14. Bonding HL

Predict the shape and bond angles for species with five and six negative charge centers using the VSEPR theory (See BONDING SL table for this objective)

Describe Sigma and Pi bonds

Sigma bonds-

·         Caused by overlapping S or P orbitals

·         Parallel to Inter-nuclear axis

·         Most electron density along in region between

nuclei along the inter-nuclear axis

                                   

Pi bonds-

·         Caused by overlapping P orbitals

·         Perpendicular to inter-nuclear axis

·         Most electron density in regions above/below
the inter-nuclear axis

Single bonds composed of 1 sigma OR 1 pi bond

Double bonds composed of 1 sigma AND 1 pi bond

Triple bonds composed of 1 sigma and 2 pi bonds


Below: image of a C-C double bond and four C-H single bonds

Explain hybridization in terms of the mixing of atomic orbitals to form new orbitals for bonding

Identify and explain the relationships between Lewis structures, molecular shapes, and types of hybridization

SP hybridization exists in linear molecules. Example: Ethyne ( HC=CH )

SP² hybridization exists in trigonal planar molecules. Example: Ethene ( H₂C=CH₂ )

SP³ hybridization exists in tetrahydral molecules. Example: Methane (CH₄)

Describe the delocalization of pi electrons and explain how this can account for the structures of some species.

   

Resonance Structures- Molecules that can be represented by multiple Lewis structures due to the various placements of double and single bonds.

Resonance Hybrid- Lewis Structures drawn with delocalized electrons so that no one part is charged or inaccurately drawn