Chemical Bonding

A. Introduction

• Chemical Bond: It is the force that holds two or more atoms together to form a stable molecule.
• Atoms: The smallest unit taking part in a chemical reaction, consisting of protons (+ charge), neutrons (no charge), and electrons (- charge).
• Metallic Elements: Have 1, 2, or 3 valence electrons. They tend to lose these electrons to become positively charged ions called cations.
• Non-Metallic Elements: Have 4, 5, 6, or 7 valence electrons. They tend to gain electrons to become negatively charged ions called anions.
• Chemical Combination: Occurs in two ways:
  • Transfer of valence electrons (metallic to non-metallic atom).
  • Sharing of valence electrons (between two non-metallic atoms).

B. Chemical Bond & Types of Bonding

• Noble Gases: Inherently stable, chemically inert, and unreactive because their outermost shell is complete (2 electrons for Helium, 8 for others).
• Reason for Bonding: Atoms combine to achieve a stable electronic configuration matching their nearest noble gas.
  • Duplet Rule: Having two electrons in the outermost shell (like Helium).
  • Octet Rule: Having eight electrons in the outermost shell.
• Methods of Bonding:
  • Electrovalent/Ionic Bonding: By transfer of electrons.
  • Covalent Bonding: By sharing of electron pairs.
• Periodic Properties Affecting Bonding:
  • Ionic Bonds form easily when there is a large difference in electronegativity (low ionization potential for metals, high electron affinity for non-metals).
  • Covalent Bonds form easily when both atoms have high ionization potential and electron affinity, with a negligible electronegativity difference.

C. Electrovalent Bonding

• Formation: Involves the complete transfer of valence electrons from a metallic electropositive atom to a non-metallic electronegative atom.
• Electrostatic Force: The strong attraction between the resulting positively charged cation and negatively charged anion forms the electrovalent (ionic) bond.
• Electrovalency: The exact number of electrons donated or accepted by an atom to achieve a stable configuration.
• Redox Reactions in Bonding:
  • Oxidation: The process where an atom or ion loses electrons (e.g., Na loses 1e⁻ to become Na⁺).
  • Reduction: The process where an atom or ion gains electrons (e.g., Cl gains 1e⁻ to become Cl⁻).

D. Structure of Electrovalent Compounds

• Sodium Chloride (NaCl):
  • Sodium (2, 8, 1) loses 1 electron to achieve Neon's stability, becoming Na⁺.
  • Chlorine (2, 8, 7) gains 1 electron to achieve Argon's stability, becoming Cl⁻.
  • The oppositely charged ions attract to form NaCl.
• Calcium Oxide (CaO):
  • Calcium (2, 8, 8, 2) loses 2 electrons to become Ca²⁺.
  • Oxygen (2, 6) gains 2 electrons to become O²⁻.
• Magnesium Chloride (MgCl₂):
  • Magnesium (2, 8, 2) loses 2 electrons (becoming Mg²⁺).
  • Since one Chlorine atom only needs 1 electron, two Chlorine atoms each accept one electron to form two Cl⁻ ions.

E. Covalent Bonding

• Formation: Formed by the mutual sharing of electron pairs between two generally non-metallic atoms.
• Types of Covalent Bonds:
  • Single Bond (-): Sharing of one electron pair.
  • Double Bond (=): Sharing of two electron pairs.
  • Triple Bond (≡): Sharing of three electron pairs.
• Covalency: The number of electron pairs an atom shares to achieve a stable electronic configuration.
• Non-Polar vs. Polar Covalent Compounds:
  • Non-Polar: The shared pair of electrons is equally distributed (e.g., H₂, Cl₂, O₂, N₂, CH₄). Molecules are symmetrical and electrically neutral.
  • Polar: The shared pair is unequally distributed because one atom strongly attracts electrons, creating a slight charge separation or poles (e.g., H₂O, NH₃, HCl).

F. Structure of Covalent Compounds

• Hydrogen (H₂): Shares one electron pair (single bond) to achieve a stable duplet.
• Chlorine (Cl₂): Shares one electron pair (single bond) to achieve a stable octet.
• Oxygen (O₂): Shares two electron pairs (double bond) for stable octets.
• Nitrogen (N₂): Shares three electron pairs (triple bond) for stable octets.
• Carbon Tetrachloride (CCl₄): Carbon shares its 4 valence electrons with 4 Chlorine atoms, forming 4 single bonds.
• Methane (CH₄): Carbon shares its 4 valence electrons with 4 Hydrogen atoms, forming 4 single bonds.
• Water (H₂O) - Polar: Oxygen shares single bonds with two Hydrogens, leaving two lone pairs of unshared electrons.
• Ammonia (NH₃) - Polar: Nitrogen shares single bonds with three Hydrogens, leaving one lone pair of unshared electrons.

G. Coordinate Bond

• Coordinate Bond Definition: A special type of covalent bonding where the shared pair of electrons is contributed by only one of the combining atoms.
• Lone Pair: A pair of electrons in the valence shell that is not shared with any other atom (e.g., the oxygen in water has two, the nitrogen in ammonia has one).
• Formation of Hydronium Ion (H₃O⁺): A water molecule uses one of oxygen's lone pairs to bond with a hydrogen ion (H⁺), creating H₃O⁺.
• Formation of Ammonium Ion (NH₄⁺): An ammonia molecule uses nitrogen's lone pair to bond with a hydrogen ion (H⁺) from an acid or water, creating NH₄⁺.

H. Properties & Comparison