Molecular Orbital Theory Diagrams to Know for General Chemistry II

Molecular Orbital Theory Diagrams help us understand how atoms bond and form molecules. By examining bonding and antibonding orbitals, we can predict molecular stability, reactivity, and properties like magnetism in various molecules, from diatomic to complex structures.

1. Homonuclear diatomic molecules (H2, N2, O2, F2)

   • Formed from two identical atoms, leading to symmetrical molecular orbitals.
   • Bonding and antibonding orbitals are created from the combination of atomic orbitals.
   • The bond order can be calculated to determine the stability of the molecule (e.g., H2 has a bond order of 1, N2 has a bond order of 3).
   • Molecular orbital diagrams show the energy levels of bonding and antibonding orbitals.
   • The presence of unpaired electrons in O2 indicates it is paramagnetic.

2. Heteronuclear diatomic molecules (CO, NO)

   • Composed of two different atoms, leading to unequal sharing of electrons.
   • The molecular orbital energy levels are influenced by the electronegativity of the atoms.
   • Bonding and antibonding orbitals still form, but their energy levels differ from homonuclear molecules.
   • The bond order can also be calculated, providing insight into the stability of the molecule.
   • Examples like CO show a strong triple bond and significant dipole moment due to polarity.

3. Linear triatomic molecules (BeH2, CO2)

   • Molecules with three atoms arranged in a straight line.
   • In CO2, the linear shape results from the double bonds between carbon and oxygen.
   • The molecular orbital theory helps explain the bond angles and hybridization (sp hybridization in CO2).
   • BeH2 exhibits sp hybridization with two Be-H sigma bonds.
   • Linear triatomic molecules often have a bond angle of 180°.

4. Angular triatomic molecules (H2O)

   • Molecules with three atoms arranged in a bent shape.
   • The angle between the hydrogen atoms in H2O is approximately 104.5° due to lone pair repulsion.
   • Molecular orbital theory explains the hybridization (sp3) and the formation of sigma bonds.
   • The presence of lone pairs affects the molecular geometry and bond angles.
   • H2O is polar, leading to significant hydrogen bonding.

5. Tetrahedral molecules (CH4)

   • Molecules with four atoms arranged around a central atom in a tetrahedral shape.
   • The bond angles are approximately 109.5°, resulting from sp3 hybridization of the central carbon atom.
   • All four C-H bonds are sigma bonds formed from the overlap of atomic orbitals.
   • Tetrahedral geometry minimizes electron pair repulsion.
   • CH4 is nonpolar due to its symmetrical shape.

6. Octahedral molecules (SF6)

   • Molecules with six atoms symmetrically arranged around a central atom.
   • The bond angles are 90°, resulting from sp3d2 hybridization of the central sulfur atom.
   • All bonds in SF6 are sigma bonds, formed from the overlap of atomic orbitals.
   • Octahedral geometry allows for maximum separation of electron pairs.
   • SF6 is nonpolar due to its symmetrical arrangement.

7. Square planar molecules (XeF4)

   • Molecules with four atoms arranged in a square plane around a central atom.
   • The bond angles are 90°, resulting from dsp2 hybridization of the central xenon atom.
   • Two lone pairs on the central atom occupy axial positions, minimizing repulsion.
   • Square planar geometry is a result of the arrangement of bonding and non-bonding electron pairs.
   • XeF4 is polar due to the presence of electronegative fluorine atoms.

8. Pi-bonding in organic molecules (ethylene, benzene)

   • Pi bonds form from the side-to-side overlap of p orbitals.
   • In ethylene (C2H4), a double bond consists of one sigma and one pi bond.
   • Benzene (C6H6) features resonance, where pi electrons are delocalized over the ring structure.
   • Pi-bonding contributes to the stability and reactivity of organic compounds.
   • The presence of pi bonds affects molecular geometry and hybridization (sp2 in ethylene, sp2 in benzene).

9. Antibonding and bonding orbitals

   • Bonding orbitals are lower in energy and stabilize the molecule, while antibonding orbitals are higher in energy and destabilize it.
   • Electrons in bonding orbitals contribute to bond formation, while those in antibonding orbitals can weaken or break bonds.
   • The filling of molecular orbitals follows the Aufbau principle, Hund's rule, and the Pauli exclusion principle.
   • The bond order can be calculated using the difference between the number of electrons in bonding and antibonding orbitals.
   • Understanding these concepts is crucial for predicting molecular stability and reactivity.

10. Molecular orbital energy level diagrams

    • Diagrams visually represent the relative energy levels of molecular orbitals.
    • They illustrate the filling of orbitals with electrons and help determine bond order and magnetic properties.
    • The arrangement of orbitals reflects the symmetry and type of bonding in the molecule.
    • Diagrams differ for homonuclear and heteronuclear molecules due to variations in atomic orbital energies.
    • Analyzing these diagrams is essential for understanding molecular structure and behavior in chemical reactions.