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UNIT 8: ADVANCED BONDING
Valence Bond Theory (Linus Pauling & others - 1930s)
- Half-filled atomic orbitals (containing single electrons of opposite spins) OVERLAP to form covalent bonds, which results when: - (1) an orbital on one atom overlaps an orbital on a second atom - (2) the single electrons in each orbital combine to form an electron pair.
- This overlap leads to a DECREASE in energy of system (due to the atoms being widely separated) , stabilizing the system and forming a covalent bond
- Atoms reach optimum distance when the energy between them reaches its lowest ( most stable ) value due to the attractive and repulsive forces combined to create the lowest possible energy configuration.
of bonds formed by an atom is determined by the # of unpaired electrons
- Greater overlap = stronger bonds
- LIMITATIONS: CH 4
σ Bond π Bond
Orbital Overlap Head-on ( end-to-end ) Side-by-side of 2 p orbitals
Electron Density Along the internuclear axis Above and below the axis
Strength Stronger due to greater overlap
Weaker due to less effective overlap
Formation Priority Always forms first Forms after a σ bond in double/triple bonds
Examples SINGLE bonds (H₂, CH₄) C₂H₄ ( DOUBLE bond, 1 σ bond and 1 π bond) C₂H₂ ( TRIPLE bond, 1 σ bond and 2 π bonds)
Q: Identify the number of σ and π bonds contained in this molecule.
A: There are 6 σ C–H bonds and 1 σ C–C bond, for a total of 7 from the single bonds. There are 2 double bonds that each have a π bond in addition to the σ bond. This gives a total nine σ and two π bonds overall.
Q: Identify each illustration as depicting a σ or π bond:
A: (a) is a π bond with a node along the axis connecting the nuclei while (b) and (c) are σ bonds that overlap along the axis.
HYBRID ATOMIC ORBITALS
- Explains electron density around covalently bonded atoms. bond theory..
- Hybridization describes the changes in the atomic orbitals of an atom when it forms a covalent compound
- # of hybrid orbitals formed = # of atomic orbitals mixed
- The type of hybridization depends on the electron pair geometry around the central atom - LONE PAIRS COUNT AS A BOND
Hybridization Orbitals Mixed (# bonds)
Geometry Bond Angles Examples
sp 1s + 1p = 2 Linear 180° BeCl₂, CO₂, C₂H₂ (ethyne)
sp² 1s + 2p = 3 Trigonal planar 120° BF₃, SO₃, C₂H₄ (ethene)
sp³ 1s + 3p = 4 Tetrahedral 109.5° CH₄, NH₃, H₂O
MOLECULAR ORBITAL THEORY (Walther Kohn)
- Describes how atoms come together to form molecules.
- When atoms get close, their atomic orbitals (where their electrons "live") combine to form molecular orbitals, which belong to the whole molecule, not just one atom
- Equal numbers of bonding and antibonding electrons = unstable molecule
Types of Molecular Orbitals:
○ Bonding orbitals are formed when atomic orbitals combine constructively (adding together), creating a region where electrons can exist between the atoms and hold them together. ○ Antibonding orbitals are formed when atomic orbitals combine destructively (subtracting from each other), creating a region with no electrons between the atoms, which tries to push them apart.
BONDING ORBITALS ANTI-BONDING
Electrons are between the nuclei Electrons are outside the nuclei
Lower in energy (stabilizing) Higher in energy (destabilizing)
Help hold atoms together Work against bonding
σ (sigma) or π (pi) σ (sigma-star)* or π*
Bond Order Formula: Bond Order = [(Electrons in bonding MOs)−(Electrons in antibonding MOs)]/
Interpretation:
● Bond Order > 0 (+): Molecule is EXIST and is STABLE ● Bond Order <= 0 (0 or - ): Molecule is unstable and DNE ● Higher Bond Order = MOST STABLE ( Stronger bond and shorter bond length)
The molecular orbital energy diagram for O 2 predicts two unpaired electrons.
Bond order = (8 - 4)/2 = 2
Magnetic Properties:
- Paramagnetic Molecules: Have unpaired electrons and attracted to magnetic fields (e.g., O₂ with two unpaired electrons).
- Diamagnetic Molecules: Have all electrons paired and repelled by magnetic fields (e.g., N₂ and F₂)
The orbital energies and atomic radius decreases across the period while the effective nuclear charge increases.
- Molecular orbitals in solids are so closely spaced that they are described as bands.
