Electron-Deficient and Odd-Electron Molecules

Electron-Deficient Molecules

• Contain either Be or B as the central atom
• Have fewer than 8 electrons around the central atom (4 e- around Be and 6 e- around B)
• Examples: BF3

Odd-Electron Molecules

• Have an odd number of valence electrons
• Examples: free radicals, NO2, CH3, OH, H, etc.
• Contain a lone (unpaired) electron and are paramagnetic

Expanded Valence Shells

• Occur when a central atom has more than 8 electrons around it
• Use empty outer d orbitals
• Occurs only with central atoms from Period 3 or higher
• Examples: SF6, PCl5, H2SO4, H3PO4

VSEPR Theory

• Enables prediction of molecular structure and approximate bond angles
• Assumes electron pairs in the valence shell of a central atom will adopt an arrangement that minimizes repulsions
• Electrons form either bonding pairs or lone pairs
• Electron pairs arrange to maximize distance between them

Octet Rule

• States that atoms tend to gain, lose, or share electrons to attain the electronic configuration of the nearest noble gas
• Applies to all types of bonding
• Atoms tend to achieve an octet (8) electrons in their highest occupied energy level

Bonding

• The joining of at least two atoms to form a molecule or compound
• Involves the valence shell electrons
• A chemical bond is a mutual electrical attraction between nuclei and valence electrons

Ionic Radii

• Increase down a group in the periodic table
• Decrease across a period in the periodic table
• Example: Al3+ < Ca2+

Electron Affinity (EA)

• Defined as the energy released when an electron is added to an isolated gaseous atom
• Can be either endothermic or exothermic
• Negative values indicate energy is released, while positive values indicate energy is required

Electro-negativity (EN)

• Measures the relative tendency of an atom to attract electrons
• Elements with high electro-negativities (nonmetals) often gain electrons to form anions
• Example: F (4.0) > O (3.5) > Cl=N (3.0)