Chemical Bonding and Molecular Shape

Ionic Bonding

Ionic bonds form between metals and non-metals through electron transfer.

• Metals lose electrons → cations (e.g. Na⁺)
• Non-metals gain electrons → anions (e.g. Cl⁻)
• Strong electrostatic attraction between oppositely charged ions
• Forms giant ionic lattices

Covalent Bonding

Covalent bonds form between non-metals through sharing electron pairs.

• Each atom contributes one electron to the shared pair
• Bond strength depends on orbital overlap
• Can form single (1 pair), double (2 pairs), or triple (3 pairs) bonds

Dative (Coordinate) Bonding

A special type of covalent bond where one atom provides both electrons in the shared pair.

• The atom donating the pair must have a lone pair
• Represented by an arrow (→) from donor to acceptor

Examples:

• $\text{NH}_4^+$: NH₃ donates a lone pair to H⁺
• $\text{H}_3\text{O}^+$: H₂O donates a lone pair to H⁺
• $\text{Al}_2\text{Cl}_6$: Cl donates a lone pair to Al
• CO: carbon to oxygen dative bond (triple bond with one dative component)

Metallic Bonding

Positive metal ions in a regular lattice surrounded by a sea of delocalised electrons. Strong electrostatic attraction between ions and delocalised electrons.

VSEPR (Valence Shell Electron Pair Repulsion) theory states:

• Electron pairs around a central atom repel each other
• They arrange themselves to be as far apart as possible
• This determines the shape and bond angles of the molecule

Types of Electron Pairs

• Bonding pairs (bp): shared between atoms
• Lone pairs (lp): not involved in bonding

Lone pairs repel more than bonding pairs: $\text{lp-lp} > \text{lp-bp} > \text{bp-bp}$

Lone pairs occupy more space, so they compress bond angles.

Shapes and Bond Angles

Why Lone Pairs Reduce Bond Angles

• CH₄: 4 bp → 109.5°
• NH₃: 3 bp + 1 lp → 107° (lp compresses)
• H₂O: 2 bp + 2 lp → 104.5° (two lps compress further)

Electronegativity is the ability of an atom to attract the bonding pair of electrons in a covalent bond.

The Pauling scale measures electronegativity:

• Fluorine is the most electronegative element (4.0)
• Electronegativity increases across a period and up a group

Polar Bonds

When two atoms with different electronegativities form a covalent bond, the shared electrons are pulled towards the more electronegative atom:

• Creates a polar bond with partial charges: $\delta^+$ and $\delta^-$
• The greater the electronegativity difference, the more polar the bond
• If the difference is very large (>1.7), the bond becomes essentially ionic

Polar and Non-polar Molecules

A molecule can have polar bonds but be non-polar overall if the bond dipoles cancel out due to symmetry:

• CO₂: polar C=O bonds but linear → dipoles cancel → non-polar
• CCl₄: polar C-Cl bonds but tetrahedral → dipoles cancel → non-polar
• H₂O: polar O-H bonds and bent → dipoles don't cancel → polar
• CHCl₃: asymmetric → polar

4.1 London (Dispersion) Forces

• Present in all molecules
• Caused by temporary, instantaneous dipoles from random electron movement
• Temporary dipole induces a dipole in a neighbouring molecule
• Strength increases with number of electrons (larger molecules)
• The weakest type of intermolecular force

4.2 Permanent Dipole-Dipole Forces

• Between molecules with permanent dipoles (polar molecules)
• The $\delta^+$ end of one molecule attracts the $\delta^-$ end of another
• Stronger than London forces alone

4.3 Hydrogen Bonding

The strongest type of intermolecular force. Occurs when hydrogen is bonded to a highly electronegative atom (F, O, or N) and interacts with a lone pair on another F, O, or N atom.

Conditions for hydrogen bonding:

1. H bonded to F, O, or N (very polar bond)
2. The F, O, or N on the neighbouring molecule has a lone pair

Effects of hydrogen bonding:

• Water has an anomalously high boiling point for its molecular mass
• Ice is less dense than liquid water (hydrogen bonds create an open lattice)
• Ethanol is soluble in water (forms H-bonds with water)
• DNA double helix is held together by H-bonds between base pairs

• For VSEPR, count ALL electron pairs (bonding + lone pairs) around the central atom
• State the shape name AND the bond angle — both are usually required
• For polarity: consider both bond polarity AND molecular symmetry
• Hydrogen bonding requires H bonded to F, O, or N — not just any electronegative atom
• Remember: lone pairs compress bond angles

• Ionic bonding: electron transfer → lattice of ions
• Covalent bonding: electron sharing; dative: one atom provides both electrons
• VSEPR: electron pairs repel → molecular shape; lp > bp repulsion
• Electronegativity determines bond polarity; molecular shape determines molecular polarity
• Intermolecular forces: London < dipole-dipole < hydrogen bonding
• Hydrogen bonds form between H (on F/O/N) and lone pair on F/O/N