What does VSEPR theory say?

Electron pairs around a central atom repel and spread out as far as possible

Molecular Shapes and Polarity

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.

Count the total valence electrons from all the atoms.
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.
Join each outer atom to the centre with one single bond (a line = 2 electrons).
Put the remaining electrons on the outer atoms as lone pairs until each has 8 (hydrogen only needs 2).
Put any electrons left over on the central atom.
If the central atom has fewer than 8, move a lone pair from an outer atom to make a double or triple bond.

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:

Draw the Lewis structure.
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.
Use the table.

Bonded atoms	Lone pairs on centre	Shape	Bond angle	Example
2	0	Linear	180°	$C O_{2}$ , $BeC l_{2}$
3	0	Trigonal planar	120°	$B F_{3}$
4	0	Tetrahedral	109.5°	$C H_{4}$
3	1	Trigonal pyramidal	about 107°	$N H_{3}$
2	2	Bent (V-shaped)	about 104.5°	$H_{2} O$

A wedge points out of the page, a dashed line points into it. Water's two lone pairs (red) push the hydrogens closer together.

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 ( $δ -$ ) and the other a partial positive charge ( $δ +$ ). The Greek letter $δ$ (delta) means "a little bit of".

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. $C O_{2}$ (linear) and $CC l_{4}$ (tetrahedral) have polar bonds but are non-polar.
If the shape is not symmetrical, the pulls do not cancel and the molecule is polar. $H_{2} O$ (bent), $N H_{3}$ (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.

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

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

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How many valence electrons are in a molecule of $H_{2} O$ ?
1. A 6
2. B 8
3. C 10
4. D 18
Show answer
B: 8
O has 6 and each H has 1: 6 + 2 = 8.
In a Lewis structure, which atom is never the central atom?
1. A Carbon
2. B Nitrogen
3. C Hydrogen
4. D Oxygen
Show answer
C: Hydrogen
Hydrogen can only form one bond, so it is always on the outside.
What does VSEPR theory say?
1. A Electron pairs around a central atom repel and spread out as far as possible
2. B Atoms always form ionic bonds
3. C Molecules are always flat
4. D Lone pairs attract each other
Show answer
A: Electron pairs around a central atom repel and spread out as far as possible
Repulsion between electron pairs decides the shape.