3.6 Predicting Molecular Shape

Learning Objectives

  • Utilising VSEPR theory, predict the geometries of molecules through the presence of lone-pairs and electron-dense groups.
  • Determine the polarity of a molecule, considering bond types and geometry.

You may have noticed that when drawing or representing molecules, we sometimes draw them with unexpected geometry. Some of the most well-known organic molecules, such as methane, formaldehyde and carbon dioxide all have different shapes.

Ball and stick molecular models. Carbon dioxide, on the left, has a linear structure, a central carbon with two oxygen atoms; in the middle, formaldehyde has a trigonal planar structure, a central carbon with two hydrogen atoms and one oxygen atom; and methane with a tetrahedral shape, a central carbon with four surrounding hydrogen atoms.
Figure 3.6.1: Carbon dioxide, formaldehyde and methane are all common organic compounds, with vastly different shapes from one another.

There seems to be sound reasoning behind this as all three compounds have different numbers of atoms present. However, consider carbon dioxide compared to a molecule of water; even though both molecules have the same number of atoms, we represent them differently. Water is found to have a bent structure, while carbon dioxide is perfectly linear (see Figure 3.6.2); why is this the case?

Ball and stick depiction of a molecule of water showing a 'bent' molecular shape: a larger red hydrogen is attached to two smaller white oxygens that are situated slightly lower. The bonds are angled downwards at 45 degrees.

Molecule of carbon dioxide with a linear shape.
Figure 3.6.2: Water and carbon dioxide have the same number of atoms, but present in different shapes.

VSEPR Theory

To understand why certain molecules present in this way, let us consider the repulsive effect of electrons through the Valence Shell Electron-pair Repulsion model, also known as VSEPR theory. Recall that electrons have negative charges. Electrons will, therefore, emit repelling forces upon each other and will try to roughly position themselves where the minimum repulsive force is experienced. By understanding how many high-density regions of electrons exist, we can best spread them out and predict geometry.

5 molecules represented by a central blue atom with pink electron dense regions. 2 electron dense regions give a linear structure. 3 electron dense regions give a trigonal planar structure. 4 electron dense regions give a tetrahedral structure. 5 electron dense regions gives a trigonal bipyramidal structure. 6 electron dense regions give a octahedral structure.
Figure 3.6.3: Different bond angles. Image attribution: Vector illustration of Bond Angles © aboabdelah – stock.adobe.com.

In Figure 3.6.3, we observe the preferred orientation of molecules depending on the number of electron-dense regions. The proposed shapes have all electron-dense regions positioned as far as possible from each other. These regions can either be bond pairs or lone electron pairs (on the central atom). In the case of carbon dioxide, we have two regions of high-electron density, each found at the bonds between the central carbon and oxygen atoms present. As such, a linear formation naturally occurs (see Figure 3.6.4).

A lewis dot diagram of carbon dioxide: between two oxygens is a central carbon atom with a delta plus sign indicating a partial positive charge—the oxygen and carbon are double-bonded. The oxygen atoms each have a delta minus sign signifying a partial negative charge, with two pairs of valence electrons drawn as four dots. We see two regions of electron density at each oxygen.
Figure 3.6.4: Lewis dot diagram of carbon dioxide; we see two regions of electron density (at each oxygen).

In the case of water, however, we detect the two electron-dense groups from the two O-H bonds and the two lone pairs of electrons on the oxygen atoms. This means there are four regions of electron density — inferring that the preferred configuration would be a tetrahedral arrangement. Naturally, however, the paired electrons do not generate bonds of their own. Removing two of the electron-dense regions produces a final ‘bent’ or angular geometry (see Figure 3.6.5).

Diagram of a water molecule with labeled dimensions. The oxygen atom at the center has two pairs of valence electrons, and is bonded to two hydrogen atoms situated below, the bonds drawn as single lines. The bond length of oxygen to hydrogen is marked as 95.7 picometers (pm), and the H-O-H bond angle is noted as 104.5 degrees, showing a bent molecular structure due to the repulsion between the pairs of valence electrons.
Figure 3.6.5: The bent angle of a water molecule is generated from the lone pairs present.
Three dimensional geometric illustration of a tetrahedral shape with a sphere inside the shapes centre, representing an atom at the intersection of four triangular faces.
Figure 3.6.6 “Tetrahedral Geometry.” Four electron groups orient themselves in the shape of a tetrahedron. Image attribution: Chem&121: Introduction to Chemistry Copyright © 2023 by Lake Washington Institute of Technology is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License.

As such, it is important to note that in VSEPR theory, all molecules take up their preferred orientations – however, the presence of lone pairs makes this appear as if it does not happen. For instance, methane, ammonia and water all have 4 regions of electron density and form tetrahedral arrangements (see Figure 3.6.6). Methane is able to accomplish this, while nitrogen (with one lone pair of electrons) forms a trigonal pyramid without another hydrogen bond. Water (with two lone pairs of electrons) forms a “bent” geometry – there is a clear angle between the two bonds.

Three diagrams: methane, ammonia and water. <a href="https://rmit.pressbooks.pub/rmitchemistrybridgingcourse/chapter/3-6-predicting-molecular-shape-2/">Transcript</a>.
Figure 3.6.7: Methane, ammonia and water. As the number of lone pairs of electrons increases, the shape changes. Transcript.

Therefore, the overall geometry of the molecule can be categorised by the number of atoms and the number of lone electron pairs present on the central atom. Figure 3.6.8 details a helpful guide to predict shape and can be relied upon as a general rule for most molecules. Let’s use the table to help us ascertain the geometry of another organic molecule: formaldehyde.

Table detailing the types geometries that can occur.
Figure 3.6.8: Table of Molecular Geometry. Image attribution: vector illustration of VSEPR Theory CHART. © aboabdelah – stock.adobe.com. Transcript.


Formaldehyde (CH2O) is represented in the following Lewis structure:
A lewis dot diagram of formaldehyde, with a central carbon atom single-bonded to two hydrogen atoms and double-bonded to an oxygen atom with two lone pairs of electrons.
As such, we observe three electron-dense areas, one for each of the C-H bonds, and one for the C-O bond (note how a double bond is still counted as a single region). There are no lone pairs on the central carbon. Therefore, the preferred geometry is trigonal planar.

3D molecular model of formaldehyde: a large black sphere representing the carbon atom is centrally located with two smaller white spheres for hydrogen atoms at an angle and one larger red sphere for the oxygen atom positioned above the carbon, showing the trigonal planar arrangement of atoms.

VSPER theory represents an idealised set of interactions between electron-dense areas within a molecule. Bond angles and lengths differ slightly between molecules dependant on the types of elements within. However, they are a great generalisation and should be used moving forward in undergraduate studies. Have a look at the simulation below for how real molecules differ to the model predictions.

Simulation by PhET Interactive Simulations, University of Colorado Boulder, licensed under CC-BY-4.0 (https://phet.colorado.edu).

Now that we understand the geometry of molecules, we can begin to determine if a molecule is polar or non-polar. Recall that these terms were used when discussing bond formation. Polarity refers to the presence of a (whole or partial) magnetic charge generated through atomic interactions, which can be determined through a difference in electronegativity.

Carbon dioxide is a linear molecule consisting of two double-bonded C=O’s. These bonds are polar covalent due to their difference in electronegativity. The carbon is electropositive, while the oxygens are electronegative. You would therefore expect the entire molecule to be considered polar. However, if we draw out these negative forces as charges, they directly oppose each other – cancelling each other out, as seen in Figure 3.6.9.

Figure 3.6.9: Although the C=O bonds are polar, they directly oppose each other. The net result is that the molecule remains non-polar. Transcript.

As such, while there are polar covalent bonds with electronegative ends, the entire molecule does not experience a net charge. This makes carbon dioxide non-polar.

Conversely, for water, the O-H bonds are also polar covalent. However, the introduction of the angular bend means that the forces experienced are not directly opposing each other. While the horizontal components of the force do cancel each other out, the vertical vectors ensure that a permanent dipole is present, meaning that the molecule is polar. <a href="https://rmit.pressbooks.pub/rmitchemistrybridgingcourse/chapter/3-6-predicting-molecular-shape-2/">Transcript.</a><a href="https://rmit.pressbooks.pub/rmitchemistrybridgingcourse/chapter/3-6-predicting-molecular-shape-2/">Transcript.</a><a href="https://rmit.pressbooks.pub/rmitchemistrybridgingcourse/chapter/3-6-predicting-molecular-shape-2/">Transcript.</a>

<a href="https://rmit.pressbooks.pub/rmitchemistrybridgingcourse/chapter/3-6-predicting-molecular-shape-2/">Transcript.</a>
Figure 3.6.10: Determining the polarity/net dipole of a molecule requires us to look at the horizontal and vertical components of the forces at play – cancelling out those that directly oppose. Transcript.

To understand how the polarity of a molecule is determined by the electronegativity of atoms, have a look at the following interactive activity’s “Three Atoms” simulation with partial charges enabled:

Simulation by PhET Interactive Simulations, University of Colorado Boulder, licensed under CC-BY-4.0 (https://phet.colorado.edu).

Key Takeaways

  • VSEPR Theory allows for the prediction of molecule geometry according to the number of electron-dense regions and lone pair electrons present.
  • Molecular geometry can result in certain molecules, which feature polar bonds, to be non-polar due to cancelling out forces.



Practice Questions




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