How to Study the Chemical Theory of Valence Bonds

Co-authored by Meredith Juncker, PhD

Last Updated: July 23, 2020 References

Atoms are considered the building blocks of matter. As such, the properties and interactions of atoms are of great interest to scientists. One important property of an atom is how many electrons it has in its outermost shell. These are known as valence electrons and are responsible for the bonding interactions of that atom. The valence bond theory aims to describe and predict these interactions. To study valence bond theory, you will need to visualize atomic orbitals, overlap them, and understand their geometries.

Method 1

Method 1 of 4:

Visualizing Atomic Orbitals

Atoms are made up of protons (positively charged particles), neutrons (particles with no charge), and electrons (negatively charged particles). Protons and neutrons make up the mass of the atom and rest at the center of the atom. Electrons are so small that their mass is negligible, and they orbit around the center of the atom.^{[1]
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Know that electrons reside at different levels. Electrons do not orbit the nucleus randomly. Instead, they remain in orbitals that reside at specific distances from the nucleus (this distance varies by atom). Orbitals closer to the nucleus are considered low orbitals and those further away are high orbitals. The more energy an electron has, the higher orbital state it will occupy.^{[2]
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- Orbitals refer to the probable zone in which you can find the electron.
- Electrons are most stable in the lowest possible energy state, also known as the ground state.
3
Remember that electrons have different orbital patterns. When conceptualizing the electron cloud (the space in which electrons can be found), many people naturally envision a sphere around the nucleus. While some orbitals are spherical (s orbitals), others are shaped like dumbbells with the nucleus in the center (p orbitals). These different shapes are important to the concept of valence bonds, and must be taken into account when you analyze the bonds between two atoms.^{[3]
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- There are also d orbitals and f orbitals that have more complex geometry.
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Method 2

Method 2 of 4:

Overlapping Atomic Orbitals

1
Picture single bonds. Single bonds, or sigma (𝝈) bonds, are the result of two s orbitals overlapping. Electrons are shared in the overlap region, and this region can be found between the two nuclei. For this reason, the area is referred to as the internuclear axis.^{[4]
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- Sigma bonds overlap head-on. This means they have the most effective overlap, and thus form the strongest bond.
2
Imagine a bond outside the internuclear axis. While sigma bonds all reside between the nuclei of two atoms, p orbitals form a different kind of bond. Because of the shape of a p orbital, it forms what is known as a pi (𝝅) bond. The pi bond exists above and below the nuclei of the atoms, and therefore, is outside of the internuclear axis.^{[5]
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- P orbitals do not overlap as well as s orbitals, so pi bonds are easier to break (weaker) than sigma bonds.
- Above and below the nuclei is the accepted orientation for the first pi bond. However, it is possible to have another pi bond that is perpendicular to the first. This bond would be considered to reside on either side of the nuclei.
To visualize these different bonds, you have to superimpose the orbital of one atom on the orbital of the other. To visualize pi bonds, imagine two dumbbells being pushed together. The tops and bottoms would touch, but the centers would not. Sigma bonds can be compared to two balls being forced together. They meet head on and the bond resides in the internuclear axis, which could be compared to the space between the centers of the two balls.^{[6]
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Method 3

Method 3 of 4:

Correcting for Geometry

1
Consider the charge of electrons. Electrons are attracted to the nucleus at the center of the atom because it is positive and they are negative. That also means that electrons are repelled by each other. An atom is at its lowest energy state (most stable) when the electrons are as far from each other as possible. This makes the geometry of electron orbitals very important to the valence bond model.^{[7]
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- Electrons repelling from one another is commonly referred to as the Valence Shell Electron Pair Repulsion theory, or the VSEPR theory.
- Common types of geometry for atomic orbitals are linear, bent, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.
2
Factor in lone pairs. Lone pair electrons are valence electrons in an atom that do not bond with another atom. Since they are not bonding with other atoms and being pulled outward by other nuclei, the lone pairs orbit closer to the center of the atom. This exerts slightly more repulsive force on the other electrons, and alters the shape of the atom or molecule.^{[8]
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- For example, water could be expected to be linear (H-O-H), but the oxygen has two lone pairs of electrons that interact with the shape of the molecule. This pushes the hydrogens closer together than they would be otherwise, and gives the molecule a bent geometry.
3
Hybridize orbitals. In some cases, an atom’s bonding is not accurately predicted by the s and p orbitals it possesses. When this happens, valence bond theory suggests that the orbitals of the atom have been hybridized. In short, that’s a way of saying that some s and p orbitals merged together to form orbitals that share characteristics of both and increase the stability of the atom. This phenomenon helps to predict the shape and bonding activity of some atoms.^{[9]
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- For example, carbon is sp³ hybridized (1 s and 3 p orbitals merged). This allows the orbitals to spread out optimally and reduce electron-electron repulsion. It also allows the carbon atom to form four bonds.
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Method 4

Method 4 of 4:

Keeping Successful Study Habits

If you are just starting your first chemistry class, be warned that they are not for the faint of heart. You should read the chapter ahead of time and take notes on what you read. This will help you identify any questions you have about the valence bond theory in advance.^{[10]
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Pay attention to the class. Chemistry students are rarely successful at memorizing everything needed to do well with the valence bond theory. Instead, engage in the class and practice thinking in a scientific way. You should also take notes in the class so that you remember the important points that are covered pertaining to the valence electrons and bonding.^{[11]
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- Draw models of orbitals and molecular geometry to help you visualize what is happening.
Lab classes are a critical part of understanding chemistry. They provide a hands on application for many of the concepts that you will see in the textbook. It is important to come to lab prepared, and to follow all safety precautions. This is where you can see some of the chemical reactions predicted by valence bond theory in action.^{[12]
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Study regularly. Cramming the night before a chemistry exam is rarely successful, and always miserable. Instead, set aside an hour or two after every class to review the material covered in the class and refresh yourself on older material. This will help you stay focused and confident in your chemistry class.
- Study groups are a great way to make studying more fun.^{[13]
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- Find practice exams online or create your own to test your knowledge on the content.
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Tips

Review the valence bond theory several times.

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Ask your instructor if you need additional help.

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Warnings

Do not try to learn valence bond theory in a single cramming session.

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New discoveries in chemistry and physics could alter the valence bond theory in the future.

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The valence bond theory is a model. It describes and predicts the outcomes of atomic interactions in many cases. However, it is still a theory and is subject to error.

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