Covalent Bonding: Orbitals Chapter 9 Covalent Bonding: Orbitals

9.1: I. Hybridization A. Bond only includes valence e- B. When alone, an atom has orbitals based on where its e- need to be for lowest energy and least amount of repulsions C. As a molecule, atoms form a mixture of orbitals (“hybrid”) from its valence orbitals

II. Carbon in CH4 A. Carbon has s and p valence orbitals that form sp3 hybrid orbitals when making 4 single bonds B. Called sp3 because there is one s and three p orbitals involved, bond angle of ~109º C. Carbon is “sp3 hybridized” in CH4

III. sp2 Hybridization A. Carbon in C2H4 has one double and 2 single bonds B. Carbon with three bonds is sp2 hybridized C. Uses one s and two p orbitals to give trigonal planar structure,120º angle D. Remaining p orbitals are oriented perpendicular to double bond to minimize electron repulsions

E. Overlapping orbitals are sigma (σ) orbitals with sigma (σ) bond F. Parallel orbitals are pi () orbitals with pi () bond G. Need a sigma and pi bond to form double bond

IV. SP Hybridization A. Carbon in CO2 has 2 effective bonds B. Involves one s and one p (“sp” orbital) with 180º bond angle C. Oxygen has sp2 hybridization in CO2 (1 bond, 2 lone pairs)

V. sp3d, sp3d2 Hybridization A. When atoms exceed octet rule, we need more than s and p orbitals B. PCl5 forms 5 bonds so we use sp3d hybridization C. Sulfur in SF6 has 6 bonds, sp3d2 hybridization

9.2: I. Molecular Orbital Model A. Unlike hybrid orbitals which are combinations of an atoms own orbitals (s,p, or d), molecular orbitals are combinations between the orbitals of two atoms B. Mathematically determined locations where electrons involved in molecules exist C. Get two different types: “bonding” (when e- between atoms) and “anti-bonding” (lone pairs or e- not shared)

II. Example H2 A. H has 1s orbital, combined with itself you get: MO1 = 1sA + 1sB MO2 = 1sA – 1sB B. Bonding MO has a sigma bond (σ) and is lower energy than separate atoms (more stable) C. Anti-bonding MO has repelling force, is called σ* and has higher energy than separate atoms

III. Molecular Orbital Rules Sigma bonds between molecular orbitals in same plane Pi bonds between parallel molecular orbitals Bonding orbitals always lower energy than anti-bonding Anti-bonding orbitals indicated by asterisk Molecular e- configuration written same way as atomic e- config. 6. Each molecular orbital can hold 2 e-s (opposite spin) 7. # of molecular orbitals always same as number of atomic orbitals used to create them

IV. Molecular Orbital Energy Diagram A. Shows electron placement as separate atoms and as joined atoms, includes both bonding and anti-bonding molecular orbitals

V. Using MO Energy Diagrams A. Fill in e- in lowest energy states first when combined from both atoms, extra e- go in anti-bonding state

VI. Bond Order A. Indicates relative strength of bond B. Difference between number of bonding e- and anti-bonding e- divided by two C. B.O. = (# bonding “e-” - # of anti-bonding “e-”)/2 D. Larger number = stronger bond E. H2 = 1, H2+ = ½, H2- = ½ F. Helium’s bond order bonded to itself is 0, so it is not as stable as separate atoms which is why it doesn’t form a pair

9.3: I. Homonuclear Diatomic Molecules A. Diatomic molecules are made of two identical atoms (H2, O2, Cl2, F2, etc.) B. We only consider the valence e- because the inner shells are not close enough to the other atom to form molecular orbitals C. Any molecule with only s orbitals in the valence shell will interact similarly to H2

II. What Happens to P Orbitals? A. P orbitals interact in two ways depending on their orientations Sigma (σ) Bond Pi (π) Bond

III. P Molecular Orbital Energy Diagrams A. P orbitals that align form sigma bonds, P orbitals that are parallel form pi bonds B. We need energy levels for the sigma (σ) bond, anti-sigma (σ*) bond, pi () bond, and anti-pi (*) bond C. This is the energy diagram for just the p orbitals (we can’t put any s orbital e- here) D. Each line holds two e- just like each orientation of the p orbitals

IV. Bond Order for P orbital A. To determine bond order for something with a p orbital, we also have to consider s orbitals in terms of bonding vs. anti-bonding electrons Ex. Boron Bond order = (bonding electrons–non-bonding)/2 B.O. = (4 – 2)/2 = 1, stable atom, higher numbers better

V. Paramagnetism/Diamagnetism A. Most materials aren’t magnetic, but when in a magnetic field, two types of magnetism can result: Paramagnetism: substance attracted to whatever causes magnetic field Diamagnetism: substance repelled from source B. If a substance has both paired and unpaired electrons, the paramagnetic effect will usually over-power diamagnetism ***Paramagnetism is related to unpaired electrons, diamagnetism related to paired electrons***

VI. Paramagnetic or Diamagnetic? Boron Carbon

VII. Bonding Summary A. As bond order increases, bond energy increases and bond length decreases B. Bond orders cannot be used definitively for bond strength because e- repulsion can weaken the strength C. Bond order is somewhat related to # of covalent bonds D. B.O.= 1 is single bond, 2 is double bond, 3 is triple bond, ½ numbers are in between E. The molecular orbital theory better at explaining paramagnetism and diamagnetism than the localized e- model

9.4: I. Heteronuclear Diatomic Molecules A. Different atoms bonded B. Can use the homonuclear energy diagram for atoms near each other on periodic table (ex. N, O) C. For farther apart atoms we need different energy chart

9.5: I. Combining Localized and Molecular Orbital Models A. Resonance structures for a molecule are different possible e- positions that actually occur simultaneously B. Combining these two models eliminates the need for the idea of resonance by assuming that e- are not localized (in isolated positions)

II. Sigma and Pi bonds in Resonance A. In resonance structures, the only aspect of the molecule that changes is the placement of double bonds B. Since double bonds are made of a sigma and a pi bond, the only thing actually changing in resonance structures is the pi bond placement