Molecular Shapes and Electron Pair Repulsion Theory

Ever wondered why water molecules bend whilst carbon dioxide stays straight? It's all about electron pair repulsion theory - electrons hate being near each other because they're all negatively charged.

There are two types of electron pairs around atoms: bonding pairs (BP) that connect atoms together, and lone pairs (LP) that just hang around the central atom. Lone pairs are bullies - they're held closer to the atom and push harder against other electron pairs than bonding pairs do.

Here are the key molecular shapes you need to know:

• Linear molecules like CO₂ have 180° bond angles (2 BP, 0 LP)
• Trigonal planar like BF₃ have 120° angles (3 BP, 0 LP)
• Tetrahedral like CH₄ have 109.5° angles (4 BP, 0 LP)
• Bent/non-linear like H₂O have 104.5° angles (2 BP, 2 LP)
• Pyramidal like NH₃ have 107° angles (3 BP, 1 LP)

Quick Tip: The repulsion strength order is LP-LP > BP-LP > BP-BP. This explains why bond angles get smaller when lone pairs are present!

When drawing 3D molecular shapes, remember that wedged lines (▲) come towards you, dashed lines (---) go away from you, and straight lines stay in the plane of the paper.

Electronegativity and Bond Types

Electronegativity is basically how greedy an atom is for electrons - it measures an atom's ability to attract bonding electron pairs. The Pauling scale gives us numerical values to compare different elements.

Electronegativity increases as you go across periods (left to right) but decreases as you go down groups. This happens because nuclear charge increases across periods, whilst atomic size increases down groups.

The electronegativity difference between bonded atoms determines the bond type:

• Non-polar covalent bonds $difference 0-0.4$: electrons shared equally, like in hydrocarbons
• Polar covalent bonds $difference 0.4-1.8$: unequal electron sharing creates dipoles
• Ionic bonds (difference >1.8): complete electron transfer

Remember: Dipoles are just separations of charge, marked with δ+ and δ- symbols to show which end is slightly positive or negative.