Chemistry 1050- Chapter 10: Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory

VSPER Theory (Valence Shell Électron Pair Repulsion)

A model used to predict molecular shapes by assuming that electron groups around a central atom repel each other and arrange themselves as far apart as possible to minimize repulsion

Electron-Pair Geometry

The 3D arrangement of all electron groups (bonding and lone pairs) around a central atom, where the groups spread out to maximize space between regions of high electron density around the central atom and minimize repulsion, resulting in one of five basic shape

Molecular Geometry

Geometry describing only the placement of the atoms in the molecule

Polar Covalent Bonds

Formed between two atoms with different electronegativites and results in separation of charge giving rise to bond dipole moment.

Polar Molecule

A molecule that has an overall separation of charge, caused by the combination of its bond polarities and its molecular shape.

Diatomic Molecules (Only one bond)

Molecules made of two atoms where molecular polarity is determined entirely by the polarity of the single bond—if the bond is polar, the molecule is polar; if the bond is nonpolar, the molecule is nonpolar.

Polyatomic Molecule

The presence of polar bonds may or may not result in a polar molecule

Quantum Mechanics

A model that describes where electrons are most likely to be found around atoms using orbitals (s, p, d, f), and explains that atoms must use hybrid orbitals when forming certain molecular bonds.

Hybrid Atomic Orbitals

Blend or combination of two or more standard atomic orbitals

Valance Bond Theory

A theory that says covalent bonds form when two half-filled atomic orbitals from different atoms overlap and share their electrons, with stronger bonds resulting from greater orbital overlap.

Hybridization (mixing)

the mixing of two or more atomic orbitals on the same atom to form new hybrid orbitals that better match how electrons are distributed in bonded atoms; these hybrid orbitals have new shapes and energies but remain localized on a single atom.

Degenerate Oribitals

All with the same energy

Sigma Bonds

1) s-s orbital overlap

2) s-p orbital overlap

3) p-p head-on overlap

Pi Bonds

Side-to-side orbital overlap and tends to be less efficient than end-to-end orbital overlap

- bond in double bond generally easier to break than the s bond

Molecular Orbital Theory

An advanced bonding model where atomic orbitals combine to form molecular orbitals that spread out over the entire molecule; electrons occupy these orbitals (two per orbital with opposite spins), and the orbitals are formed by adding or subtracting atomic wave functions (LCAO), creating regions of high or low electron probability.