Deviation From Ideal Bond Angles — Overview & Examples

Ideal Bond Angles

Ideal bond angles are the angles formed from the valence shell electron repulsion (VSEPR) theory. This theory states that the lone and bonding electron pairs want to be as far away from each other as possible in molecules. The ideal bond angles arise from the electron repulsion and create the maximum distance between all electron pairs. To see a chart of the ideal bond angles and electron geometries check out VSEPR Theory.

Deviation From Ideal Bond Angles

Ideal bond angles are determined by ideal geometry. However, when analytical chemists evaluate some molecule geometries experimentally, the bond angles deviate from the ideal angles. Usually, deviation of bond angles occurs in molecules with lone pairs on the central atom. Lone pair electrons spread out more than bonding electrons. That's because, in a covalent bond, the electrons are confined to the space between two atoms. The lone pairs are therefore larger and can push the bonding pair electrons closer together. This causes a decrease in the bond angle compared to the ideal angle.

Water Molecules and Bond Deviation

Water molecules are a great example of bond angle deviation. Water is known to have a bent geometry from the two hydrogen bonds to the central oxygen atom and two lone lairs. Water is not a linear molecule where the bonding pairs are 180$^{\circ}$ from each other. This is because the repulsion between lone pairs and bonding pairs causes a bent geometry. The ideal angle of 109.5$^{\circ}$ is expected for bent geometries. However, experiments show that the bond angle between the hydrogens atoms is 104.5$^{\circ}$. The angle decrease is caused by the slightly larger lone pair electron clouds.

Deviation From Ideal Bond Angles: Why is there a deviation from ideal bond angles?

(Video) VSEPR Theory Practice Problems

by Tyler DeWitt

In this video, Tyler DeWitt walks you through VSEPR theory practice problems. You'll learn the steps to take to solve any VSEPR problem. You'll also see examples of molecules with ideal bond angles and deviations from ideal bond angles. If you need help with predicting bond angles, this is a great explanation!

More Causes of Deviations from the Ideal Angles

Size of Central Atom

Do you remember the atomic radius trend? It increases down a group of the periodic table. Why? We are adding energy levels. Each additional level pushes the valence electrons further from the nucleus. How does this affect bonding? The two bonding atoms will be further from each other. So, the bond length also expands. How does this affect the ligands? There will be more space between ligands. The electrostatic repulsion gets reduced! So, a larger central atom causes a decrease in the bond angle. For example, let's consider ammonia ($NH_{3}$) and arsenic trihydride ($AsH_{3}$). They are both in group five. But arsenic is further down. So, it's larger. Arsenic trihydride has a bond angle of 91.8 degrees. Ammonia is much larger at 107.3 degrees.

Size of Ligand

The size of the ligand also affects electrostatic repulsion. As ligand size increases, so does the bond angle. Why? Large ligands have more prominent electron clouds. So, there will be more electrostatic repulsion. The bond angle expands to minimize the repulsion. For example, consider trifluoromethane ($CHF_{3}$), tribromomethane ($CHBr_{3}$), and triiodomethane ($CHI_{3}$). The halogen-carbon-halogen bond angles are 108.0, 110.0, and 113.0, respectively. The bond angle increases as we move down the halogen series!

Electronegativity

Unfortunately, size isn't the only factor. We also have to consider electronegativity. The greater an element's electronegativity, the more it will pull electron density. In the VSEPR model, electron lone pairs are the prominent source of electron density. In polar bonds, the more electronegative atom pulls the electron lone pairs. So, the repulsion caused by the lone pairs gets limited. What's the result? Our bond angle decreases.

As a side note, size and electronegativity are often competing factors. But, electronegativity tends to be more prominent. Why? Because the electronegativity recasts the electron cloud. More specifically, it alters the lone pairs. In VSEPR, the lone pairs contribute the most to repulsion. Unfortunately, the trend is not universal. So, on some occasions, size matters more.

Multiple Bonds

When we first learned VSEPR, we assumed double and triple bonds had little impact. Sadly, that's not entirely true. More bonds equal more electron density. Remember, in the VSEPR model, we assume bonded electrons are localized. They get stuck between the atoms. With multiple bonds, that means we have more electrons squeezed into place. So, there is more electron density that repels other groups. A great example of this effect is the organic compound formaldehyde. The IUPAC name is methanal ($CH_{2}O$). Carbon is the central atom. It's double-bonded to the oxygen and single-bonded to each hydrogen. We would expect a trigonal planar geometry with 120-degree angles. But, the double bond increases the repulsion. So, the hydrogen-carbon-hydrogen bond angle is 118 degrees.

Exception to the Octet Rule: Trigonal Bipyrimidal

Before discussing the exceptions to the octet rule, we should learn some new terms. When the bonding atoms are directly above or below the central atom, we call them axial. The other option is to have the ligands distributed around the center plane of the molecule. We call this position equatorial. Let's consider the trigonal bipyramidal geometry. It has two axial and three equatorial positions. It turns out that lone pairs and electronegative elements favor the equatorial position. Why? Around the center, the bond angle is 120 degrees. In the axial position, the bond angle is 90 degrees. So, the equatorial position minimizes the repulsion.