Valence Bond Theory Considers the interaction of separate atoms brought together as they form a molecule. Lewis structures Resonance considerations
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1 CHEM 511 chapter 2 page 1 of 11 Chapter 2 Molecular Structure and Bonding Read the section on Lewis dot structures, we will not cover this in class. If you have problems, seek out a general chemistry text. Things of which you should already know/be aware Lewis structures Resonance considerations o resonance averages bond characteristics (length, strength, order) o the energy of a resonance hybrid is lower than that of any single contributing structure o molecules and/or ions do not resonate! Oxidation states Hypervalence (seen for atoms in period 3 and higher) VSEPR model: predicts shape of covalent structures based on the number of regions of electron density around an atom. Should know electronic and molecular shapes and the distortions caused by lone pairs of electrons. & see-saw/distorted tetrahedral/disphenoidal Valence Bond Theory Considers the interaction of separate atoms brought together as they form a molecule. EX. Simplest case is H2
2 CHEM 511 chapter 2 page 2 of 11 Ex. More complex case: O2 Diamagnetic vs. paramagnetic species Diamagnetic: a material repelled by a magnetic field (why?) Paramagnetic: a material attracted to a magnetic field (why?) Rationalization of Valence Bond and VSEPR There are some problems with these theories 1. Bond angles don't match expectation EX. H2O 2. Number of bonds formed doesn't always match prediction EX. Carbon 3. Bond strength and length is not always predicted accurately EX. Carbon (with promotion) Hybridization How can VSEPR account for all of these differences? Mix orbitals together to create new orbitals.
3 CHEM 511 chapter 2 page 3 of 11 Molecular Orbital Theory A way of describing the bonding characteristics of a molecule Instead of electrons being localized in atomic orbitals, molecular-sized orbitals (orbitals covering the entire molecule) are involved Types of Orbitals 1. Bonding orbitals (+) -positive overlap of atomic orbitals (AO) -when filled with electrons, the energy of the molecule is lowered relative to that of the separated atoms 2. Antibonding orbitals (-) -when filled with electrons, the energy of the molecule is raised relative to that of the separated atoms -antibonding orbitals are more destabilizing than a bonding orbital is stabilizing 3. Nonbonding orbitals (n) -generally are not involved in bonding and are usually localized on a particular atom -often occurs because no atomic orbitals of the correct symmetry can overlap Further orbital distinctions -orbital -formed from head-to-head overlap of orbitals -will have cylindrical symmetry -orbital -formed from the sideways overlap of orbitals -will have one nodal plane and the sign of ψ will change with a 180 rotation -orbital -formed from the sideways overlap of d-orbitals -will have 2 nodal planes
4 CHEM 511 chapter 2 page 4 of 11 MO diagram for H2 Bond order: this is equivalent to a single, double, triple, etc. bond in VB theory. # of bonding e BO # of antibondin g e 2 MO and BO for He2? Rules for diatomic molecules 1. Using n AO's, you will create n MO's 2. For period 2 molecules there are 4 and 4 orbitals 3. The orbitals form 1 doubly degenerate pair of bonding orbitals and 1 doubly degenerate pair of antibonding orbitals 4. The orbitals form 1 strongly bonding, 1 strongly antibonding, and two in between 5. If the orbital has a center of inversion designate with a g (gerade, even) or u (ungerade, odd) 6. To determine the actual energies of the orbitals, experimental evidence must be given (rigorous computation may get you close, but it should be confirmed with photoelectron)
5 CHEM 511 chapter 2 page 5 of 11 MO for Li2 through N2 EX. Bond order for N2? EX. Bond order for F2?
6 CHEM 511 chapter 2 page 6 of 11 Evidence for MO theory? Photoelectron spectroscopy! Aufbau Principle for Molecules Follows the same principles as for atoms lower energy orbitals fill first obey Hund's rule obey the Pauli exclusion principle What is the electron configuration of Li2? Be2? B2? F2? Important distinctions of orbitals HOMO: highest occupied molecular orbital LUMO: lowest unoccupied molecular orbital SOMO: singly occupied MO
7 CHEM 511 chapter 2 page 7 of 11 Heteronuclear diatomic molecules Since 2 different atoms, electronegativities will differ The more electronegative atom will contribute more to the bonding MO The less electronegative atoms will contribute more to the antibonding MO Bonding in HF Electron configuration of H? Electron configuration of F?
8 CHEM 511 chapter 2 page 8 of 11 EX. Bond order for CO? Electron configuration of C? of O? Bond Correlations Bond length is inversely proportional to bond order Bond enthalpy (strength) increases as bond order increases Bond Bond strength (kj/mol) Relative strength N-N 163 x N=N x N N x P-P 201 x P=P x P P x C-C 348 x C=C x C C x
9 CHEM 511 chapter 2 page 9 of 11 Molecular Orbitals of Polyatomic Molecules New nomenclature for orbital overlaps of non-linear molecules a, b = non-degenerate orbitals (in place of ) e = doubly degenerate orbitals (in place of ) t = triply degenerate orbitals Bonding in NH3 Use nitrogen atomic orbitals and a triangular H3 as a basis set Recall: the number of AO's = number of MO's Hypervalence with MOs Consider SF6 Note: I won't ask something this complicated for you to construct. Note: It does show how MO theory can be used to "expand the octet" without using d-orbitals (So why don t 2nd period atoms undergo hypervalence?)
10 CHEM 511 chapter 2 page 10 of 11 Electron deficient compounds: Molecules or ions that contain too few electrons to allow their bonding to be described exclusively in terms of two-center, two-electron bonds (IUPAC Gold Book) EX. B2H6 Contains 2 electron, 3 center bonds Bond lengths Covalent radius: contribution of an atom in a covalent bond. Orbitals must overlap! van der Waals radius: size of atoms near a molecule, but not bonded with that molecule Bond strength Can measure bond dissociation energy (ΔH o A-B or B) Enthalpy for the dissociation of all bonds in an isolated, gaseous molecule into gaseous, isolated atoms
11 CHEM 511 chapter 2 page 11 of 11 Electronegativity and bond type From Chapter 1 we saw Pauling calculated electronegativity derived from ideas relating to bond strength. First, let s calculate the Pauling electronegativity of hydrogen based on the following bond energies: H-H = 436 kj/mol F-F = 155 kj/mol H-F = 565 kj/mol Instead of a simple difference in electronegativity to determine covalent/ionic character, a Ketelaar triangle is used (this was fully developed by Gordon Sproul at USC Beaufort) Classify the following bonds according to the figure above Bond Δ avg Br-C (2.96 vs 2.55) Br-Li (2.96 vs 0.98)
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