2.1 · Shapes & polarity

VSEPR: counting electron domains

Electron domains repel; lone pairs repel most.

By the end you should be able to:

Predict electron-domain and molecular geometry with VSEPR
Give AXE notation, lone pairs on the central atom and approximate bond angles

Interactive

VSEPR shape explorer

Rotate 3-D models, add lone pairs and bonding domains, and see how electron geometry becomes molecular geometry.

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Key idea

Electron domains and steric number

VSEPR (valence-shell electron-pair repulsion): the electron domains around a central atom spread out to be as far apart as possible.

An electron domain is any region of electron density on the central atom:

a lone pair;
a single bond;
a double or triple bond, which counts as one domain;
a single unpaired electron (in radicals such as $\ce{NO2}$ ).

The steric number SN = (atoms bonded to the centre) + (lone pairs on the centre). SN fixes the electron-domain geometry; the molecular geometry describes only where the atoms are.

Method

VSEPR in five steps

Draw the Lewis structure. Any valid resonance form gives the same shape.
Count the domains on the central atom: bonded atoms (X) and lone pairs (E). Each multiple bond counts once.
Write the AXE class, e.g. $\ce{NH3}$ is $\mathrm{AX_3E}$ and $\ce{H2O}$ is $\mathrm{AX_2E_2}$ .
SN gives the electron-domain geometry: 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral.
Name the molecular geometry from the atom positions only, then adjust the ideal angles for lone pairs.

Key idea

Classes with 2 to 4 electron domains

SN	AXE	Electron geometry	Molecular geometry	Angle	Example
2	$\mathrm{AX_2}$	linear	linear	180°	$\ce{CO2}$ , $\ce{BeCl2}$
3	$\mathrm{AX_3}$	trigonal planar	trigonal planar	120°	$\ce{BF3}$ , $\ce{SO3}$
3	$\mathrm{AX_2E}$	trigonal planar	bent	< 120°	$\ce{SO2}$ , $\ce{O3}$
4	$\mathrm{AX_4}$	tetrahedral	tetrahedral	109.5°	$\ce{CH4}$ , $\ce{NH4+}$
4	$\mathrm{AX_3E}$	tetrahedral	trigonal pyramidal	≈ 107°	$\ce{NH3}$
4	$\mathrm{AX_2E_2}$	tetrahedral	bent	≈ 104.5°	$\ce{H2O}$

Key idea

Classes with 5 and 6 electron domains

SN	AXE	Electron geometry	Molecular geometry	Angles	Example
5	$\mathrm{AX_5}$	trigonal bipyramidal	trigonal bipyramidal	90°, 120°, 180°	$\ce{PCl5}$
5	$\mathrm{AX_4E}$	trigonal bipyramidal	seesaw	< 90°, < 120°	$\ce{SF4}$
5	$\mathrm{AX_3E_2}$	trigonal bipyramidal	T-shaped	< 90°	$\ce{ClF3}$
5	$\mathrm{AX_2E_3}$	trigonal bipyramidal	linear	180°	$\ce{XeF2}$ , $\ce{I3-}$
6	$\mathrm{AX_6}$	octahedral	octahedral	90°, 180°	$\ce{SF6}$
6	$\mathrm{AX_5E}$	octahedral	square pyramidal	< 90°	$\ce{BrF5}$
6	$\mathrm{AX_4E_2}$	octahedral	square planar	90°	$\ce{XeF4}$

In a trigonal bipyramid, lone pairs always take equatorial positions (two 90° neighbours instead of three). In an octahedron, two lone pairs sit opposite each other.

Key idea

Lone pairs squeeze bond angles

Repulsion strength: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair. A lone pair is held by only one nucleus, so it spreads out and pushes the bonds closer together.

Molecule	Class	Lone pairs	H–X–H angle
$\ce{CH4}$	$\mathrm{AX_4}$	0	109.5°
$\ce{NH3}$	$\mathrm{AX_3E}$	1	107°
$\ce{H2O}$	$\mathrm{AX_2E_2}$	2	104.5°

Multiple bonds hold more electron density and also squeeze nearby angles slightly: in $\ce{CH2O}$ the H–C–H angle is about 116°, below the ideal 120°.

Common mistake

VSEPR traps

Wrong: calling $\ce{H2O}$ tetrahedral. Right: its electron-domain geometry is tetrahedral; its molecular geometry is bent.

Wrong: counting each double bond in $\ce{CO2}$ as two domains. Right: a multiple bond is one domain, so $\ce{CO2}$ is $\mathrm{AX_2}$ , linear.

Wrong: putting the lone pair of $\ce{SF4}$ in an axial position. Right: lone pairs in a trigonal bipyramid sit equatorial, giving a seesaw; with two or three lone pairs the shape is T-shaped ( $\ce{ClF3}$ ) or linear ( $\ce{XeF2}$ ).

Wrong: calling $\ce{XeF4}$ tetrahedral because it has four atoms attached. Right: Xe also has two lone pairs, so it is $\mathrm{AX_4E_2}$ , square planar.

Worked example

Worked example: SF₄ and XeF₄

Sulfur tetrafluoride. 34 valence e⁻; four S–F bonds and one lone pair on S, so SN = 5 and the class is $\mathrm{AX_4E}$ . Electron geometry trigonal bipyramidal; the lone pair is equatorial, so the molecule is a seesaw with angles a little under 90° and 120°.

Xenon tetrafluoride. 36 valence e⁻; four Xe–F bonds and two lone pairs on Xe, so SN = 6 and the class is $\mathrm{AX_4E_2}$ . Electron geometry octahedral; the lone pairs sit opposite each other, so the molecule is square planar with F–Xe–F angles of 90°.

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