| Lecture
#29 |
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| CURMUDGEON
GENERAL'S WARNING. These "slides"
represent highlights from lecture and are neither
complete nor meant to replace lecture. It is
advised not to use
these as a reliable means to replace missed
lecture material. Do so at risk to healthy
academic performance in 09-105. |
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Exam III Review
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Review for Exam III |
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| The following arrangement of flasks containing
methane (MW = 16.0 g/mol), propane (MW = 44.1 g/mol) and
ethane (MW = 30.0 g/mol) is set up at 25 C. Assuming
constant temperature, zero volume for the connecting
tubes and ideal gas behavior... Determine the total
pressure after the valves are open for a long time.
What is the final partial pressure of ethane (C2H6)
in the left flask?
What is the molar compressibility of the final gas
mixture?
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| Reviewing some interesting odds and ends before EXAM
III, let's look at why di-carbon does not have a
quadruple bond, despite the contrary indications of the
Lewis structure. |
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| The more rigorous molecular orbital theory approach
to chemical bonding shows that di-carbon has a double
bond, both being pi bonds, and that there are lone pairs
of 2s electrons at each carbon atom. |
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| An exotic excited state of di-carbon can be
hypothesized in which the total bond order is now 4, but
actually making such a beast is unlikely to be successful
in practice. |
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| Reviewing heteronuclear diatomic molecules and the
atomic orbitals from which molecular orbitals are
constructed, we examined the HCl molecule as shown. |
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| Some heteronuclear diatomic molecules, not
surprisingly, differ only slightly in their participating
molecular orbitals from what we've seen for homonuclear
diatomic molecules. |
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| If we add an electron to ozone to produce the ozonide
ion, what happens to the total bond order? Bond length?
Bond strength? |
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| Here is an underlying feature of a large collection
of molecules of extreme importance. There is a ring of
delocalized electrons, shown as the red alternating
(conjugated) single bond double bond sequence. |
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| A very simple model conceptually identical to that
for the particle-in-a-box can be applied to electrons
delocalized about a ring. The formula for the allowed
energy levels is reminiscent of the one we had for chains
of delocalized electrons and the calculations of
transition energies follow exactly the same line of
reasoning. The ring structure shown absorbs light of
approximately 600 nm wavelength. |
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| Constructing molecular orbitals starts with a look at
the alignment of the valence atomic orbitals. |
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| Polyatomic orbitals starting from a Lewis structure. |
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