Covalent Bonding: Orbitals

1. Chapter 9 Covalent Bonding: Orbitals

2. Types of bonds Sigma bonds Overlap of orbitals from two atoms Between the atoms Pi bonds Above and below atoms Between p orbitals Second bond of all double bonds

3. Hybridization and Electrons Atomic orbitals don’t work to explain molecular geometry (methane - CH4) Shape is tetrahedral Carbon has two s and two p orbital electrons to bond with hydrogen These bonds and orbitals are identical Atomic orbitals change when making a new molecule

4. Hybridization s and p orbitals can be blended together Blend of s and p orbitals can combine in a variety of manners depending on the number of pairs present Dictates the molecular geometry of a molecule

5. sp3 Hybridization s and p orbitals of two atoms overlap One s orbital and three p orbitals give a tetrahedral shape Example is methane

6. sp2 Hybridization Occurs with three bonding pairs with bond arrangements of 120 degrees Has trigonal planar geometry One of the p orbital electrons is shared with another p orbital (sigma bond) Double bonds always contain a sigma bond and a pi bond Example is C2H4 (ethylene)

7. sp Hybridization Involves one s orbital and one p orbital Two bonding pairs are arranged at 180 degrees Produces a linear molecule Examples are CO2 and C2H2

8. 2p sp Hybridization In terms of energy 2p Energy 2s

9. d level Hybridizations dsp3 hybridizations Molecules with five sets of electron pairs (Exceeds octet rule) Follows the trigonal bipyramidal geometry d2sp3 hybridizations Molecules with six sets of electron pairs (Exceeds octet rule) Follows the octahedral geometry

10. Hybridization summary Steps for using the localizied electron model (hybridizated orbitals) Draw the Lewis structure. Determine the arrangement of electron pairs using the VSEPR model. Specify the hybrid orbitals needed to accommodate the electron pairs.

11. Molecular Orbital Model Similar to quantum mechanical model of atoms Molecular orbitals (MO) are similar to atomic orbitals Each MO can hold two electrons, but they must have opposite spins MO1 = 1sA + 1sB MO2 = 1sA - 1sB

12. The Molecular Orbital Model In the molecule only the molecular orbitals exist, the atomic orbitals are gone MO1 is lower in energy than the 1s orbitals they came from. This favors molecule formation Called an bonding orbital MO2 is higher in energy This goes against bonding (less stable) Called an antibonding orbital

13. The Molecular Orbital Model MO2 Energy 1s 1s MO1

14. The Molecular Orbital Model The molecular orbitals are centered on a line through the nuclei MO1 is the greatest probability of electrons between the nuclei. MO2 is the electrons on either side of the nuclei. This shape is called a sigma molecular orbital

15. The Molecular Orbital Model We use labels to indicate whether the MO’s are bonding or antibonding. MO1 =s1s MO2 =s1s* (* indicates antibonding) Can write them the same way as atomic orbitals H2 =s1s2 New molecular orbitals are formed (no longer atomic orbitals) - have lower energy associated with them

16. Bond Order Concept to indicate bond strength Larger bond order, the greater the bond strength Bond energy is the difference between the number of bonding electrons and the number of antibonding electrons divided by 2. Use 2 because bonding occurs in pairs H2 versus H2- example

17. Paramagnetism Paramagetism causes a substance to be attracted into a magnetic field Associated with unpaired electrons Diamagnetism cause the substance to be repelled from the magnetic field Associated with paired electrons