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1/9/2015
1
In the mid 1800s, it was first suggested that substances
are defined by a specific arrangement of atoms.
Why is a compounds formula NOT adequate to define it?
What term do we use to describe different substances
with the same formula?
Chapter 1 Structural Theory
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-1
Atoms that are most commonly bonded to carbon
include N, O, H, and halides (F, Cl, Br, I).
With some exceptions, each element generally forms a
specific number of bonds with other atoms:
Practice with SKILLBUILDER 1.1.
1.2 Structural Theory
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-2
How do potential energy and stability relate?
What forces keep the bond at the optimal
length?
1.3 Covalent Bonding
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-3
For simple Lewis structures:
1. Draw the individual atoms using dots to represent the
valence electrons.
2. Put the atoms together so they share PAIRS of electrons to
make complete octets. WHAT is an octet?
Take NH3, for example:
Practice with SKILLBUILDER 1.3.
1.3 Covalent Bonding Simple Lewis
Structures
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-4
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2
Consider the formal charge example below. Calculate the formal
charge on each atom.
Carbon SHOULD have 4 valence electrons because it is in Group IVA on the
periodic table.
Carbon ACTUALLY has 8 valence electrons. It needs 8 for its octet, but only
4 count towards its charge. WHY?
The 4 it ACTUALLY has balance out the 4 it SHOULD have, so it does not
have a formal charge. Its neutral.
1.4 Formal Charge
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-5
or
Covalent bonds are either polar or nonpolar:
Nonpolar covalent bonds: bonded atoms share electrons evenly
Polar covalent bonds: one of the atoms attracts electrons more than the
other
Electronegativity:
How strongly an atom attracts shared electrons
1.5 Polar Covalent Bonds
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-6
Electrons tend to shift away from lower
electronegativity atoms to higher electronegativity
atoms.
The greater the difference in electronegativity, the more
polar the bond.
1.5 Polar Covalent Bonds
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-7
Electrons behave as BOTH particles and waves. How can they be
both?
Maybe the theory is not yet complete.
The theory does match experimental data, and it has predictive
capability.
Like a wave on a lake, an electrons wavefunction can be (+), (-), or ZERO.
1.6 Atomic Orbitals
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-8
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Because they are generated
mathematically from wavefunctions,
orbital regions can also be (-), (+), or
ZERO.
The sign of the wave function has nothing
to do with electrical charge.
In this p-orbital, there is a nodal plane.
The sign of the wavefunction will be
important when we look at orbital
overlapping in bonds.
1.6 Atomic Orbitals
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-9
Once the 2s is full, electrons fill into the three
DEGENERATE 2p orbitals.
Where are the nodes in each of the 2p orbitals?
1.6 Atomic Orbitals
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-10
The bond for an H2 molecule results from constructive
interference.
Where do the bonded electrons spend most of their
time?
1.7 Valence Bond Theory
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-11
Atomic orbital wave functions
overlap to form molecular
orbitals (MOs) that extend
over the entire molecule.
MOs are a more complete
analysis of bonds because they
include both constructive and
destructive interference.
The number of MOs created
must be equal to the number
of atomic orbitals that were
used.
1.8 Molecular Orbital Theory
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-12
H2 MOs
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Imagine a He2 molecule.
How would its MOs
compare to those for H2?
How would the energy of
the He2 compare to 2 He?
Why does helium exist in its
atomic form rather than in
molecular form?
In general, if a molecule has
all of it MOs occupied, will it
be stable or unstable?
1.8 Molecular Orbital Theory
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-13
H2 MOs
To make CH4, the 1s atomic orbitals of the H atoms will
overlap with the four sp3 hybrid atomic orbitals of C.
1.9 Hybridized Atomic Orbitals
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-14
Consider ethene (ethylene).
Each carbon in ethene must bond to THREE other
atoms, so only THREE hybridized atomic orbitals are
needed.
1.9 Hybridized Atomic Orbitals
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-15
An sp2 hybridized carbon will have three equal-energy
sp2 orbitals and one unhybridized p orbital.
1.9 Hybridized Atomic Orbitals
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-16
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The sp2 atomic orbitals overlap to form sigma () bonds.
Sigma bonds provide maximum HEAD-ON
overlap.
1.9 Hybridized Atomic Orbitals
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-17
The unhybridized p orbitals in ethene form pi () bonds,
created by SIDE-BY-SIDE orbital overlap.
Practice with CONCEPTUAL CHECKPOINT 1.20.
1.9 Hybridized Atomic Orbitals
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-18
The sp atomic orbitals overlap HEAD-ON to form sigma
() bonds while the unhybridized p orbitals overlap
SIDE-BY-SIDE to form pi () bonds.
Practice with SKILLBUILDER 1.7.
1.9 Hybridized Atomic Orbitals
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-19
Explain the different strengths and lengths below.
Practice with CONCEPTUAL CHECKPOINT 1.24.
1.9 Hybridized Atomic Orbitals
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-20
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Valence shell electron pair repulsion (VSEPR theory):
Valence electrons (bonded and lone pairs) repel each other.
To determine molecular geometry:
1. Determine the steric number.
1.10 Molecular Geometry
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-21
To determine molecular geometry:
2. Predict the hybridization of the central atom:
If the steric number is 4, then it is sp3.
If the steric number is 3, then it is sp2.
If the steric number is 2, then it is sp.
1.10 Molecular Geometry
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-22
For any sp3 hybridized atom, the four valence electron pairs
will form a TETRAHEDRAL ELECTRON GROUP geometry:
Methane has four
equal bonds, so the
bond angles are
equal.
HOW does the
lone pair of ammonia
affect its geometry?
The bond angles in
oxygen are even
smaller. WHY?
1.10 Molecular Geometry
sp3 Geometry
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-23
The MOLECULAR geometry is different from the
ELECTRON GROUP geometry. HOW?
ExampleSteric
numberHybridization
Arrangement
of electron
pairs
Arrangement of
atoms
(geometry)
CH4 4 sp3 Tetrahedral Tetrahedral
NH3 4 sp3 Tetrahedral Trigonal pyramidal
H2O 4 sp3 Tetrahedral Bent
1.10 Molecular Geometry
sp3 Geometry
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-24
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7
Calculate the steric number for BF3.
Electron pairs that are located in sp2 hybridized orbitals
will form a trigonal planar ELECTRON GROUP geometry.
What will be the molecular geometry?
1.10 Molecular Geometry
sp2 Geometry
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-25 Practice with SKILLBUILDER 1.8.
1.10 Molecular Geometry
Summary
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-26
Electronegativity differences cause induction.
Induction (shifting of electrons WITHIN their orbitals)
results in a dipole moment.
Dipole moment = (the amount of partial charge) x (the
distance the + and - are separated)
Dipole moments are reported in units of debye (D)
1 debye = 10-18 esu cm
An electrostatic unit of charge (esu) is a unit of charge. One electron
has a charge of 4.80 x 10-10 esu.
Centimeters (cm) are included in the unit because the distance
between the centers of + and charges affects the dipole.
1.11 Molecular Polarity
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-27
Consider the dipole for CH3Cl
Dipole moment () = charge (e) x distance
(d)
Plug in the charge and distance
= (1.056 x 10-10 esu) x (1.772 x 10-8 cm)
Note that the amount of charge separation is
less than what it would be if it were a full
charge separation (4.80 x 10-10 esu).
= 1.87 x 10-18 esu cm
Convert to debye
= 1.87 D
1.11 Molecular Polarity
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-28
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Check out the polarity of come other common bonds:
1.11 Molecular Polarity
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-29
For molecules with multiple polar bonds, the dipole
moment is the vector sum of all of the individual bond
dipoles.
1.11 Molecular Polarity
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-30
Electrostatic potential maps are often used to give a
visual depiction of polarity.
1.11 Molecular Polarity
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-31
1.11 Molecular Polarity
Practice with SKILLBUILDER 1.9.
Klein, Organic Chemistry 1eCopyright 2012 John Wiley & Sons, Inc. 1-32
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Dipoledipole forces result when polar molecules line
up their OPPOSITE charges.
Note that acetones permanent dipole results from the
difference in electronegativity between C and O.
The dipoledipole attractions BETWEEN acetone
molecules affects acetones boiling point (BP) and
melting point (MP). HOW?
1.12 Intermolecular Forces
Dipoledipole
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-33
Why do isobutylene and acetone have such different
MPs and BPs?
1.12 Intermolecular Forces
Dipoledipole
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-34
Hydrogen bonds (H-bonds) are an especially strong type
of dipoledipole attraction.
Hydrogen bonds are strong because the partial + and
charges are relatively large.
Why are the partial charges in the H-bonding examples
below relatively large?
1.12 Intermolecular Forces
Hydrogen Bonding
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-35
Explain why the following isomers have different boiling
points.
1.12 Intermolecular Forces
Hydrogen Bonding
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-36
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The result is a fleeting attraction between the two
molecules.
Such fleeting attractions are generally weak.
However, like any weak attraction, if there are enough
of them, they can add up to a lot.
1.12 Intermolecular Forces
London Dispersion Forces
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-37
Explain why molecules with more mass generally have
higher boiling points.
Practice with SKILLBUILDER 1.10.
1.12 Intermolecular Forces
London Dispersion Forces
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-38
Explain why more highly branched molecules generally
have lower boiling points.
1.12 Intermolecular Forces
London Dispersion Forces
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-39
We use the principle that like dissolves like.
Polar compounds GENERALLY mix well with other polar
compounds:
If the compounds mixing are all capable of H-bonding and/or
strong dipoledipole interactions, then there is no reason why
they shouldnt mix.
Nonpolar compounds GENERALLY mix well with other
nonpolar compounds:
If none of the compounds are capable of forming strong
attractions, then no strong attractions would have to be
broken to allow them to mix.
1.13 Solubility
Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-40
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