Does a double bond count as two electron domains?

No. A double or triple bond counts as one electron domain. Both electron pairs in a double bond lie between the same two nuclei, so they point in one direction and behave as a single region of electron density. Treating a double bond as two domains is the single most common error in VSEPR problems and will change your predicted shape entirely.

What is the difference between electron geometry and molecular shape?

Electron geometry counts every domain, including lone pairs; molecular shape counts only the positions of atoms. They match when the central atom has no lone pairs. When lone pairs are present, the lone pairs still control the arrangement but are not part of the shape name — ammonia has tetrahedral electron geometry but trigonal pyramidal molecular shape.

Why are bond angles described as approximate?

The ideal angles come from spreading identical domains as far apart as possible. Real domains differ: a lone pair spreads wider than a bonding pair, and a multiple bond is bulkier than a single bond. Each lone pair pushes the bonds closer together, so the actual angle drops below the ideal — from 109.5^ to about 107^ in NH_3 and about 104.5^ in H_2O.

Do lone pairs on the outer atoms affect the shape?

No. Only domains on the central atom determine geometry. The three lone pairs on each chlorine in CCl_4 or on each fluorine in BF_3 are ignored. Draw them in your Lewis structure for correctness, but count only what is attached to the central atom when you apply VSEPR.

CHEM-4.4

Molecular Shapes: VSEPR Theory

Learn VSEPR theory: count electron domains, tell electron geometry from molecular shape, and predict bond angles for linear, bent, trigonal pyramidal, and tetrahedral molecules.

What you'll do in this lesson

A voice-first session with the Crimsora tutor on Molecular Shapes: VSEPR Theory, then targeted practice and FRQs — with the tutor adapting to where you get stuck.

Start Molecular Shapes: VSEPR Theory free

What this lesson covers

A Lewis structure tells you which atoms are bonded and where the lone pairs sit, but it is drawn flat on paper. Real molecules are three-dimensional, and their shape controls almost everything about how they behave — whether water is a liquid at room temperature, how enzymes recognize a drug molecule, why carbon dioxide is nonpolar even though it contains polar bonds.

VSEPR theory (valence shell electron pair repulsion) is the simple, powerful model that turns a flat Lewis structure into a 3-D shape. The whole model rests on one idea: negatively charged groups of electrons around a central atom push each other as far apart as possible. In this lesson you will learn to count electron domains, name the electron geometry, name the molecular shape (which counts only atoms), and estimate bond angles — including why lone pairs squeeze those angles smaller than the ideal values.

The One Rule Behind VSEPR: Electrons Repel

Every bond and every lone pair around a central atom is a region of negative charge. Like charges repel, so those regions arrange themselves to be as far apart as geometrically possible. That single rule is all VSEPR needs.

The regions being counted are called electron domains (also called electron groups). A domain is any of the following attached to the central atom: a single bond, a double bond, a triple bond, a lone pair, or a lone single electron. The crucial and most-missed detail is that a multiple bond counts as one domain, not two or three. The two electron pairs in a double bond are locked between the same two nuclei, so they point in one direction and repel other domains as a single fat region.

Because domains push apart, the number of domains alone determines the basic arrangement in space:

Domains	Electron geometry	Ideal angle
2	linear	$180^\circ$
3	trigonal planar	$120^\circ$
4	tetrahedral	$109.5^\circ$
5	trigonal bipyramidal	$90^\circ$ and $120^\circ$
6	octahedral	$90^\circ$

Notice that four domains do not spread out at

90^\circ

in a flat cross. Pushing one pair down out of the plane lets all four separate to

109.5^\circ

, which is farther apart. That is why tetrahedral geometry, not square, dominates carbon chemistry.

Counting Domains from a Lewis Structure

VSEPR always starts from a correct Lewis structure, so the counting procedure is mechanical once you have one.

First, identify the central atom — usually the single atom of one element, the least electronegative atom, or the atom written first in the formula (hydrogen is never central because it forms only one bond). Second, count the atoms bonded to that central atom; each bonded atom contributes exactly one domain regardless of whether the bond is single, double, or triple. Third, count the lone pairs drawn on the central atom. Lone pairs on outer atoms are ignored completely — they do not affect the shape.

A useful shorthand is

AX_mE_n

, where

A

is the central atom,

X

is a bonded atom, and

E

is a lone pair. Total domains

= m + n

Molecule	Bonded atoms ( $m$ )	Lone pairs on center ( $n$ )	Total domains
$CO_2$	2	0	2
$SO_2$	2	1	3
$CH_2O$	3	0	3
$NH_3$	3	1	4
$H_2O$	2	2	4

For an ion, adjust the electron count before drawing: add one electron per negative charge, subtract one per positive charge. In

NH_4^+

the nitrogen has four bonds and no lone pair, so it has four domains; in

H_3O^+

the oxygen has three bonds and one lone pair, also four domains. Getting the lone-pair count on the central atom right is the step that decides everything that follows.

Electron Geometry Versus Molecular Shape

These two terms are the source of most confusion in this topic. Electron geometry describes the arrangement of all domains, lone pairs included. Molecular shape (molecular geometry) describes the arrangement of the atoms only — you keep the same skeleton but you make the lone pairs invisible when you name the shape.

When there are no lone pairs on the central atom, the two names are identical. When lone pairs are present, they still occupy space and still set the arrangement, but since we cannot see them, the visible shape gets a different name.

Domains	Lone pairs	Electron geometry	Molecular shape	Example
2	0	linear	linear	$CO_2$
3	0	trigonal planar	trigonal planar	$BF_3$
3	1	trigonal planar	bent	$SO_2$
4	0	tetrahedral	tetrahedral	$CH_4$
4	1	tetrahedral	trigonal pyramidal	$NH_3$
4	2	tetrahedral	bent	$H_2O$

The classic error is calling ammonia "tetrahedral" or water "linear." Ammonia's four domains are arranged tetrahedrally, but you can only see three hydrogens and the nitrogen, and that set of four atoms looks like a tripod — trigonal pyramidal. Water has four domains arranged tetrahedrally, yet only three atoms are visible, giving a bent shape.

Note also that "bent" appears twice, from three domains and from four. Bent is a shape name, not a domain count, so always state the electron geometry too if you are asked to describe a molecule fully.

Bond Angles: Why Lone Pairs Squeeze

Ideal angles assume every domain repels equally. Real molecules deviate because domains are not all the same size.

A lone pair is held by only one nucleus, so its electron cloud spreads out wider and closer to the central atom than a bonding pair, which is pulled taut between two nuclei. The result is the repulsion ranking

\text{lone pair–lone pair} > \text{lone pair–bonding pair} > \text{bonding pair–bonding pair}

Each lone pair therefore pushes the bonding pairs together, shrinking the bond angle below the ideal value. The tetrahedral series shows this cleanly:

CH_4

has

109.5^\circ

NH_3

drops to about

107^\circ

with one lone pair, and

H_2O

drops to about

104.5^\circ

with two.

Molecule	Domains / lone pairs	Approximate angle
$CH_4$	4 / 0	$109.5^\circ$
$NH_3$	4 / 1	$107^\circ$
$H_2O$	4 / 2	$104.5^\circ$
$SO_2$	3 / 1	slightly less than $120^\circ$
$CO_2$	2 / 0	$180^\circ$

Multiple bonds are also slightly bulkier than single bonds, so in

CH_2O

the

H-C-H

angle is a little under

120^\circ

while the

H-C=O

angles open a little above it.

On homework, the safe phrasing is "approximately

107^\circ

" or "slightly less than

109.5^\circ

." What matters is that you justify the direction of the deviation with lone-pair repulsion — a bare number with no reasoning shows nothing about whether you understand the model.

A Reliable Procedure and Where Students Slip

Work every VSEPR problem in the same order and the shapes stop feeling like memorization.

Step	Action
1	Draw a valid Lewis structure, adjusting electrons for any ionic charge
2	Count domains on the central atom: bonded atoms plus lone pairs
3	Name the electron geometry from the domain count
4	Delete the lone pairs mentally and name the molecular shape
5	State the ideal angle, then reduce it slightly for each lone pair

The recurring slip-ups are worth naming. Counting a double bond as two domains turns trigonal planar molecules into tetrahedral ones; a multiple bond is always exactly one domain. Counting lone pairs on outer atoms inflates the domain count — the three lone pairs on each chlorine in

CCl_4

are irrelevant to the shape. Skipping the Lewis structure and guessing from the formula fails immediately:

CO_2

and

H_2O

are both

AX_2

-looking formulas but are linear and bent respectively, because oxygen in water carries two lone pairs. Finally, students often report the electron geometry when the question asked for molecular shape; read the wording carefully.

Shape matters beyond naming. In the next lesson on polarity, the shape decides whether individual bond dipoles cancel:

CO_2

is linear, so its two polar bonds point in exactly opposite directions and cancel, making the molecule nonpolar, while bent

H_2O

cannot cancel and is strongly polar. Same kinds of bonds, different geometry, completely different behavior — which is exactly why chemists care about VSEPR.

Key terms

VSEPR theory.: Valence shell electron pair repulsion theory: the model stating that electron domains around a central atom arrange themselves as far apart as possible to minimize repulsion.
Electron domain.: A region of electron density around the central atom — one lone pair, or one bond of any order. A double or triple bond counts as a single domain.
Electron geometry.: The three-dimensional arrangement of all electron domains around the central atom, counting both bonding domains and lone pairs.
Molecular shape.: The arrangement of the atoms only, obtained by ignoring the lone pairs on the central atom while keeping the underlying domain arrangement.
Bonding pair.: A pair of electrons shared between two atoms; it is pulled between two nuclei and occupies a narrower region than a lone pair.
Lone pair.: A nonbonding electron pair on the central atom; it spreads over a wider region and repels other domains more strongly, compressing bond angles.
Bond angle.: The angle formed at the central atom between two bonds to outer atoms, reported as an approximate value because lone pairs and multiple bonds cause deviations from ideal values.
AXE notation.: Shorthand $AX_mE_n$ where $A$ is the central atom, $m$ is the number of bonded atoms, and $n$ is the number of lone pairs; total domains equal $m+n$ .

Worked example

Predict the electron geometry, molecular shape, and approximate bond angle of the sulfite ion,

SO_3^{2-}

Step 1 — Lewis structure. Valence electrons: sulfur contributes 6, each oxygen contributes 6 for 18, and the

2-

charge adds 2 more, giving

6 + 18 + 2 = 26

electrons. Place sulfur in the center with three oxygens around it. Three

S-O

single bonds use 6 electrons, leaving 20. Completing octets on the three oxygens uses 18 of those, leaving 2 electrons, which go on sulfur as one lone pair. Every atom now has an octet, so the structure is valid.

Step 2 — Count domains on the central atom. Sulfur has 3 bonded oxygens plus 1 lone pair, so there are

3 + 1 = 4

electron domains. In AXE notation this is

AX_3E_1

.

Step 3 — Electron geometry. Four domains always spread to a tetrahedral arrangement, ideal angle

109.5^\circ

.

Step 4 — Molecular shape. Hide the lone pair and look at the atoms: sulfur sits at the apex with three oxygens forming a triangular base. That is trigonal pyramidal, the same shape as

NH_3

.

Step 5 — Bond angle. Start from the ideal

109.5^\circ

. One lone pair repels the bonding pairs more strongly than they repel each other, so the

O-S-O

angles compress to approximately

107^\circ

(any answer phrased as "slightly less than

109.5^\circ

" with that reasoning is complete).

Final answer: tetrahedral electron geometry, trigonal pyramidal molecular shape, bond angles of about

107^\circ

Practice questions

A central atom is surrounded by four electron domains, two of which are lone pairs. What is the molecular shape?

Linear
Bent
Trigonal planar
Trigonal pyramidal

Answer: Bent

Four domains give a tetrahedral electron geometry. Removing the two lone pairs from view leaves the central atom and only two bonded atoms, which sit at roughly

104.5^\circ

to each other — a bent shape, exactly like water. Trigonal pyramidal would require three bonded atoms and one lone pair, and linear would require the two bonding domains to be

180^\circ

apart, which only happens when there are no lone pairs crowding them.

Both

CO_2

and

H_2O

contain one central atom bonded to two other atoms, yet

CO_2

is linear with a

180^\circ

angle and

H_2O

is bent with an angle near

104.5^\circ

. Explain the difference using VSEPR.

Answer: Carbon in $CO_2$ has two domains (two double bonds, no lone pairs), so the domains spread to $180^\circ$ and the molecule is linear. Oxygen in $H_2O$ has four domains (two bonds plus two lone pairs), giving tetrahedral electron geometry; hiding the lone pairs leaves a bent shape, and the extra repulsion from two lone pairs compresses the angle to about $104.5^\circ$ .

The key move is counting domains from the Lewis structure rather than from the formula. Each

C=O

double bond is a single domain, so carbon has only two, and two domains always point in opposite directions. Oxygen's two lone pairs are invisible in the shape name but fully present in the geometry, which is why the molecule is bent rather than linear. This difference is also why

CO_2

is nonpolar while water is polar.

Determine the electron geometry, molecular shape, and approximate bond angle for

CH_2O

(formaldehyde), where carbon is the central atom.

Answer: Trigonal planar electron geometry, trigonal planar molecular shape, bond angles of approximately $120^\circ$ .

The Lewis structure has carbon double bonded to oxygen and single bonded to two hydrogens, with no lone pairs on carbon. The double bond counts as one domain, so carbon has three domains total:

AX_3

. Three domains spread to

120^\circ

in a flat triangle. Because there are no lone pairs on the central atom, the electron geometry and molecular shape have the same name. The double bond is slightly bulkier than the single bonds, so the actual

H-C-H

angle is a little under

120^\circ

while the

H-C=O

angles are a little over — a fine detail, but the reported answer is approximately

120^\circ

FAQ

Does a double bond count as two electron domains?: No. A double or triple bond counts as one electron domain. Both electron pairs in a double bond lie between the same two nuclei, so they point in one direction and behave as a single region of electron density. Treating a double bond as two domains is the single most common error in VSEPR problems and will change your predicted shape entirely.
What is the difference between electron geometry and molecular shape?: Electron geometry counts every domain, including lone pairs; molecular shape counts only the positions of atoms. They match when the central atom has no lone pairs. When lone pairs are present, the lone pairs still control the arrangement but are not part of the shape name — ammonia has tetrahedral electron geometry but trigonal pyramidal molecular shape.
Why are bond angles described as approximate?: The ideal angles come from spreading identical domains as far apart as possible. Real domains differ: a lone pair spreads wider than a bonding pair, and a multiple bond is bulkier than a single bond. Each lone pair pushes the bonds closer together, so the actual angle drops below the ideal — from $109.5^\circ$ to about $107^\circ$ in $NH_3$ and about $104.5^\circ$ in $H_2O$ .
Do lone pairs on the outer atoms affect the shape?: No. Only domains on the central atom determine geometry. The three lone pairs on each chlorine in $CCl_4$ or on each fluorine in $BF_3$ are ignored. Draw them in your Lewis structure for correctness, but count only what is attached to the central atom when you apply VSEPR.

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