How to Determine Molecular Geometry Using VSEPR Theory

Molecular geometry describes the three-dimensional arrangement of atoms in a molecule. It determines whether a molecule is linear, bent, trigonal planar, tetrahedral, or shaped in another way. One of the most reliable ways to predict molecular geometry is through VSEPR theory, which stands for Valence Shell Electron Pair Repulsion theory.

The central idea is simple: electron groups around a central atom repel one another and arrange themselves as far apart as possible. By drawing a Lewis structure, counting the electron groups around the central atom, and accounting for lone pairs, you can predict a molecule’s shape and estimate its bond angles.

VSEPR theory is widely used in introductory chemistry because it connects a molecule’s electron arrangement to its observable three-dimensional structure. Although it does not explain every detail of molecular bonding, it provides a practical method for predicting the geometry of many common molecules and polyatomic ions.

What VSEPR theory explains

Electrons carry negative electric charge, so regions of electron density tend to repel one another. In a molecule, the electron density associated with bonds and lone pairs influences how the atoms arrange themselves in space.

A bonding pair is a pair of electrons involved in a bond between two atoms. A lone pair is a pair of valence electrons associated primarily with one atom rather than shared in a bond. Both contribute to electron repulsion around a central atom.

VSEPR theory treats these regions of electron density as electron groups. Their arrangement produces the molecule’s electron-group geometry, while the positions of the atoms determine its molecular geometry.

These two descriptions are related but not identical. Electron-group geometry accounts for both bonding groups and lone pairs. Molecular geometry describes only the spatial arrangement of the atoms, excluding lone pairs from the shape’s name.

For example, methane (CH₄) has four bonding groups around its central carbon atom and no lone pairs. Its electron-group geometry and molecular geometry are both tetrahedral. Water (H₂O), by contrast, has two bonding groups and two lone pairs around oxygen. Its electron-group geometry is tetrahedral, but its molecular geometry is bent because only the two hydrogen atoms define the molecular shape.

Understanding this distinction is essential to applying VSEPR theory correctly.

Step 1: Draw the Lewis structure

The first step in determining molecular geometry is to draw the molecule’s Lewis structure. This diagram shows how atoms are connected and represents their valence electrons as bonding pairs and lone pairs.

Valence electrons are the electrons in an atom’s outermost occupied shell. They are the electrons most directly involved in chemical bonding.

To draw a Lewis structure, begin by counting the total number of valence electrons contributed by all the atoms. For a negatively charged ion, add electrons equal to the magnitude of the negative charge. For a positively charged ion, subtract electrons equal to the magnitude of the positive charge.

Next, select a central atom, usually the least electronegative atom other than hydrogen. Electronegativity is an atom’s tendency to attract shared electrons toward itself. Hydrogen is generally placed at the end of a structure because it forms only one ordinary covalent bond.

Connect the surrounding atoms to the central atom with single bonds, then distribute the remaining electrons to complete the outer atoms’ electron shells. Finally, place any remaining electrons on the central atom and consider multiple bonds if needed to obtain a reasonable Lewis structure.

For example, carbon dioxide (CO₂) has the Lewis structure O=C=O. Carbon is the central atom, and each oxygen is connected to it by a double bond. Although each double bond contains two shared electron pairs, VSEPR theory counts each entire bond as one electron group around carbon.

A correct Lewis structure is important because an incorrect count of bonds or lone pairs can lead to an incorrect prediction of molecular geometry.

Step 2: Identify the central atom and count its electron groups

Once the Lewis structure is complete, focus on the central atom. Count the regions of electron density surrounding it, treating each bond to another atom as one group and each lone pair as one group.

The basic counting rules are:

• A single bond counts as one electron group.
• A double bond counts as one electron group.
• A triple bond counts as one electron group.
• Each lone pair on the central atom counts as one electron group.

This means that the number of electron groups is not the same as the total number of bonding electron pairs.

Consider carbon dioxide. The central carbon has two double bonds, one to each oxygen. Each double bond counts as one group, so carbon has two electron groups.

In methane (CH₄), the central carbon has four single bonds and no lone pairs, giving four electron groups.

In ammonia (NH₃), nitrogen has three single bonds to hydrogen and one lone pair, giving four electron groups.

This count provides the starting point for predicting the electron-group geometry. The next step is to determine how those groups are arranged and which positions correspond to atoms.

Step 3: Determine the electron-group geometry

VSEPR theory predicts that electron groups arrange themselves to reduce repulsion. For common molecules with two to six electron groups around a central atom, the ideal arrangements follow recognizable geometric patterns.

These angles are ideal reference values. Actual bond angles may differ because of lone pairs, differences between surrounding atoms, and other electronic effects.

Two electron groups: Linear geometry

Two electron groups arrange themselves on opposite sides of the central atom, producing a straight line with an ideal angle of 180°.

In carbon dioxide, the two carbon–oxygen double bonds occupy these opposite positions. Because carbon has no lone pairs, the molecular geometry is also linear.

Three electron groups: Trigonal planar geometry

Three electron groups arrange themselves in one plane, with ideal angles of 120° between them. This arrangement is called trigonal planar.

Boron trifluoride (BF₃) is a common example. Boron forms three bonds to fluorine and has no lone pairs in its usual Lewis structure. The three fluorine atoms occupy the corners of a triangle around the central boron atom.

Because all three electron groups are bonding groups, the molecular geometry is trigonal planar as well.

Four electron groups: Tetrahedral geometry

Four electron groups arrange themselves toward the corners of a tetrahedron, a three-dimensional shape with four triangular faces. The ideal angle between the groups is approximately 109.5°.

Methane is the standard example. Its four hydrogen atoms surround the central carbon atom in a tetrahedral arrangement.

A tetrahedral geometry is three-dimensional, not flat. Drawing a molecule with all its bonds in one plane can therefore give a misleading impression of its actual structure.

Five electron groups: Trigonal bipyramidal geometry

Five electron groups produce a trigonal bipyramidal arrangement. Three groups occupy positions around the central atom in an equatorial plane, separated by ideal angles of 120°. Two additional groups occupy axial positions above and below that plane.

Each axial group forms an ideal 90° angle with each equatorial group, while the two axial groups are separated by 180°.

Phosphorus pentachloride (PCl₅) is a familiar example of this geometry in its gas-phase molecular form.

The distinction between axial and equatorial positions becomes especially important when lone pairs are present because these positions do not experience identical repulsions.

Six electron groups: Octahedral geometry

Six electron groups arrange themselves toward the corners of an octahedron. The ideal angles between adjacent groups are 90°, and opposite groups are separated by 180°.

Sulfur hexafluoride (SF₆) is a common example. Six sulfur–fluorine bonds surround the central sulfur atom, producing an octahedral molecular geometry.

Recognizing these five basic electron-group geometries makes it much easier to identify molecular shapes, including those modified by lone pairs.

Step 4: Account for lone pairs

After identifying the electron-group geometry, determine how many groups are bonding groups and how many are lone pairs. This distinction determines the molecular geometry.

Lone pairs generally exert stronger repulsive effects on nearby bonding groups than bonding groups exert on one another. A lone pair’s electron density is concentrated around a single atom, whereas bonding electron density is distributed between two nuclei.

A useful qualitative order of repulsion is:

Lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.

This order is a practical guideline rather than a universal numerical rule. The strength of these interactions depends on the molecule’s electronic structure.

Because lone pairs occupy space and influence bond directions, they can change the molecular shape and compress bond angles.

Consider three molecules with four electron groups around their central atoms.

Methane has four bonds arranged tetrahedrally. Ammonia has three nitrogen–hydrogen bonds and one lone pair, so the three hydrogen atoms form a trigonal pyramidal shape. Water has two oxygen–hydrogen bonds and two lone pairs, producing a bent shape.

The ideal tetrahedral angle is approximately 109.5°. In ammonia, the H–N–H bond angle is about 107°. In water, the H–O–H angle is about 104.5°. These smaller angles are consistent with the greater repulsive influence of lone pairs.

This comparison illustrates why counting electron groups alone is not enough. You must also identify which groups are bonds and which are lone pairs.

Step 5: Determine the molecular geometry from the bonding groups

Once the number of electron groups and lone pairs is known, identify the positions of the bonded atoms and name the resulting molecular geometry.

For molecules with two, three, or four electron groups, the most common patterns are straightforward:

• Two groups and no lone pairs produce a linear molecule.
• Three groups and no lone pairs produce a trigonal planar molecule.
• Three groups with one lone pair produce a bent molecule.
• Four groups and no lone pairs produce a tetrahedral molecule.
• Four groups with one lone pair produce a trigonal pyramidal molecule.
• Four groups with two lone pairs produce a bent molecule.

Five and six electron groups allow additional shapes. For example, a central atom with five electron groups and one lone pair has four bonding groups and one lone pair. In a common trigonal bipyramidal electron-group arrangement, the lone pair preferentially occupies an equatorial position, leaving the bonded atoms in a seesaw molecular geometry.

With five electron groups and two lone pairs, the three remaining bonds commonly produce a T-shaped geometry when the lone pairs occupy equatorial positions.

For six electron groups, four bonds and two lone pairs commonly produce a square planar molecular geometry when the lone pairs occupy opposite positions. Five bonds and one lone pair produce a square pyramidal shape.

These examples show that molecular geometry depends on the arrangement of the bonding groups after lone pairs are accounted for, not simply on the total number of electron groups.

How to use VSEPR theory with polyatomic ions

VSEPR theory also applies to polyatomic ions, which are groups of covalently bonded atoms carrying an overall electric charge.

The main difference is that the total charge must be included when counting valence electrons for the Lewis structure. Once the structure is established, electron groups around the central atom are counted in the same way as for a neutral molecule.

For example, the ammonium ion (NH₄⁺) has four nitrogen–hydrogen bonds and no lone pair on the central nitrogen in its usual Lewis structure. It has four electron groups and a tetrahedral molecular geometry.

The nitrate ion (NO₃⁻) has three bonding groups around its central nitrogen and no lone pair on that atom in its usual resonance description. It therefore has trigonal planar geometry.

Resonance structures are alternative Lewis representations of the same delocalized electronic structure. They can show different placements of formal double bonds, but these alternative drawings do not mean that the ion rapidly switches between distinct molecular shapes. In nitrate, for instance, the three nitrogen–oxygen bonds are equivalent in the actual resonance hybrid, and the ion is approximately planar.

The overall charge affects the electron count, but the charge alone does not determine molecular geometry. The arrangement of electron groups around the central atom remains the key factor.

Step 6: Estimate the bond angles and recognize deviations

After determining molecular geometry, use the ideal electron-group arrangement to estimate bond angles. The angles indicate the directions of bonds relative to one another.

Linear molecules have an ideal angle of 180°. Trigonal planar molecules have ideal angles of 120°, while tetrahedral molecules have ideal angles of approximately 109.5°. Trigonal bipyramidal and octahedral arrangements include ideal angles of 90° and 120°, depending on the positions involved.

Actual bond angles may differ from these ideal values. Lone pairs often compress angles between bonds, as seen in ammonia and water. Different surrounding atoms can also affect bond angles because they do not necessarily distribute electron density in the same way.

Multiple bonds count as one electron group under the basic VSEPR method, but they can exert greater repulsive influence than single bonds because they contain more electron density between the bonded atoms. This can also contribute to deviations from ideal angles.

For example, a central atom with three electron groups may have a trigonal planar electron-group geometry without producing perfectly equal 120° bond angles if the surrounding atoms differ or the electronic environment is asymmetric.

Therefore, ideal angles should be treated as starting points rather than exact measurements for every molecule.

Worked examples of molecular geometry

Applying the same procedure to several familiar molecules helps make the method more intuitive.

Carbon dioxide (CO₂)

Carbon dioxide has a central carbon atom bonded to two oxygen atoms through double bonds.

1. Draw the Lewis structure: O=C=O.
2. Count the electron groups around carbon: two.
3. Identify lone pairs on carbon: none.
4. Determine the electron-group geometry: linear.
5. Determine the molecular geometry: linear.

The O–C–O angle is 180° because the two bonding groups lie on opposite sides of the central carbon atom.

Ammonia (NH₃)

Ammonia contains a central nitrogen atom bonded to three hydrogen atoms.

1. Draw the Lewis structure with three nitrogen–hydrogen bonds and one lone pair on nitrogen.
2. Count the electron groups around nitrogen: four.
3. Identify the electron-group geometry: tetrahedral.
4. Exclude the lone pair when naming the molecular shape.
5. Determine the molecular geometry: trigonal pyramidal.

The H–N–H bond angle is approximately 107°, slightly smaller than the ideal tetrahedral angle because of the lone pair’s repulsive influence.

Water (H₂O)

Water contains a central oxygen atom bonded to two hydrogen atoms.

1. Draw the Lewis structure with two oxygen–hydrogen bonds and two lone pairs on oxygen.
2. Count the electron groups around oxygen: four.
3. Identify the electron-group geometry: tetrahedral.
4. Exclude the lone pairs when naming the molecular shape.
5. Determine the molecular geometry: bent.

The H–O–H bond angle is approximately 104.5°. The two lone pairs help explain why this angle is smaller than the ideal tetrahedral angle.

Sulfur dioxide (SO₂)

Sulfur dioxide has a central sulfur atom bonded to two oxygen atoms. Its Lewis structure is commonly represented using resonance forms, reflecting electron delocalization across the molecule.

In the usual introductory VSEPR treatment, sulfur has two bonding groups and one lone pair. The double-bond character represented in resonance does not increase the count of bonding groups beyond two.

Three electron groups correspond to a trigonal planar electron-group arrangement. Because one group is a lone pair, the molecular geometry is bent, and the O–S–O angle is somewhat less than 120°.

This example demonstrates why the electron-group count should be based on regions around the central atom rather than on the number of lines or electron pairs drawn in a Lewis structure.

Common mistakes when applying VSEPR theory

Several errors can lead to incorrect molecular shapes even when the basic principle is understood.

Confusing electron-group geometry with molecular geometry is one of the most common mistakes. Electron-group geometry includes lone pairs; molecular geometry describes the positions of atoms. Water has tetrahedral electron-group geometry but bent molecular geometry.

Counting a double or triple bond as multiple electron groups is another frequent error. Each bond connecting the central atom to a neighboring atom counts as one electron group, regardless of bond order. Multiple bonds can affect repulsion, but they do not occupy multiple positions in the basic VSEPR count.

Ignoring lone pairs can produce the wrong shape and bond-angle estimate. A central atom with four electron groups may have a tetrahedral, trigonal pyramidal, or bent molecular geometry, depending on how many of those groups are lone pairs.

Assuming every bond angle equals its ideal value can also be misleading. Ideal angles describe symmetrical reference arrangements. Lone pairs, different surrounding atoms, and multiple bonds can cause measurable deviations.

Finally, choosing the wrong central atom or drawing an incorrect Lewis structure can undermine every later step. Electron counting and connectivity must be established before molecular geometry can be predicted reliably.

What VSEPR theory can and cannot predict

VSEPR theory is a qualitative model. It is particularly useful for predicting the shapes of many small molecules and polyatomic ions when a reasonable Lewis structure is available.

Its strength is that it links electron-group arrangements to molecular shapes through a small set of geometric principles. It also helps explain why molecules with the same number of electron groups can have different shapes when their numbers of lone pairs differ.

However, VSEPR theory does not provide a complete quantum-mechanical description of bonding. It does not always predict exact bond angles, and its simple rules can become less reliable for unusual bonding arrangements, certain transition-metal compounds, and molecules whose structures are strongly influenced by electronic effects beyond the basic model.

When precise molecular structures or bond angles are required, experimental measurements and more detailed computational methods may be necessary.

Molecular geometry also has consequences beyond a molecule’s appearance. Shape influences how molecules approach and interact with one another, and it can contribute to polarity, reactivity, and biological recognition. However, molecular shape alone does not determine all these properties; bond polarity and the distribution of electron density also matter.

The practical value of VSEPR theory lies in providing a dependable first prediction. By drawing the Lewis structure, counting electron groups, identifying lone pairs, and distinguishing electron-group geometry from molecular geometry, you can determine the shapes of many common molecules systematically and explain why those shapes occur.