Question

Explain the essential difference in how the valencebond method and molecular orbital theory describe a covalent bond.

Short answer

Valence bond theory suggests that covalent bonds form from the overlap of atomic orbitals, with each bond corresponding to two electrons in an overlap region directly between two atoms. On the other hand, molecular orbital theory, treats the molecule as a whole, and proposes that atomic orbitals combine to form molecular orbitals that belong to the entire molecule, wherein the bond results from the filling of these molecular orbitals by electrons.

Step-by-step solution

Understanding Valence Bond Theory

This theory proposes that bonds form from the overlap of atomic orbitals, which allows electrons to pair up. Each bond corresponds to two electrons in an overlap region directly between the two bonded atoms.

Understanding Molecular Orbital Theory

This theory considers a molecule as a collection of nuclei and electrons, where all the electrons move under the influence of all the nuclei. The original atomic orbitals combined and form new orbitals, which are the molecular orbitals that belong to the entire molecule.

Explaining Covalent Bond Formation in Valence Bond Theory

In valence bond theory for formation of a covalent bond, one electron from each atom are brought closer such that their atomic orbitals overlap. When two atomic orbitals overlap, they create a region of space where there are two electrons of opposite spins, and this region forms a covalent bond.

Explaining Covalent Bond Formation in Molecular Orbital Theory

In molecular orbital theory, when two atoms combine, their atomic orbitals merge to form new orbitals, called molecular orbitals. These molecular orbitals are spread over the entire molecule, rather than being localised between two atoms as in valence bond theory. The covalent bond results from the filling of these molecular orbitals by electrons.

Key concepts

Valence Bond Theory

Valence Bond Theory is a model that helps us understand how covalent bonds form. It emphasizes the idea that bonds are formed when atomic orbitals overlap. This overlapping involves a straightforward union where electrons from different atoms come together to share a space.

In this theory, each bond is a pair of electrons residing in a region directly between two atoms. These electrons must have opposite spins to coexist in the same orbital. Think of it like two atoms coming close enough so that their outermost shell, or valence shell, allows their electrons to "meet and greet" to form a bond.

This model also highlights that these overlaps create a stronger, more stable situation compared to the atoms being apart. The concept of overlapping explains nicely transparent visibility in what holds atoms together in the molecule.

Molecular Orbital Theory

Molecular Orbital Theory offers a more complex perspective on covalent bonds. Rather than just focusing on the overlap between two atomic orbitals, it considers an entire molecule’s framework. This theory views electrons as not just belonging to a single pair of atoms but being part of a global interaction with the whole molecule.

When atomic orbitals from the bonding atoms merge, they create new molecular orbitals. These new orbitals cover the entire molecule, not just the space between two specific atoms.

In essence, electrons will fill these molecular orbitals, and the presence of these filled orbitals leads to the formation of the covalent bond. This holistic approach provides fascinating insights into bond strength, magnetism, and the overall behavior of molecules.

Atomic Orbitals

Atomic Orbitals are regions around a nucleus where electrons are likely to be found. Each atom has a set of atomic orbitals that can hold a specific number of electrons.

The most commonly discussed orbitals in chemistry are the s, p, d, and f orbitals, each having distinct shapes and capacities.

For example:

• the s orbital is spherical and can hold up to 2 electrons.
• the p orbitals are dumbbell-shaped and can contain up to 6 electrons across three sub-orbitals.
• the d orbitals, with their more complex shape, can hold up to 10 electrons.

Understanding atomic orbitals is essential in exploring how atoms interact and bond with each other. These individual orbitals overlap in bonds, whether localized as in Valence Bond Theory or spread out as in Molecular Orbital Theory.

Electron Pairing

Electron Pairing is a fundamental concept in forming bonds and stabilizing atoms. Electrons have a property called spin, which can either be "up" or "down." In an orbital, two electrons can pair up only if they are of opposite spin.

This pairing is crucial because it allows atoms to achieve a lower energy state and greater stability. In the context of covalent bonds, like those described in Valence Bond Theory, this pairing happens in the overlapping orbitals between two atoms.

Electron pairing also determines how the electrons fill available orbitals, whether they are atomic or molecular. By filling lower energy orbitals first with paired electrons, atoms and molecules can find stability, forming the basis for solid chemical substances.