Download Midterm 2 Review slides from November 15

17/11/2010
Reading assignment: 8.5-8.8
As you read ask yourself:
How can I use Lewis structures to account for bonding in covalent
m
molecules?
What are the differences between single, double and triple bonds in
terms of bond distance and strength?
When is it useful to assess the formal charges on atoms in a Lewis
structure?
How does a resonance hybrid structure differ from a regular Lewis
structure? How do resonance structures affect the predicted bond
distances?
What are the situations when the octet rule is disobeyed? How will
I recognize molecules that disobey the octet rule?
How can I make use of average bond enthalpies?
Chem 101
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Lewis structures
1. Find the total number of valence
electrons (account for any charges) =
TOTAL
1 + 4 + 5 = 10
2. decide connection between atoms, draw
a line to represent 1 electron pair for
each connection, count the electrons
SHARED
2+2=4
3. calculate the remaining electrons =
TOTAL – SHARED, assign these to the
terminal atoms to make octet (or 2 for H
atom)
10 - 4 = 6
Choose central atom
correctly
least electronegative atom
(not H)
oxygen rarely bonds to
itself
more than 1 central
atom (e.g. N2O4)? 
make a symmetric
arrangement
4. any electrons left? – put them on the
central atom
5. if central atom doesn’t have an octet,
make multiple bonds from nonbonded
electron pairs on terminal atoms
Chem 101
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Formal charges
often can make more than one Lewis structure
which one is
correct?
bookkeeping of electrons
calculate the charge on atom IF all bonding electrons shared equally
assign to the atom
all unshared (nonbonding) electrons
+
½ of all bonding (shared) electrons
formal charge = number of valence electrons – total assigned electrons
Evaluate Lewis structures
more stable if there are small (or no) formal charges
the most electronegative atom has the most negative
formal charge
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Resonance
actual structure is
neither A or B but a resonance hybrid structure
hybrid is intermediate between the two “parent”
parent structures
Resonance has impact on bond lengths and strengths
Chem 101
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Exceptions to the octet rule
1. less than an octet
small atoms that are too crowded with an octet
2. odd number of electrons
3. more than an octet
octet impossible, small number of stable
molecules
elements in third period (and beyond) can
expand valence shell
can expand valence shell to make a Lewis structure with lower formal charge
Chem 101
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Covalent bond strength
stability of molecule is related to strength of covalent bonds
energy change when a particular bond is broken in one mole of
gaseous substance is bond enthalpy, ΔH
HCl(g)
H(g) + Cl(g) H = 431 kJ
bond enthalpies are always positive
the greater the ΔH, the stronger the bond
depend on
atoms in the bond and type of bond
(single double or triple)
(single,
For polyatomic molecules the bond enthalpies are average values
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Estimate enthalpy change of a reaction
overall change is the difference
between bonds broken (it takes energy to break bonds)
and
bonds formed (forming bonds releases energy)
Chem 101
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ΔH is negative (rxn. is exothermic) when
Chem 101
weak bonds are
broken and strong
bonds are formed
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Bond lengths
also depend on nature of atom and type of bond
also calculated as averages
trends : shorter bonds are stronger
N
N
1.47
163 kJ/mol
N
N
1.24
418 kJ/mol
N
N
1.10
941kJ/mol
Chem 101
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Reading: Chapter 9, sections 9.1-9.2
As you read these sections ask yourself:
How does the number
mbe of bonds
bo ds and
a d nonbonded
o bo ded pairs
pai s of
electrons affect the shape of a molecule?
How is the geometry of a molecule defined?
Why is the repulsion between two domains of
nonbonded pairs of electrons greater than between two
domains of bonded pairs of electrons?
Chem 101
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Chapter 9
Molecular Shapes
Lewis structures provide
info on number and types of
bonds
no info on shape of
molecule in 3D
shape determined by bond angle:
angle between bonds joining
nuclei
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The Valence Shell Electron Pair Repulsion
Model
VSEPR
Assumptions:
 Shape determined by numbers of
valence electrons around the atoms
bonded and nonbonded pairs
 Electron domains arranged to
minimize
repulsion between
domains of electrons
VSEPR predicts
shape of electron
arrangement
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VSEPR
model
leads to approximate molecular shape, based on
electron locations
consider electron domains around central atom
domains may be
single bond
double bond
triple bond
nonbonded pair
each
h off these
th
counts
t
as 1 domain
number of domains around central atom determines the shape of
electron domains
arranged to minimize the
repulsion
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Optimal shape determined by number of
domains
Two domains: linear, bond angle = 180°
Three electron domains:
Four electron domains:
Five electron domains:
Six electron domains
Chem 101
trigonal planar
tetrahedron
bond angles = 120°
bond angles = 109.5°
trigonal bipyramid
octahedral
two bond angles 90° or 120°
all bond angles = 90º
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VSEPR steps
Draw Lewis structure
count electron domains: single, multiple, nonbonded
Determine electron domain g
geometry
y
NOT molecular g
geometry
y
Molecular shape is determined by position of bonded atoms
example:
NCl3
electrons = (7 × 3) + 5 = 26
4 domains
electron domain
shape is tetrahedral
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Nonbonding electrons, multiple bonds affect bond
angles
nonbonded electrons occupy more space than bonded pairs
in selected VSEPR geometries: there will be favoured positions for nonbonded
electron pairs
multiple bonds also have
larger electron domains
nonbonded electrons and multiple bonds compress bond angles due to
greater repulsion
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sp hybrid
sp2 hybrid
sp3 hybrid
h b id
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sp3d hybrid
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sp3d2 hybrid
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Reading: Chapter 9, sections 9.4-9.7
As you read these sections ask yourself:
How can a molecule with polar bonds be nonpolar?
Why do we need theories of bonding that differ from
VSEPR?
How does Valence Bond theory differ from the Lewis
concept of chemical bonding?
How does molecular orbital theory differ from valence
bond theory?
How d
H
does a h
hybrid
b id orbital
bit l diff
differ ffrom a pure atomic
t i
orbital?
How are hybrid orbitals related to the VSEPR shapes you
learned earlier?
How do sigma and pi bonds differ from each other?
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Molecular shape and polarity
Can now take into consideration the effect of bond polarity
on the overall molecule
Bond polarity is a measure of unequal sharing of electrons in a
bond
depends on differences in
electronegativity
dipole moment defined as a measure of the separation and
magnitude of charge
of a p
polar molecule
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The overall molecular dipole moment depends on
 polarity of the individual bonds
(electronegativity difference and direction)
 overall geometry of the molecule
bond dipoles are vector quantities
molecular dipole moment
Chem 101
is the vector sum of the bond dipoles
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To determine if a molecule is polar…
draw Lewis structure and determine geometry
determine if the bonds are polar
p
determine if the polar bonds add together (based on geometry)
to form a net dipole moment
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Bonding theories
VSEPR models and Lewis structures
good for predicting molecular shapes
do not explain why bonds exist
do not explain how the electron’s atomic orbitals are involved in
bonding
Need a theory that combines the idea of two electron bonds
with the theory of atomic orbitals
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Valence Bond Theory
 valence electrons are in the localized atomic orbitals of isolated
atoms
these are the s,p,d,f
,p, , orbitals
 bond is formed from overlap of half-filled valence orbitals, spin-pairing of
valence electrons
if the interactions
lower energy,
a bond is formed
 shape of molecule determined
by geometry of overlapping
orbitals
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Make new orbitals
orbitals in a molecule don’t have to be the same as in an atom
orbitals are (wave) functions
can make math combinations to form
new orbitals
has the effect of mixing the orbitals
Hybrid orbitals
- have shapes that match actual electron distribution in bonded atoms
- number of hybrid orbitals = number of atomic orbitals mixed
- central or interior atoms have the greatest tendency to hybridize
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Bonding
schemes:
1. Draw the Lewis structure
2. Determine the electron domain geometry using VSEPR
3 Choose hybrid orbitals for central/interior atoms based on VSEPR shape
3.
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Multiple Bonds
The sp and sp2 hybrid orbitals have unused p orbitals on the central atom
These p orbitals are perpendicular to the hybrid orbitals
The p orbitals can overlap sideways
Two kinds of
overlap
p
σ bond –
overlap along
line between
nuclei
Chem 101
π bond –
overlap above
and below the
line between
nuclei
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single bonds are
double bonds are
triple bonds are
always sigma (σ) bonds
one sigma bond and one π bond
one σ bond and two π bonds
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Resonance in valence bond
theory
two or more resonance structures with pi bonds can not
be described with localized bonding
th pii b
the
bonding
di iin resonance structures
t t
iis
d l
delocalized
li d
electrons are delocalized (smeared out) over more than 2 atoms
all atoms with delocalized π bonding must be in the same plane
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Bonding
schemes:
1. Draw the Lewis structure
2. Determine the electron domain geometry using VSEPR
3 Choose hybrid orbitals for central/interior atoms based on VSEPR shape
3.
4. sketch molecule starting with central atom and its orbitals, show overlap
5. label all bonds using the σ and π notation
Chem 101
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Molecular orbital (MO) theory (only section 9.7)
Need better theory to understand excited states and the
properties of some molecules
MO theory describes the electrons in molecules
with wave functions called molecular orbitals
wave function over entire molecule is constructed from the
atomic orbitals of all the atoms in the molecule
MO can hold a maximum of 2 electrons
MO has a definite energy
MO has an electron-density distribution
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MO diagram (energy level diagram)
antibonding MO
raises energy
bonding
g MO
lowers energy
bonding electrons
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He
He
He2
predict that He2 is unstable
and does not form
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Bond order
stability of the bond depends on the relative number of bonding
and antibonding electrons
Bond order = ½ {no. bonding electrons – no. antibonding electrons}
antibonding electrons
Chem 101
Bd.ord. = ½
Bd.ord. = 0
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Metal bonding (Sections 23.5 and 12.2)
Electron sea model
array of metal cations in
a ‘sea’ of electrons
electrons are mobile
and uniformly distributed
model explains
p
conductivity,
y,
malleability and ductility
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MO model
combines atomic orbitals to make MO over entire molecule
each MO can hold 2 electrons
number of MO = number of atomic orbitals combined
in general:
lowest energy MO are the most bonding
hi h t energy MO are mostt antibonding
highest
tib di
as no. of atoms increases energy separation between
MOs decreases  bands
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bands are not independent and can
be represented as one set of energy
levels
Bonding MOs are called the valence band
In the band structure roughly half of the
MOs (or energy levels) are BONDING
Antibonding MOs are called the
conduction band
the upper half (high energy half) are ANTIBONDING
Conduction arises when electrons are
promoted into unoccupied MOs or
energy levels
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Bonding strengthens as electrons are added to bonding orbitals
strength of bonding in the transition
metals increases until the band
structure is roughly half-full
roughly 6-7 electrons
strength decreases with more than 67 valence electrons because some
electrons are in antibonding orbitals
valence electron configurations
3B
ns2 (n-1)d
( 1)d1
6B
ns1 (n-1)d5
1B
ns1 (n-1)d10
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Metals, insulators and semiconductors
differ in the size of the gap between the valence and conduction bands
metals
t l
insulators
h
have
partially
ti ll fill
filled
d band
b d with
ith no gap
have a large gap, filled valence band and empty conduction band
semiconductors
Chem 101
have a small gap, that electrons can be induced across
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