Molecular Orbital Theory and Diatomic Molecules

A comprehensive set of flashcards covering key concepts of Molecular Orbital Theory as discussed in the lecture notes for CHEM 201.

Molecular Orbital (MO) Theory

A theory explaining how atomic orbitals combine to form molecular orbitals, which determine the bonding and properties of molecules.

• Atomic orbitals combines to create molecular orbitals

• number of eletrons of molecular orbitals = number of eletron of comning AOs

• overlap = similar phase, energy and orientation 

• Bondng 1s in MOs has higher eletron density between 2 nuclei

• antibonding 1s MOs has 0 eletorn density between 2 nuclei due to node

• THe enrgy level of antibonding is raised slightly more than bonding ( in-phase) is lowered

Bonding Molecular Orbitals (MOs)

Molecular orbitals formed by constructive interference of atomic orbitals; they stabilize the molecule by allowing electron sharing between atoms.

Antibonding MOs (σ, π)

Molecular orbitals formed by destructive interference of atomic orbitals; they DESTABILIZE the molecule by reflecting a HIGHER ENERGY STATE

Bond Order

number of eletrons pairs shared between two atoms 

Paramagnetic

A property of a substance that has unpaired electrons, causing it to be attracted to a external magnetic field.

Diamagnetic

A property of a substance that has all electrons paired, resulting in no net magnetic moment, thus repelling magnetic fields.

σ Bond

• overlap between orbital “ HEAD TO HEAD” 

• one region overlap 

• high electron density between two nuclei 

• rotation arround bond axis of signma is allowed 

• same orientation, enrgy, phase, to be overlapped’

• sigma bond is stronger than pie bond. because it sigma the orbital overlaps more efficiently, and pie bond is less efficient

π Bond

• overlap between orbital “ SIDE BY SIDE” 

• two region overlap

• eletron density is divided into two region, above and below. Along it, there is node cause 0 eletrons density

• rotation arround bond axis of pie bond is restricted because cause breaking

•  same orientation and phase adn energy to be overlaped.

Constructive Interference

A phenomenon where overlapping wave functions reinforce each other, leading to increased amplitude in molecular orbitals.

Destructive Interference

A phenomenon where overlapping wave functions cancel each other out, leading to decreased amplitude in molecular orbitals.

MO Diagram for H2

A graphical representation showing the energy levels of molecular orbitals for the H2 molecule, indicating bonding and antibonding orbitals.

Unpaired Electrons

Electrons that are alone in an orbital and not paired with another electron with opposite spin, contributing to a paramagnetic property.

Aufbau Principle

• fill lowest energy MO available

Pauli's Exclusion Principle

• two eletrons must be opposite spin in one orbital

Hund’s Rule

A rule stating that electrons will fill degenerate orbitals singly before pairing up, to maximize total spin.

He2 Molecule Stability

Cannot exist because it has an equal number of bonding and antibonding electrons, resulting in a bond order of zero.

Electron Configuration

A notation that shows the arrangement of electrons in atomic orbitals, critical for predicting molecular behavior and properties.

Molecular Orbital Diagram for O2

A diagram that illustrates the arrangement of molecular orbitals in O2 and indicates its magnetic property as paramagnetic.

Bond stregth ( Bond order and Bond dissociation energy)

The energy required to break a covalent bond in a molecule and is related to bond order and bond length.

BDE= change in enthalpy, when one mole of covalent gas phase is “ Broken” —> gas fragments

• Formation of bond: exothermic

• Breaking bond: endothermic

—> BDE varies with bond type( double, single or triple)

Homopolar Diatomic Molecules

Molecules composed of two identical atoms, such as O2 or N2, that have specific bonding properties based on molecular orbital theory.

Energy Levels of σ and π Orbitals

In different molecules, the relative energy levels of σ and π orbitals can change based on bonding interactions and atomic configurations.

Relationship between bond length, order, and strength

• Longer the bond length, weaker bond strength, and bond order decreases

• If same element bonds: bond increases, the bond dissociation energy increases, and the bond length decrease

( remember: giữa hai element giống nhau thì nếu atom nhỏ in size, thí sẽ có bond length nhỏ và BDE sẽ lớn)