# Molecular Shapes and Polarity

> Drawing Lewis structures step by step, predicting shapes with VSEPR, and telling polar from non-polar bonds and molecules.

From HSC Chemistry on WhizzNotes, free study notes for NSW students: https://whizznotes.com/chemistry/molecular-shapes-and-polarity/. Updated 3 November 2026.
Year 11 and Year 12 boxes go deeper than the Year 10 text around them.

### Story time: Why oil and water fight
Hana made salad dressing by pouring olive oil into vinegar and shaking it hard. For a moment it looked mixed. Then, while she set the table, the oil crept back up to the top as a separate layer.

Her grandfather said, "Oil and water are like two groups of friends who only like their own kind."

He was more right than he knew. Water molecules are **bent**, with a slightly negative end and a slightly positive end. They cling to each other. Oil molecules have no charged ends at all. The shape of a molecule decides whether it has charged ends, and that decides what it mixes with, how high it boils, and even how it acts in your body.

**In short:** VSEPR theory predicts a molecule's shape: the electron groups around the central atom spread out as far apart as possible, giving shapes like linear, tetrahedral and bent. A molecule is polar if it has polar bonds and is not symmetrical.

## Drawing Lewis structures step by step

A **Lewis structure** (electron dot structure) shows every valence electron in a molecule. Use this method for any small molecule.

1. **Count** the total valence electrons from all the atoms.
2. Pick the **central atom**: usually the one there is only one of, or the least electronegative. **Hydrogen and the halogens are always on the outside**, because they make only one bond.
3. Join each outer atom to the centre with **one single bond** (a line = 2 electrons).
4. Put the remaining electrons on the **outer** atoms as lone pairs until each has 8 (hydrogen only needs 2).
5. Put any electrons left over on the **central** atom.
6. If the central atom has fewer than 8, **move a lone pair** from an outer atom to make a **double or triple bond**.

### Worked example: Ammonia, NH₃
1. Valence electrons: N has 5, each H has 1. Total $5 + 3 = 8$.
2. N is central.
3. Three N−H bonds use 6 electrons.
4. H atoms are full with 2 each.
5. 2 electrons left go on N as **one lone pair**.
6. N has 3 bonds (6) + 1 lone pair (2) = 8. Done.

### Worked example: Carbon dioxide, CO₂
1. Valence electrons: C has 4, each O has 6. Total $4 + 12 = 16$.
2. C is central: O−C−O uses 4 electrons.
3. 12 left: 6 on each O (three lone pairs each). 0 left.
4. C has only 4 electrons. Move one lone pair from each O into a bond.

Result: O=C=O, with two lone pairs on each oxygen. Every atom now has 8.

## VSEPR: predicting shapes

**VSEPR** stands for **Valence Shell Electron Pair Repulsion**. The idea is simple: electron pairs are all negative, so they **push each other as far apart as possible** around the central atom.

To find a shape:

1. Draw the Lewis structure.
2. Count the **electron groups** around the central atom. A single, double or triple bond each counts as **one** group. Each lone pair counts as one group.
3. Use the table.

| Bonded atoms | Lone pairs on centre | Shape | Bond angle | Example |
|---|---|---|---|---|
| 2 | 0 | **Linear** | 180° | $\mathrm{CO_2}$, $\mathrm{BeCl_2}$ |
| 3 | 0 | **Trigonal planar** | 120° | $\mathrm{BF_3}$ |
| 4 | 0 | **Tetrahedral** | 109.5° | $\mathrm{CH_4}$ |
| 3 | 1 | **Trigonal pyramidal** | about 107° | $\mathrm{NH_3}$ |
| 2 | 2 | **Bent** (V-shaped) | about 104.5° | $\mathrm{H_2O}$ |

*Diagram: Three shapes: linear carbon dioxide, tetrahedral methane drawn with a wedge and dash, and bent water with two lone pairs*
*A wedge points out of the page, a dashed line points into it. Water's two lone pairs (red) push the hydrogens closer together.*

### Key idea: Lone pairs push harder
A lone pair is held by only one nucleus, so it spreads out more and pushes harder than a bonding pair. Each lone pair squeezes the bond angle by about 2 to 3°. That is why methane is 109.5°, ammonia about 107° and water about 104.5°, even though all three have 4 electron groups.

## Polar and non-polar bonds

**Electronegativity** is how strongly an atom pulls on shared electrons (see the periodic trends chapter). When two atoms share electrons, compare their electronegativities.

| Electronegativity difference | Bond type | Example |
|---|---|---|
| 0 to about 0.4 | **Non-polar covalent**: shared equally | H−H, C−H, Cl−Cl |
| about 0.4 to 1.8 | **Polar covalent**: shared unequally | O−H, H−Cl, C−O |
| more than about 1.8 | **Ionic**: electrons transferred | Na−Cl, Mg−O |

In a polar bond, the more electronegative atom gets a **partial negative charge** ($\delta-$) and the other a **partial positive charge** ($\delta+$). The Greek letter $\delta$ (delta) means "a little bit of".

### Worked example: Which bond is it?
Electronegativities: H 2.1, Cl 3.0, Na 0.9.

- H−Cl: $3.0 - 2.1 = 0.9$, so **polar covalent**, with $\delta-$ on Cl.
- Na−Cl: $3.0 - 0.9 = 2.1$, so **ionic**.
- Cl−Cl: $3.0 - 3.0 = 0$, so **non-polar covalent**.

## Polar and non-polar molecules

A molecule with polar bonds is not always polar overall. **Shape decides.**

- If the polar bonds are arranged **symmetrically**, their pulls **cancel** and the molecule is **non-polar**. $\mathrm{CO_2}$ (linear) and $\mathrm{CCl_4}$ (tetrahedral) have polar bonds but are non-polar.
- If the shape is **not symmetrical**, the pulls do not cancel and the molecule is **polar**. $\mathrm{H_2O}$ (bent), $\mathrm{NH_3}$ (pyramidal) and HCl are polar.

Think of a tug of war. If equal teams pull in exactly opposite directions, nothing moves. If the teams pull at an angle, the rope moves.

### Exam tip: Like dissolves like
Polar substances dissolve in polar solvents (salt and sugar in water). Non-polar substances dissolve in non-polar solvents (grease in petrol). This is why water alone cannot wash oil off your hands, and why soap, which has a polar end and a non-polar end, can.

### Watch out: Polar bonds do not always make a polar molecule
$\mathrm{CO_2}$ has two very polar C=O bonds, but it is a **non-polar** molecule. Always check the shape before deciding.

**Mind map**

```text
Shapes and polarity
  Lewis structures
    Count valence electrons
    Central atom, single bonds
    Fill outer atoms to 8
    Make double bonds if short
  VSEPR
    Electron groups repel
    Linear 180°
    Trigonal planar 120°
    Tetrahedral 109.5°
    Pyramidal 107°, bent 104.5°
  Bonds
    Non-polar: equal sharing
    Polar: δ+ and δ−
    Ionic: big difference
  Molecules
    Symmetrical: non-polar
    Not symmetrical: polar
    Like dissolves like
```

### Year 11: Lewis structures of polyatomic ions
For an ion, adjust the electron count before you start:

- **Anion**: **add** one electron for each negative charge.
- **Cation**: **subtract** one electron for each positive charge.

Then draw it the normal way and put the whole structure in **square brackets** with the charge outside.

Hydroxide, $\mathrm{OH^-}$: $6 + 1 + 1 = 8$ electrons. One O−H bond and three lone pairs on O, written $[\mathrm{O{-}H}]^-$.

Ammonium, $\mathrm{NH_4^+}$: $5 + 4 - 1 = 8$ electrons. Four N−H bonds, no lone pairs, so it is **tetrahedral** at 109.5°.

### Year 12: Resonance and expanded octets
Sometimes more than one correct Lewis structure can be drawn. In the nitrate ion, $\mathrm{NO_3^-}$, the double bond could be to any of the three oxygens. The real ion is a blend of all three, called **resonance**. Measurements show all three N−O bonds are the **same length**, between a single and a double bond.

Atoms in **period 3 and below** can have **more than 8** electrons around them, because they have empty d orbitals. Sulfur in $\mathrm{SF_6}$ has 12 (6 bonds, **octahedral**, 90°). Phosphorus in $\mathrm{PCl_5}$ has 10 (**trigonal bipyramidal**, 90° and 120°). Period 2 atoms such as C, N and O can never go past 8.

## Quick summary

- Lewis structures: count valence electrons, join atoms with single bonds, fill outer atoms to 8, then make double or triple bonds if the centre is short.
- **VSEPR**: electron groups around the central atom repel and spread out as far as possible.
- Shapes: linear (180°), trigonal planar (120°), tetrahedral (109.5°), trigonal pyramidal (about 107°), bent (about 104.5°).
- Lone pairs push harder and squeeze bond angles.
- Electronegativity difference: small means non-polar, medium means polar, large means ionic.
- Symmetrical molecules are non-polar even with polar bonds. Unsymmetrical ones are polar.

## Quiz

Try these questions from the chapter quiz.

1. How many valence electrons are in a molecule of $\mathrm{H_2O}$?
   - A. 6
   - B. 8
   - C. 10
   - D. 18
   - Answer: B. 8. O has 6 and each H has 1: 6 + 2 = 8.

2. In a Lewis structure, which atom is never the central atom?
   - A. Carbon
   - B. Nitrogen
   - C. Hydrogen
   - D. Oxygen
   - Answer: C. Hydrogen. Hydrogen can only form one bond, so it is always on the outside.

3. What does VSEPR theory say?
   - A. Electron pairs around a central atom repel and spread out as far as possible
   - B. Atoms always form ionic bonds
   - C. Molecules are always flat
   - D. Lone pairs attract each other
   - Answer: A. Electron pairs around a central atom repel and spread out as far as possible. Repulsion between electron pairs decides the shape.

The whole chapter quiz is on the page, for signed-in students: https://whizznotes.com/chemistry/molecular-shapes-and-polarity/#quiz