Chemistry 20 · Chemical bonding
Why substances behave differently
Salt is a hard crystal that melts at 801 °C and water is a liquid you can drink. Both are made of a metal-or-not and a non-metal; the difference is entirely in how the electrons are arranged. That is what this unit is for.
- 1. The words, first
- 2. Ionic, polar and non-polar
- 3. Lewis structures and shape
- 4. Forces between molecules
- 5. What costs marks
The words, first
The idea: Bond type is a continuum, not three boxes — and the force inside a molecule is a different thing from the force between molecules.
| Word | What it means |
|---|---|
| Ionic bond | The attraction between oppositely charged ions after electrons have transferred. Forms a lattice, not molecules. |
| Covalent bond | A shared pair of electrons between two atoms, both usually non-metals. |
| Electronegativity | How strongly an atom pulls on a shared pair. Increases up and to the right of the periodic table; fluorine is the highest. |
| Polar covalent bond | A shared pair held unequally, so one end is slightly negative and the other slightly positive. |
| Lewis structure | A drawing showing valence electrons as dots and bonding pairs as lines. |
| VSEPR | Valence Shell Electron Pair Repulsion — electron groups arrange themselves as far apart as possible, and that fixes the shape. |
| Lone pair | A valence pair not involved in bonding. It takes up space and pushes bonds together, but it is not part of the shape's name. |
| Dipole | A separation of charge across a bond or a whole molecule. |
| Intermolecular force | An attraction between separate molecules — always far weaker than the covalent bonds inside them. |
| Hydrogen bond | The strongest kind of dipole attraction, occurring when hydrogen is bonded to N, O or F. |
| London dispersion force | A weak attraction from momentary uneven electron distribution. Present in everything, and the only force between non-polar molecules. |
Ionic, polar and non-polar
The idea: Subtract the two electronegativities. A small difference means sharing, a large one means transfer, and most real bonds are somewhere in between.
Why ionic compounds behave as they do. There are no NaCl molecules. Each Na⁺ is surrounded by Cl⁻ ions and vice versa, throughout the crystal. Melting it means overcoming all of those attractions at once, which is why the melting point is 801 °C. Molten or dissolved, the ions are free to move, which is why the liquid conducts and the solid does not.
Why molecular substances behave as they do. Methane melts at −182 °C not because its C–H bonds are weak — they are strong — but because melting only has to separate whole molecules from each other, and the forces between them are feeble.
Lewis structures and shape
The idea: Count the electron groups around the central atom. Their arrangement is fixed by repulsion; the shape's name comes from where the atoms are, ignoring the lone pairs.
Drawing a Lewis structure. Count all valence electrons, join the atoms with single bonds, then distribute the rest as lone pairs to complete octets. If something is short, make a double or triple bond.
CO₂. Carbon has 4 valence electrons and each oxygen needs 2 more, so carbon double-bonds to each oxygen. Two electron groups on carbon, no lone pairs → 180° apart → linear.
4 groups, 0 lone pairs → tetrahedral (CH₄) · 4 groups, 1 lone pair → trigonal pyramidal (NH₃) · 4 groups, 2 lone pairs → bent (H₂O) · 3 groups, 0 lone pairs → trigonal planar · 2 groups → linear
Ammonia and methane have the same electron arrangement and different shapes, because a shape is named for atoms. That distinction is tested constantly.
Shape decides polarity. CO₂ and H₂O both have polar bonds. Carbon dioxide is linear, so the two dipoles point in exactly opposite directions and cancel — the molecule is non-polar. Water is bent, so they add, and the molecule has a net dipole. CCl₄ is the same story: four polar bonds arranged tetrahedrally cancel completely.
Forces between molecules
The idea: Boiling means separating whole molecules, so boiling point measures the forces between them — not the bonds inside them.
Weakest → strongest: London dispersion < dipole–dipole < hydrogen bonding < ionic lattice
Worked comparison. CH₄ (−162 °C) < H₂S (−60 °C) < H₂O (100 °C). Methane is non-polar, so only dispersion forces hold it; hydrogen sulfide is polar; water hydrogen bonds. All three are similar in size, so the difference is entirely the type of force.
The HF puzzle. HF boils at 20 °C and HCl at −85 °C, even though HCl is heavier. Mass predicts the wrong order because fluorine is electronegative enough for HF to hydrogen bond and chlorine is not.
Like dissolves like. Polar solvents dissolve polar and ionic solutes, because the solvent can form comparable attractions to the solute. Non-polar solvents dissolve non-polar solutes. Oil and water separate because neither can offer water anything as good as another water molecule.
What costs marks
The idea: Almost all of them come from mixing up two levels: inside a molecule versus between molecules.
- Saying covalent bonds break when something boils. They do not. Boiling separates molecules; the molecules survive.
- Calling a molecule polar because it has polar bonds. Check the shape first — symmetry can cancel them.
- Naming a shape from the electron arrangement. NH₃ has a tetrahedral arrangement and a pyramidal shape.
- Describing an ionic compound as molecules. NaCl is a formula unit — the simplest ratio in a lattice.