Molecular Geometry — Shapes & Examples

When we studied VSEPR theory, we learned that molecules have different geometries. Valence bond theory taught us about sigma vs pi bonds and hybridization. They demonstrate how atomic orbitals overlap. The various combinations of overlap generate unique geometries.

We determined molecular geometries by counting the central atom's bonds and lone pairs. Molecular geometries have five main classifications. For each classification, we assume that the central atom's bonds are equal. The equivalent bonds will tell us the ideal bond angle. Deviations from the ideal bond angle happen when the central atom has lone pairs.

A Few Terms You Should Know

When we describe molecular geometries, we need to know a couple of terms. When the bonding atoms are directly above or below the central atom, we call them axial. When the atoms distribute around the center plane of the molecule, we call them equatorial. Many of the molecular geometries are combinations of axial and equatorial positions. However, some have bonds that angle away because of electrostatic interactions.

Linear Molecules

The most simple geometry is linear. In linear molecules, the central atom bonds with two other atoms. The molecule is a straight line with bond angles of 180 degrees. The central atom cannot have lone pairs of electrons. So, linear molecules cannot deviate.

An example of linear geometry is carbon dioxide ( $CO_{2}$ ). Carbon is the central atom. The carbon is double bonded on both sides to oxygen. The bond angle between the oxygens is 180 degrees. The bond angle maximizes the distance between the oxygen atoms.

Trigonal Planar

Adding a third atom to the central atom creates the trigonal planar geometry. The trigonal planar's three bonds are on the same plane. They are equatorial. If you looked straight at a trigonal planar molecule, you would see a triangle. The bond angle is 120 degrees. The trigonal planar geometry can have one deviation. If you replace one of the bonded atoms with a lone pair, the geometry becomes angular.

Boron trifluoride ( $BF_{3}$ ) is an example of a trigonal planar molecule. The distance between the fluorines is 120 degrees. An angular deviation example is sulfur dioxide. Remember, it's an exception to the octet rule. Why? Because it has access to the 3d orbitals. Drawing the lewis structure of sulfur dioxide shows us that it has one lone pair. The angle between the two oxygen atoms is 119 degrees. We see a minimal bond angle deviation.

Tetrahedral

The most common geometry is tetrahedral. In a tetrahedral geometry, the central atom bonds to four other atoms. One bond is axial, directly above the central atom. The other three bonds angle down from the center of the atom. The angle between the top atom and the bottom atoms is 109.5 degrees. If you draw a line between each bonded atom, you form a three-sided pyramid.

Let's look at the two deviations for the tetrahedral geometry. The trigonal pyramidal has one lone pair and three bonds. A familiar example is ammonia ( $NH_{3}$ ). In ammonia, the lone pair sits on top of the nitrogen. The three hydrogens angle away, creating a 107-degree angle. The second deviation is the bent geometry. The most famous example is water ( $H_{2}O$ ). In water, the oxygen atom has two lone pairs that repel the hydrogens. The bond angle in water is 104.5 degrees. We are beginning to have more significant deviations from the ideal bond angles.

Trigonal Bipyramidal

Molecules that have a trigonal bipyramidal shape are exceptions to the octet rule. They have five bonds! Some examples are when phosphorous or sulfur bond to five halogens. The trigonal bipyramidal has two axial bonds and three equatorial bonds. If you were to connect the bonded atoms, it would have two three-sided pyramids. One pyramid would point up. The other pyramid would point down.

The trigonal pyramidal geometry has unique deviations. The less complex deviation has three lone pairs and two bonds. The three electron lone pairs distribute around the molecule's center. One bond is directly above and the other below. We know that three atoms in a straight line are linear! An example of the trigonal bipyramidal linear deviation is xenon difluoride ( $XeF_{2}$ ). Xenon starts with a full octet. So, with two bonds, it has three lone pairs.

The deviation with two lone pairs is T-shaped. An example of a T-shaped molecule is chlorine trifluoride ( $ClF_{3}$ ). The three fluorine atoms form a "T." Two of the fluorines are axial. The third fluorine and the two lone pairs are equatorial.

Finally, the deviation with three lone pairs is the seesaw. An example is sulfur tetrafluoride ( $SF_{4}$ ). In sulfur tetrafluoride, the lone pair sits directly above the sulfur. The four bonding fluorines angle down, away from the lone pair.

Octahedral

The final class of molecular geometries is octahedral. Octahedral molecules have six bonds! Two of the bonds are axial. The other four bonds are equatorial. In sulfur hexafluoride ( $SF_{6}$ ), sulfur uses six valence electrons to bond with six fluorines. The bond angle between each fluorine is 90 degrees.

When we replace one of the bonds with a lone pair, we have a square pyramidal. In a square pyramidal molecule, one bond and the lone pair are axial. The remaining four bonds are equatorial. If you drew lines between the bonding fluorines, you would form a four-sided pyramid. An example is bromine pentafluoride ( $BrF_{5}$ ). Bromine uses five of its seven valence electrons to bond, leaving it with one lone pair.

If an octahedral molecule has two lone pairs, it forms a square planar. In xenon tetrafluoride ( $XeF_{4}$ ), the four fluorines are equatorial. The two lone pairs are axial.

Molecular Geometry and Polarity

Knowing the molecular geometry of a molecule can help us determine its polarity. Polar covalent bonds form between atoms with large electronegativity differences. But, that does not guarantee the molecule is polar. For example, carbon dioxide has two polar carbon-oxygen bonds. But, carbon is a linear molecule. The two dipoles are in opposite directions. The opposite dipoles cancel, so carbon dioxide is not a polar molecule. Now, let's consider ammonia. Ammonia has trigonal pyramidal geometry with polar nitrogen-hydrogen bonds. The hydrogens angle down, away from the lone pair. The dipoles in ammonia are not in opposite directions. So, ammonia is a polar molecule. In fact, the electronegativity difference between nitrogen and hydrogen is massive. So, it's an example of hydrogen bonding.