Chapter 9 Molecular Geometries and Bonding Theories

Multiple Bonds

Formation of two π bonds in acetylene Fig 9.26

Fig 9.27 Formation of σ and π bonds in formaldehyde, CH 2 O Describing σ and π bonds in a molecule formaldehyde σ σ σ π

Sigma (  ) and Pi Bonds (  ) Single bond 1 sigma bond Double bond 1 sigma bond and 1 pi bond Triple bond 1 sigma bond and 2 pi bonds How many  and  bonds are in the acetic acid (vinegar) molecule CH 3 COOH? C H H CH O OH  bonds = = 7  bonds = 1

Molecular Orbital (MO) Theory If waves interact constructively, the resulting orbital is lower in energy: a bonding molecular orbital. If waves interact destructively, the resulting orbital is higher in energy: an antibonding molecular orbital. In MO theory, we invoke the wave nature of electrons

MO Theory In H 2 the two electrons go into the bonding molecular orbital. The bond order is one half the difference between the number of bonding and antibonding electrons: Bond order = ½ (no. of bonding e − – no. of antibonding e − ) Here: ½ (2-0) = 1

MO Theory In the case of He 2, the bond order would be: Therefore, MO theory predicts that He 2 does not exist, which we know to be true. Here: ½ (2-2) = 0 Fig 9.35

In the case of He 2 +, the bond order would be: MO Theory He 2 + ½ (2-1) = 1/2 Therefore, MO theory predicts that He 2 + does exist and it will be relatively stable

MO Theory – Second-Row Diatomics Consider only homonuclear diatomic molecules Number of MOs = number of AOs combined AOs combine most effectively with other AOs of similar energy The greater the overlap of AOs, the lower the energy of MO Each MO can hold max of 2 electrons (Pauli exclusion) Hund’s rule applies (same spin in degenerate orbitals)

MOs for Li 2 and Be 2 Fig 9.37 Energy-level diagram for the Li 2 molecule

 For atoms with both s and p orbitals, there are two types of interactions:  The p orbitals that are head to head overlap in  fashion.  The other two sets of p orbitals overlap in  fashion. MOs from 2p Atomic Orbitals Fig 9.38

There are both  and  bonding molecular orbitals and  * and  * antibonding molecular orbitals Diagram fits only O 2 and F 2 MO Theory – Second-Row Diatomics Fig 9.43

MO Theory The smaller p-block elements in the second period have a sizeable interaction between the s and p orbitals: This flips the order of the  and  molecular orbitals in O 2 and F 2 Fig 9.45 Fig 9.44

Fig 9.48 Paramagentism of O 2

Figure Fig 9.48 Paramagentism of O 2