1	Chapter 9 Orbitals and Covalent Bond
2	Molecular Orbitals The overlap of atomic orbitals from separate atoms makes molecular orbitals Each molecular orbital has room for two electrons Two types of MO Sigma ( s ) between atoms Pi ( p ) above and below atoms
3	Sigma bonding orbitals From s orbitals on separate atoms
4	Sigma bonding orbitals From p orbitals on separate atoms
5	Pi bonding orbitals p orbitals on separate atoms
6	Sigma and pi bonds All single bonds are sigma bonds A double bond is one sigma and one pi bond A triple bond is one sigma and two pi bonds.
7	Atomic Orbitals Don’t Work to explain molecular geometry. In methane, CH₄, the shape is tetrahedral. The valence electrons of carbon should be two in s, and two in p. the p orbitals would have to be at right angles. The atomic orbitals change when making a molecule
8	Hybridization We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry. We combine one s orbital and 3 p orbitals. sp³ hybridization has tetrahedral geometry.
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11	In terms of energy
12	How we get to hybridization We know the geometry from experiment. We know the orbitals of the atom hybridizing atomic orbitals can explain the geometry. So if the geometry requires a tetrahedral shape, it is sp³ hybridized This includes bent and trigonal pyramidal molecules because one of the sp³ lobes holds the lone pair.
13	sp² hybridization C₂H₄ Double bond acts as one pair. trigonal planar Have to end up with three blended orbitals. Use one s and two p orbitals to make sp² orbitals. Leaves one p orbital perpendicular.
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16	In terms of energy
17	Where is the P orbital? Perpendicular The overlap of orbitals makes a sigma bond (s bond)
18	Two types of Bonds Sigma bonds from overlap of orbitals. Between the atoms. Pi bond (p bond) above and below atoms Between adjacent p orbitals. The two bonds of a double bond.
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20	sp² hybridization When three things come off atom. trigonal planar 120º One p bond, s + lp =3
21	What about two When two things come off. One s and one p hybridize. linear
22	sp hybridization End up with two lobes 180º apart. p orbitals are at right angles Makes room for two p bonds and two sigma bonds. A triple bond or two double bonds.
23	In terms of energy
24	CO₂ C can make two s and two p O can make one s and one p
25	N₂
26	N₂
27	Breaking the octet PCl₅ The model predicts that we must use the d orbitals. dsp³ hybridization There is some controversy about how involved the d orbitals are.
28	dsp³ Trigonal bipyrimidal can only s bond. can’t p bond. basic shape for five things.
29	PCl₅
30	d²sp³ gets us to six things around Octahedral Only σ bond
31	Molecular Orbital Model Localized Model we have learned explains much about bonding. It doesn’t deal well with the ideal of resonance, unpaired electrons, and bond energy. The MO model is a parallel of the atomic orbital, using quantum mechanics. Each MO can hold two electrons with opposite spins Square of wave function tells probability
32	What do you get? Solve the equations for H₂ H_A H_B get two orbitals MO₂ = 1s_A - 1s_B MO₁ = 1s_A + 1s_B
33	The Molecular Orbital Model The molecular orbitals are centered on a line through the nuclei MO₁ the greatest probability is between the nuclei MO₂ it is on either side of the nuclei this shape is called a sigma molecular orbital
34	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 antibonding orbital
35	The Molecular Orbital Model
36	The Molecular Orbital Model We use labels to indicate shapes, and whether the MO’s are bonding or antibonding. MO₁ = s_1s MO₂ = s_1s* (* indicates antibonding) Can write them the same way as atomic orbitals H₂ = s_1s²
37	The Molecular Orbital Model Each MO can hold two electrons, but they must have opposite spins Orbitals are conserved. The number of molecular orbitals must equal the number atomic orbitals that are used to make them.
38	H₂^-
39	Bond Order The difference between the number of bonding electrons and the number of antibonding electrons divided by two
40	Only outer orbitals bond The 1s orbital is much smaller than the 2s orbital When only the 2s orbitals are involved in bonding Don’t use the s_1s or s_1s* for Li₂ Li₂ = (s_2s)² In order to participate in bonds the orbitals must overlap in space.
41	Bonding in Homonuclear Diatomic Molecules Need to use Homonuclear so that we know the relative energies. Li₂^- (s_2s)²(s_2s)¹ Be₂ (s_2s)²(s_2s)² What about the p orbitals? How do they form orbitals? Remember that orbitals must be conserved.
42	B₂
43	B₂
44	Expected Energy Diagram
45	B₂
46	B₂ (s_2s)²(s_2s*)²(s_2p)² Bond order = (4-2) / 2 Should be stable. This assumes there is no interaction between the s and p orbitals. Hard to believe since they overlap proof comes from magnetism.
47	Magnetism Magnetism has to do with electrons. Remember that spin is how an electron reacts to a magnetic field Paramagnetism attracted by a magnet. associated with unpaired electrons. Diamagnetism repelled by a magnet. associated with paired electrons. B₂ is paramagnetic.
48	Magnetism The energies of of the p_2p and the s_2p are reversed by p and s interacting The s_2s and the s_2s* are no longer equally spaced. Here’s what it looks like.
49	Correct energy diagram
50	B₂
51	Patterns As bond order increases, bond energy increases. As bond order increases, bond length decreases. Supports basis of MO model. There is not a direct correlation of bond order to bond energy. O₂ is known to be paramagnetic. Movie.
52	Magnetism Ferromagnetic strongly attracted Paramagnetic weakly attracted Liquid Oxygen Diamagnetic weakly repelled Graphite Water Frog
53	Examples C₂ N₂ O₂ F₂ P₂
54	Heteronuclear Diatomic Species Simple type has them in the same energy level, so can use the orbitals we already know. Slight energy differences. NO
55	NO
56	You try NO⁺ CN^- What if they come from completely different orbitals and energy? HF Simplify first by assuming that F only uses one if its 2p orbitals. F holds onto its electrons, so they have low energy
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58	Consequences Paramagnetic Since 2p is lower in energy, favored by electrons. Electrons spend time closer to fluorine. Compatible with polarity and electronegativity.
59	Names sp orbitals are called the Localized electron model s and p Molecular orbital model Localized is good for geometry, doesn’t deal well with resonance. seeing s bonds as localized works well It is the p bonds in the resonance structures that can move.
60	p delocalized bonding C₆H₆
61	C₂H₆
62	NO₃^-