Electron Shells and Bonding Basics

Think of electron shells as floors in a building - each one can hold a specific number of residents (electrons). Shell 1 holds just 2 electrons, shell 2 can fit 8, shell 3 holds 18, and shell 4 can accommodate 32. As you go higher up the building (higher shell numbers), the energy levels increase too.

Within each shell, you'll find different types of orbitals - think of these as individual rooms that can hold up to 2 electrons each. Every shell has an s-orbital, shells 2 and above have p-orbitals (3 of them), shells 3+ have d-orbitals (5 of them), and shells 4+ have f-orbitals (7 of them).

When filling orbitals, electrons are like people choosing seats on a bus - they'll sit alone first before sharing. They fill up singly before pairing, and paired electrons must spin in opposite directions. The filling order follows a specific pattern: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p.

Quick Tip: Remember that electrons are lazy - they always fill the lowest energy orbitals first!

Ionic and Covalent Bonding

Ionic bonds form when metals give electrons to non-metals, creating a powerful electrostatic attraction between positive and negative ions. Group 1 metals lose 1 electron (forming +1 ions), Group 2 lose 2 electrons $+2 ions$, whilst Group 6 non-metals gain 2 electrons $-2 ions$ and Group 7 gain 1 electron $-1 ions$.

Ionic compounds form giant ionic lattices where each ion is surrounded by oppositely charged neighbours. This gives them high melting points, makes them soluble in water, and means they conduct electricity when molten or dissolved (but not as solids).

Covalent bonding happens between non-metals that share electrons rather than transferring them. A single covalent bond represents one shared pair of electrons, shown as a line between atoms (like Cl-Cl). Double and triple bonds involve sharing 2 or 3 pairs of electrons respectively.

Co-ordinate bonds are special covalent bonds where both electrons come from just one atom - the donor atom gives a lone pair to an electron-deficient acceptor. Despite their different formation, they're just as strong as regular covalent bonds.

Remember: Metals give, non-metals take (ionic), but non-metals share with each other (covalent)!

Polarity and Intermolecular Forces

Electronegativity measures how greedily an atom pulls electron density towards itself in a covalent bond. It increases across periods and decreases down groups, with fluorine being the most electronegative at 4.0 on the Pauling scale.

When atoms with different electronegativities bond, you get polar bonds with partial charges $δ+ and δ-$. However, a molecule can have polar bonds but still be non-polar overall if the electron cloud is evenly distributed across the whole structure.

Intermolecular forces are the weak attractions between molecules that determine physical properties. Van der Waals forces exist between all molecules and get stronger with molecular size - this is why larger molecules have higher boiling points. Branched molecules have weaker forces than straight chains due to less surface contact.

Permanent dipole-dipole forces occur between polar molecules, whilst hydrogen bonding is the strongest intermolecular force, forming when hydrogen bonds to oxygen, nitrogen, or fluorine.

Key Insight: Intermolecular forces are much weaker than actual bonds, but they control whether something is a solid, liquid, or gas at room temperature!

Molecular Shapes and Material Structures

Molecular shapes depend on electron pair repulsion - electrons push apart to get as far from each other as possible. Lone pairs repel more strongly than bonding pairs, affecting the final shape. You can represent 3D structures using solid lines (in the plane), solid wedges (coming out), and dotted wedges (going in).

Different structure types explain material properties. Giant ionic structures have high melting points and conduct when molten. Simple covalent structures have low melting points due to weak intermolecular forces. Giant covalent structures like diamond and graphite have very high melting points.

Diamond is incredibly hard because every carbon atom bonds to four others in a rigid 3D network, making it perfect for cutting tools. Graphite has layers that slide over each other (making it soft and useful for pencils) plus delocalised electrons that conduct electricity (ideal for electrodes).

The arrangement and movement of particles differs dramatically between solids (regular, vibrating), liquids (random, jostling), and gases (random, rapid movement), explaining their different properties like compressibility and diffusion rates.

Amazing Fact: Diamond and graphite are both pure carbon - their completely different properties come entirely from how the atoms are arranged and bonded!