Chapter 7.3-7.4 – Triple Bonds and Molecular Orbital Theory CHM1111 Section 04 Instructor: Dr. Jules Carlson Class Time: M/W/F 1:30-2:20 Friday, November 4th

Look at Chloroethylene Also, work through example 7-6 on p. 377 Chloroethylene H2CCHCl Carbons hybridize sp2 H bonds with 1s, Cl with 3p

Look at Acetic Acid Acetic acid: CH3COOH Inner atoms hybridize as shown H bonds with 1s sp3 sp3 Outer O bonds with 2p Carbon bound to outer O bonds with sp2 (σ) bond and p (π) bond. sp2

Carbon vs. Silicon Carbon forms π bonds readily, but Silicon does not. Bond length in Si bonds is larger – reduced stability of π bonds.

Other atoms that form π bonds π bonds can be formed by second row and upper row elements as inner atoms. In this course we will only look at species in second row + a class called oxoanions. Examples of oxoanions: Nitrate NO3-, Phosphate PO43-, Carbonate CO32-, Sulfate SO42-, Chlorate ClO3-

Bonding in Ethyne Each carbon atom has a steric number of 2. Hybridizes sp with 2 remaining p orbitals.

Bonding Overlap Example Construct a complete bonding picture for Hydrogen Cyanide - HCN and sketch the various orbitals.

I Clicker Question Which of the following statements are true about butenoic acid, CH3CHCHCOOH? It has 3 π bonds. The bonding between CH and CH is 2sp2 hybridized. The σ bond in C=O is a 2sp2-2sp2 bond. Both (a) and (b) are true. (a), (b), and (c) are true.

Molecular Orbital Theory The orbital overlap model, VSEPR, and hybridization, which all fit under Localized Bonding Theory (also called Valence Bond Theory), do an excellent job of rationalizing and predicting chemical structures. Electrons can be delocalized and spread out over several atoms, and described by molecular orbitals (MOs). Advantages of MOs over Localized Bonding Theory: - predicts magnetic properties - bond energy trends well described. - correctly explains electronic structures of molecules which do not follow Lewis Dot structure.

Molecular Orbitals of H2 Orbital overlap for H2 is between overlapping spherical 1s orbitals. Since electrons have wave-like properties, we can add or subtract the wavefunctions used to describe their orbitals. If you sum orbitals, their amplitudes may be added or subtracted depending upon the sign of the waves. Orbital Overlap: As two hydrogen atoms approach each other, the overlap of their 1s atomic orbitals increases. The wave amplitudes add, generating a new orbital with high electron density between the nuclei.

Molecular Orbitals of H2 Bonding molecular orbital: an additive interaction when high electron density results from more attraction than repulsion Antibonding molecular orbital: a subtractive interaction when low electron density results from more repulsion than attraction. Bonding orbital referred to as σ1s and antibonding orbital referred to as σ1s*

Types of Orbitals σ orbitals – electron density lies along the bond axis. π orbitals – electron density lies above and below the bond axis. Can have σ bonding orbitals and σ* antibonding orbitals Can also have π bonding orbitals and π* antibonding orbitals Antibonding orbitals are always higher in energy than bonding orbitals.

Filling and Energy of Molecular Orbitals for H2 H2 forms σ1s and σ1s* orbitals Aufbau principle, Pauli Exclusion principle and Hund’s rule all still apply for electron filling. Remember antibonding orbitals are higher in energy than bonding orbitals.

Rules for Building and Filling Molecular Orbitals The total number of molecular orbitals produced by a set of interacting atomic orbitals is always equal to the number of interacting atomic orbitals. The bonding molecular orbital is lower in energy than the parent orbitals and the antibonding orbital is higher in energy. Electrons of the molecule are assigned to orbitals of successively higher energy according to the aufbau principle and Hund’s Rule.

Bond Order Bond order – describes the relative effect of bonding and antibonding orbitals. Bond Order = ½(number of electrons in bonding MOs – number of electrons in antibonding MOs). Bonding is stable if Bond Order > 0. He2 Bond Order = 0 H2 Bond Order = 1

Second-Row Diatomic Molecules For second row diatomics (i.e. O2) Additive and subtractive overlaps need to be considered. Additive overlap of p σ bonds Subtractive overlap of p σ bonds Additive overlap of p π bonds Subtractive overlap of p π bonds