Chemical Bonds & Molecular Orbitals

• Chemical bonds involve shared electrons, which is favorable due to Coulomb's Law.
• In atoms, electrons are in atomic orbitals (1s, 2s, 2p, etc.).
• Bonds form through atomic orbital overlap, creating molecular orbitals; the number of orbitals is conserved.
• Two atomic orbitals overlapping results in two molecular orbitals.
• Constructive overlap between atomic orbitals leads to a bonding σ (sigma) molecular orbital.
• Electron density is between nuclei, stabilizing the molecule.
• Destructive overlap between atomic orbitals results in an antibonding σ* (sigma star) molecular orbital.
• There is no electron density between nuclei when destructive overlap occurs, destabilizing the bond.
• Unstable configurations arise when positively charged nuclei are close without "glue" electrons.

Molecular Orbital Diagram for H₂

• The molecular orbital diagram for H₂ depicts bonding and antibonding orbitals.
• Antibonding orbitals have high energy, while bonding orbitals have low energy.
• ΔE represents the energy difference between atomic and molecular orbitals.
• σ* is the lowest unoccupied molecular orbital (LUMO).
• σ is the highest occupied molecular orbital (HOMO).

Pi (π) Bonding

• Pi bonding occurs through "side-on" overlap of p orbitals.
• Occurs when neighboring atoms contain electrons in unhybridized p-orbitals, seen in alkenes and alkynes.
• Electron density is shared between neighboring p orbitals, forming a π bond.
• P orbitals must be aligned in the same plane for pi bonding.
• Atoms involved in pi bonding cannot freely rotate due to orbital overlap.
• Pi bonding isn't possible when p orbitals are at 90° to each other.
• Two pi bonds can exist on the same atom through side-on overlap of two sets of p orbitals at 90° to each other with alkynes as an example.

Bonding and Antibonding Orbitals

• The number of molecular orbitals must equal the number of atomic orbitals.
• Combining two atomic orbitals always yields one bonding and one antibonding molecular orbital.

Antibonding Orbitals

• In H₂, a third electron would occupy the antibonding orbital.
• Antibonding orbitals fill when electrons gain energy or with extra electrons.

Hybridization

• Second-row elements (B, C, N, O) possess four valence orbitals being one 2s and three 2p orbitals (px, py, pz).
• Atoms arrange in tetrahedral (~109°) or trigonal (~120°) formations, so bonding can't involve pure p-orbitals.
• Valence orbitals mix to form hybrid orbitals of both s and p character.
• Hybrid orbitals arrange to minimize electron repulsion.

Types of Hybridization

• sp³: 1 s orbital and 3 p orbitals lead to a tetrahedral geometry (109°).
• sp²: 1 s orbital and 2 p orbitals lead to a trigonal planar geometry (120°).
• sp: 1 s orbital and 1 p orbital lead to a linear geometry (180°).

sp³ Hybridization

• In CH₄, the four C-H bonds have identical lengths (1.50 Å) and bond angles (109.5°).
• The 2s orbital and three 2p orbitals hybridize into four identical sp³ orbitals, each with 25% s-character and 75% p-character.
• Four sp³ orbitals result in a "tetrahedral" orbital geometry.

sp² Hybridization

• Combining one s orbital with two p-orbitals creates three hybrid orbitals with 33% s-character and 66% p-character, and one unhybridized p-orbital.
• Each sp² orbital is 120° apart ("trigonal planar" geometry).
• It's the geometry for central atoms in trivalent boron compounds (e.g., BH₃), carbocations, and atoms with single pi bonds (e.g., alkenes).

sp Hybridization

• One s orbital hybridized with one p-orbital results in two hybrid orbitals 180° apart and two "leftover" unhybridized p orbitals.
• Each sp orbital pair is 180° apart ("linear" geometry).
• Found in the central atom of beryllium compounds (e.g., BeCl₂) and atoms with triple bonds.

Shortcut for Hybridization

• Add the number of surrounding atoms (A) and lone pairs (LP).
• If A+LP = 4, the hybridization is sp³.
• If A+LP = 3, the hybridization is sp².
• If A+LP = 2, the hybridization is sp.

Exception to Hybridization

• If an atom with a lone pair is adjacent to a pi bond, rehybridization occurs to put the lone pair in a p orbital, so it's sp² not sp³.

Orbital vs. Molecular Geometry

• In CH₄, the geometry of orbitals around the central atom ("orbital geometry") and the geometry of hydrogens ("molecular geometry") are identical.
• Molecules like NH₃ and H₂O have four electron pair domains arranged tetrahedrally around the central atom.
• Lone pairs are not bonded to an atom.
• While orbital geometry is tetrahedral, molecular geometry is trigonal pyramidal in NH₃ and bent in H₂O.
• Lone pairs compress bond angles.

Examples of Hybridization

Hybridization and Bond Strength

• Higher s-character in a bond results in a stronger and shorter bond.
• sp-hybridized C-H bonds are stronger compared to sp²-hybridized, and sp² stronger than sp³-hybridized C-H bonds.
• With C-C bonds, a sp³-sp bond is stronger than a sp³-sp² bond, while sp³-sp² is stronger than a sp³-sp³ C-C bond.

Learn about chemical bonds formed through shared electrons and atomic orbital overlap, which is energetically favorable because of Coulomb's Law. Explore molecular orbitals of Hydrogen (H₂) resulting from constructive and destructive interference. Understand the difference between bonding and antibonding orbitals.