Molecular Orbital Theory: Videos & Practice Problems

Molecular orbital theory (MOT) is a fundamental concept in organic chemistry that helps explain the behavior of electrons in molecules. Understanding MOT is crucial for grasping various chemical reactions, particularly those involving resonance and hybridization. This theory provides a framework for visualizing how atomic orbitals combine to form molecular orbitals, which can be bonding, antibonding, or nonbonding.

At the core of MOT is the idea of conjugation, which refers to the ability of electrons to resonate between adjacent nonbonding orbitals. Resonance allows for the delocalization of electrons, enhancing the stability of molecules. Nonbonding orbitals are those that do not participate in bonding with other atoms, and they play a significant role in resonance structures. It is important to note that resonance involves the movement of electrons, not atoms, which remain fixed in their positions.

In organic chemistry, the outermost electron shell is typically the focus, particularly the second shell for carbon-based compounds. For example, carbon, with an atomic number of 6, has an electron configuration that begins filling from the lowest energy orbitals according to the Aufbau principle. The first two electrons occupy the 1s orbital, the next two fill the 2s orbital, and the remaining two electrons occupy the 2p orbitals, following Hund's rule. This configuration leads to the understanding of hybridization, where carbon can form four bonds by blending its 2s and 2p orbitals.

When carbon has three bond sites, it undergoes sp² hybridization, resulting in three equivalent sp² orbitals. This hybridization allows for the formation of sigma bonds with other atoms, such as hydrogen or other carbon atoms. The remaining unhybridized 2p orbital, which is higher in energy, can participate in pi bonding, particularly in double bonds. The presence of this nonbonding orbital is essential for conjugation, as it can interact with other nonbonding orbitals to form new molecular orbitals.

Conjugation can occur with various types of nonbonding orbitals, including those from radicals, cations, anions, or lone pairs. When these orbitals interact, they can create new molecular orbitals that facilitate electron sharing, enhancing the stability and reactivity of the molecule. Understanding how atomic orbitals combine to form molecular orbitals is key to mastering molecular orbital theory and its applications in organic chemistry.

When adjacent non-bonded atomic orbitals come together, they create molecular orbitals through a process known as the linear combination of atomic orbitals (LCAO). This method allows us to predict the shape and characteristics of molecular orbitals formed from the overlap of atomic orbitals. For instance, in ethene, each conjugated atom contributes one atomic orbital, which can be visualized as a cloud of electron density where electrons are likely to be found.

Importantly, atomic orbitals do not collide like solid objects; instead, they interact as waves. This interaction can lead to two types of interference: constructive and destructive. Constructive interference occurs when the waves of the atomic orbitals align positively, enhancing the probability of finding electrons between the atoms, resulting in a bonding interaction. This is referred to as an in-phase overlap, where the combined wave functions increase the electron density in the region between the two atoms.

Conversely, destructive interference happens when the atomic orbitals are out of phase, causing their wave functions to cancel each other out. This results in a node, a region where the probability of finding electrons is zero, leading to an antibonding interaction. In this scenario, the atoms repel each other, making the bond unstable.

It is crucial to note that the positive and negative lobes of atomic orbitals do not indicate actual charges; they simply represent different phases of the wave functions. The Pauli exclusion principle plays a significant role in electron configuration, stating that no more than two electrons can occupy a single orbital. Therefore, while electrons prefer to fill bonding molecular orbitals, they may occupy antibonding orbitals only when necessary, such as when there are excess electrons.

In the case of alkenes, the presence of two electrons that can be shared in a bonding molecular orbital leads to the formation of a strong double bond. This bonding interaction is favored because it stabilizes the molecule, allowing for effective sharing of electrons between the atoms involved. Understanding these principles of molecular orbital theory is essential for predicting the behavior of molecules and their interactions in various chemical contexts.