Chemical Bonding
Easy Overview
Why do atoms stick together? Why is water wet and salt salty? This chapter is all about the glue that holds matter together. You'll learn how atoms share or transfer electrons, why some bonds are strong and others weak, and how to predict the shape of molecules. It's like relationship counseling for atoms.
Ionic Bond — The Generous Handover
Some atoms really want to give away electrons (metals), and others really want to take them (non-metals). When they meet, the metal donates an electron to the non-metal, becoming a positive ion, and the non-metal becomes a negative ion. Opposite charges attract — that's an ionic bond. Sodium and chlorine do this to make NaCl.
Covalent Bond — Sharing Is Caring
When two non-metals meet, neither wants to give up electrons. So they share. Each shared pair of electrons forms a covalent bond. H₂, O₂, N₂ — all held together by shared electrons. Single bond is one shared pair, double bond is two, triple bond is three. The more you share, the stronger the bond.
VSEPR Theory — Shape of Molecules
Electron pairs repel each other like magnets of the same pole. So they arrange themselves as far apart as possible. That's VSEPR. Two pairs? Linear (180°). Three pairs? Trigonal planar (120°). Four pairs? Tetrahedral (109.5°). The shape of a molecule determines a lot about its properties — including whether it's polar or not.
Valence Bond Theory and Hybridization
VBT says bonds form when atomic orbitals overlap. But carbon's orbitals aren't arranged right for making 4 equal bonds — so they 'mix' to form hybrid orbitals. sp³ hybridization gives methane its tetrahedral shape. sp² gives ethene's trigonal planar geometry. sp gives ethyne's linear structure. Think of it like mixing paint colors to get a new shade.
Molecular Orbital Theory
MOT says when two atoms bond, their atomic orbitals combine to form molecular orbitals — one lower energy (bonding) and one higher energy (antibonding). Electrons fill these like they fill atomic orbitals. If more electrons are in bonding orbitals than antibonding, the molecule is stable. This explains why He₂ doesn't exist but O₂ is paramagnetic.
Polarity and Intermolecular Forces
Some molecules have positive and negative ends — like a battery. That's polarity. Water is polar, oil isn't. Polar molecules stick to each other through dipole-dipole forces. Hydrogen bonding is a super strong version of this. London forces are weak attractions between all molecules. These forces determine melting points, boiling points, and solubility.
Key Points
- •Ionic bond: transfer of electrons between metal and non-metal
- •Covalent bond: sharing of electrons between non-metals
- •Lewis structures show valence electrons as dots around atomic symbols
- •VSEPR: electron pairs repel and arrange for maximum separation
- •Hybridization: sp³ (4 bonds, tetrahedral), sp² (3 bonds, trigonal planar), sp (2 bonds, linear)
- •MOT: bonding orbitals are lower energy, antibonding are higher; bond order = (bonding - antibonding)/2
- •Polar bonds arise from electronegativity differences between atoms
- •Hydrogen bonding is the strongest intermolecular force and explains water's high boiling point
Practice Questions
- Draw Lewis structures for H₂O, NH₃, and CO₂.
- Explain the formation of H₂ molecule using MOT. Calculate its bond order.
- What is hybridization? Predict the shape of CH₄, C₂H₄, and C₂H₂ using hybridization.
- Differentiate between ionic and covalent bonds with two examples each.
- Why does ice float on water? Explain using hydrogen bonding.