Valence Bond Theory

If you only remember one question about valence bond theory, remember this: which orbitals are overlapping to make the bond? Valence bond theory explains a covalent bond as overlap between atomic orbitals on neighboring atoms, with the bonding electrons localized mainly between the two nuclei.

That one question gets you much closer to the chemistry than drawing a line between two atoms and stopping there. It is also what separates valence bond theory from its main rival, molecular orbital theory.

Valence Bond Theory vs. Molecular Orbital Theory At A Glance

Both theories describe covalent bonding, but they emphasize different pictures.

Valence bond theory focuses on localized bonds between specific pairs of atoms. Molecular orbital theory uses orbitals that can extend across the whole molecule.

When To Reach For Each Model

For many introductory problems, valence bond theory gives a fast local picture. If electrons are strongly delocalized, molecular orbital theory often describes the distribution more naturally.

Neither model should be treated as the only correct language for every molecule. The better model depends on what feature you are trying to explain. Use valence bond theory when you want to explain:

• why a covalent bond forms between two atoms
• why a bond is a $\sigma$ bond or a $\pi$ bond
• why bonding has direction
• why hybridization helps account for bond angles in many common molecules

It is most practical when a molecule can be described reasonably well with localized bonds and a Lewis-structure-style picture, which is why it shows up so often in introductory organic chemistry.

What Valence Bond Theory Says

In introductory chemistry, valence bond theory usually emphasizes three ideas:

• a covalent bond comes from overlap of atomic orbitals
• the bonding electrons are treated as localized mainly between two atoms
• bond type and bond direction depend on how the orbitals overlap in space

That is why orbital shape matters. A head-on overlap gives a $\sigma$ bond. A side-by-side overlap of parallel unhybridized $p$ orbitals gives a $\pi$ bond.

Not every overlap is equally effective. More effective overlap usually means more electron density between the nuclei and, within this model, a stronger bonding interaction. Because orbitals point in specific directions, valence bond theory helps explain why many covalent bonds have predictable shapes instead of random arrangements.

Hybridization extends the model by allowing orbitals on the same atom to mix before bonding. Labels such as $sp$, $sp^2$, and $sp^3$ help explain common linear, trigonal planar, and tetrahedral bonding arrangements. In most first-course chemistry settings, hybridization is taught as part of the broader valence bond framework, not a separate competing theory.

Applying The Model: Ethene And Why A Double Bond Has Two Parts

Ethene, $C_2H_4$, is a strong example because it shows both hybridization and orbital overlap in one molecule.

Each carbon in ethene is commonly described in introductory valence bond theory as $sp^2$ hybridized. That gives each carbon three $sp^2$ orbitals in one plane and one unhybridized $p$ orbital perpendicular to that plane.

Then the bonding picture is:

• one $sp^2$ orbital from one carbon overlaps head-on with an $sp^2$ orbital on the other carbon to make the $C-C$ $\sigma$ bond
• the remaining $sp^2$ orbitals on each carbon overlap with hydrogen $1s$ orbitals to make the four $C-H$ $\sigma$ bonds
• the unhybridized $p$ orbitals on the two carbons overlap side-by-side to make one $\pi$ bond

So the carbon-carbon double bond in ethene is not two identical bonds. It is one $\sigma$ bond plus one $\pi$ bond. That is a direct valence bond explanation of both bond type and geometry.

Confusion Points Worth Avoiding

Treating the theory as just a Lewis structure with new words

The point is not the bond line by itself. The point is the orbital overlap that creates bonding electron density between atoms.

Assuming every molecule is best described by fully localized bonds

Valence bond theory works best as a localized bonding model. If a molecule has strong delocalization, a purely local picture can miss important behavior, and molecular orbital theory becomes the better choice.

Forgetting the condition behind a $\pi$ bond

A $\pi$ bond needs parallel unhybridized $p$ orbitals that can overlap side-by-side. If that geometry is not available, the usual $\pi$-bond picture does not apply.

Treating hybridization as a separate theory

In most introductory courses, hybridization is an extension inside valence bond theory, not a completely unrelated theory.

A good way to test your grip on the model is acetylene, $C_2H_2$: ask which orbitals overlap to make the $C-C$ bond, how many $\pi$ bonds are present, and why the geometry is linear. Comparing $sp$, $sp^2$, and $sp^3$ across molecules is the natural next step.