Covalent Bonding

Introduction to Covalent Bonding

Covalent Bonding: Involves the sharing of electron pairs between atoms, enabling them to achieve complete outer electron shells and attain molecular stability.

• Shared Electron Pairs: These are pairs of electrons shared between two atoms.
• Molecular Stability: This is achieved when atoms gain complete outer electron shells through electron sharing.

Forces in Covalent Bonding

• Attractive Forces: These forces act between the shared electron pairs and the nuclei of the bonded atoms, similar to how magnets are drawn together.
• Repulsive Forces: These occur between like-charged components, such as atomic nuclei, resembling the repulsion experienced when similar poles of magnets face each other.

Illustrating Covalent Bonds

Lewis Structures

• Lewis structures visually represent how atoms are bonded in a molecule, utilising dots to depict electrons.
• Steps to Draw Lewis Structures:
  • Count the total valence electrons.
  • Use symbols for atoms to connect them with single bonds.
  • Distribute any remaining electrons to fulfil octets or duets.

Example: Water (H₂O) and Methane (CH₄)

Electronegativity and Gradient of Bonds

What is Electronegativity?

Electronegativity: Refers to an atom's capacity to attract and retain electrons, akin to a tug-of-war where the stronger side draws electrons towards itself.

• Electronegativity Differences:
  • Less than 0.4: Nonpolar covalent (equal electron sharing)
  • 0.4 to 2.0: Polar covalent (unequal electron sharing)
  • Greater than 2.0: Ionic bonds (electrons are fully transferred)

Bond Strength and Length

Variability Based on Shared Electrons

• Bond Types:
  • Single Bonds: One pair of electrons shared; they are longer and weaker.
  • Double Bonds: Two pairs of electrons shared; they are shorter and stronger.
  • Triple Bonds: Three pairs of electrons shared; these are the shortest and strongest.

Common Misconceptions

Clarifying Misunderstandings

• Misconception: Electrons in covalent bonds are exclusively owned by the atoms.
• Reality: Electrons are shared, and are not solely held by any one atom in a covalent bond.

Understanding Bond Polarity

• Bond Polarity: Describes the unequal distribution of electronic charge around atoms within a molecule due to differences in electronegativity.

• Electronegativity Differences:

  • Atoms vary in their capacity to attract electrons.
  • Bonds are categorised based on the distribution of shared electrons.

• Polar vs. Nonpolar Bonds:

  • Polar bonds arise due to notable electronegativity differences.
  • Nonpolar bonds occur when the differences are minimal, resulting in equal sharing.

Periodic Trends and Their Impact

• Trends in Electronegativity:
  • An increase across periods as atoms contain more protons, drawing electrons closer.
  • A decrease down groups resulting from additional electron shells.

Comparison Examples

• HCl (Hydrogen Chloride): Considered polar due to chlorine's high electronegativity.
• CH₄ (Methane): Nonpolar, as carbon and hydrogen exhibit nearly equal electronegativity.

Worked Examples

Example with NH₃:

• Let's determine the polarity of the N-H bond in ammonia:
  • Find the electronegativity values: Nitrogen (3.04) and Hydrogen (2.20)
  • Calculate the difference: $3.04 - 2.20 = 0.84$
  • Since 0.84 is greater than 0.5, the N-H bond is classified as polar covalent
  • This explains why ammonia is soluble in water and has a dipole moment

Example with CO₂:

• For carbon dioxide:
  • Electronegativity values: Carbon (2.55), Oxygen (3.44)
  • Difference: $3.44 - 2.55 = 0.89$ (polar bonds)
  • However, due to CO₂'s linear structure, the polar bonds cancel each other out
  • Therefore, CO₂ is a nonpolar molecule despite having polar bonds

Differences Between Ionic and Covalent Compounds

1. Physical Properties Comparison

• Melting and Boiling Points: Ionic compounds generally exhibit higher melting and boiling points compared to covalent compounds.
• Solubility: Ionic compounds dissolve in polar solvents like water, while covalent compounds are more soluble in non-polar solvents like oil.
• Electrical Conductivity: Ionic compounds conduct electricity when they are melted or dissolved in solutions.

2. Structural Differences

• Ionic Compounds: Exist in a lattice structure, such as NaCl.
• Covalent Compounds: Feature discrete molecules, exemplified by H₂O.

Covalent Network Structures

Overview and Characteristics

Covalent Network Structures: Involve atoms bonded in a continuous framework through covalent bonds.

• Attributes:
  • Extreme Hardness: Resulting from the strength of the covalent bonds.
  • High Melting/Boiling Points: A substantial amount of energy is required to disrupt the network.

Diamond Structure Analysis:

• Carbon Lattice Explanation: Carbon atoms in an sp³ hybridised state form a three-dimensional tetrahedral network.
• Properties:
  • Hardness: Utilised in both jewellery and cutting tools.

Comparison with Molecular Covalent Compounds:

• Covalent Network: Robust, forming solid structures.
• Molecular Covalent: Maintained by weaker forces.