and are primarily involved in chemical bonding. They occupy the highest energy level and determine the chemical reactivity and bonding capacity (valency) of the atom.

• Valence electrons: Electrons in the outermost shell, involved in bonding.

• Valency: The number of electrons an atom can share, lose, or gain to achieve a stable configuration.

• Example: Hydrogen has 1 valence electron (1s1), so its valency is 1.

Ionic Bonding

Formation and Characteristics

Ionic bonds are formed between metals and non-metals through the transfer of electrons, resulting in the formation of positively and negatively charged ions. This process leads to filled valence orbitals and stable electronic configurations.

• Occurs between: Metal (+ively charged, e.g., +1 or +2) and non-metal (-ively charged, e.g., -1 or -2).

• Mechanism: Atoms bond ionically by giving up or accepting electrons.

• Example: Sodium chloride (NaCl): Sodium (Na) electron configuration: Na loses one electron to become Na+, Cl gains one to become Cl-.

Covalent Bonding

Formation and Valency

Covalent bonds involve the sharing of electrons between two atoms, typically non-metals. The number of shared electrons corresponds to the valency of the atoms involved.

• Sharing of electrons: Electrons are shared between atoms to achieve stable configurations.

• Valency: Determined by the number of electrons needed to complete the outer shell.

• Examples:

  • Hydrogen: 1 electron (), valency = 1

  • Oxygen: 8 electrons (), valency = 2

  • Nitrogen: 7 electrons (), valency = 3

  • Carbon: 6 electrons (), valency = 4

Types of Covalent Bonds

• Single bonds: Sharing of one pair of electrons (e.g., H–H, C–H).

• Double bonds: Sharing of two pairs of electrons (e.g., C=C, C=O).

• Triple bonds: Sharing of three pairs of electrons (e.g., C≡C, C≡N).

Polarity and Electronegativity

Polar and Nonpolar Covalent Bonds

The polarity of a covalent bond depends on the difference in electronegativity between the bonded atoms. Unequal sharing leads to polar covalent bonds, while equal sharing results in nonpolar covalent bonds.

• Nonpolar covalent bond: Electronegativity difference < 0.5 (e.g., C–H, C–C).

• Polar covalent bond: Electronegativity difference between 0.5 and 1.9 (e.g., C–O, C–Cl).

• Ionic bond: Electronegativity difference > 1.9 (e.g., NaCl).

Electronegativity Values

Example: In HCl, H (2.1) and Cl (3.0) have a difference of 0.9, resulting in a polar covalent bond and a dipole moment.

Formal Charge (FC)

Calculation and Examples

Formal charge is used to determine the distribution of electrons in molecules and ions. It helps identify the most stable Lewis structure.

• Formula:

• Examples:

  • Hydronium ion (H3O+):

  • Ammonium ion (NH4+):

  • Methane (CH4):

  • Methyl anion (CH3-):

  • Acetonitrile (CH3CN): , ,

3D Structure of Molecules

VSEPR Theory and Molecular Geometry

The three-dimensional structure of molecules is determined by the number of bonds and lone pairs around the central atom, known as the steric number (SN). VSEPR theory predicts the geometry based on electron pair repulsion.

• Steric number (SN): Number of bonds plus lone pairs around an atom.

• Common geometries:

  • SN = 4: Tetrahedral, bond angle ≈ 109.5°

  • SN = 3: Trigonal planar, bond angle ≈ 120°

  • SN = 2: Linear, bond angle ≈ 180°

• Example: Methane (CH4) is tetrahedral; ethylene (C2H4) is trigonal planar; acetylene (C2H2) is linear.

Types of Covalent Bonds: Sigma and Pi Bonds

Sigma (σ) Bonds

Sigma bonds are formed by the head-on overlap of atomic orbitals, typically s-orbitals. All single bonds are sigma bonds.

• Formation: Overlap of two s-orbitals.

• Properties: Strongest type of covalent bond; allows free rotation.

Pi (π) Bonds

Pi bonds are formed by the side-to-side overlap of two parallel p-orbitals. They occur in double and triple bonds, in addition to a sigma bond.

• Formation: Overlap of two parallel p-orbitals.

• Properties: Weaker than sigma bonds; restrict rotation.

• Example: Ethylene (C2H4) has one sigma and one pi bond between the carbons; acetylene (C2H2) has one sigma and two pi bonds.

Hybridization of Orbitals

Concept and Application

Hybridization is the mixing of atomic orbitals to form new hybrid orbitals suitable for bonding. This explains the observed shapes and bond angles in molecules.

• Carbon hybridization:

  • sp3: Tetrahedral (e.g., methane, CH4)

  • sp2: Trigonal planar (e.g., ethylene, C2H4)

  • sp: Linear (e.g., acetylene, C2H2)

• Example: Carbon's ground state electron configuration is . In methane, carbon undergoes sp3 hybridization to form four equivalent bonds.

Summary Table: Bond Types and Properties

Additional info: These notes provide foundational concepts for understanding chemical bonding, molecular structure, and electron distribution, which are essential for organic chemistry studies.