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MIDTERM 2
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Atomic orbitals
CHEMISTRY
Most chemistry is
in the electrons
(the valence
electrons)
Atomic orbitals
Atomic orbitals
Representation of
the 2p orbitals.
Atomic orbitals
Wolfgang Pauli
Atomic orbitals
The orbitals filled for
elements in various parts
intra molecular bonding
CHEMISTRY
Full shells make
the most stable
atoms
intra molecular bonding
intra molecular bonding
The HCL molecule
has a dipole moment
intra molecular bonding
The Pauling electronegativity values
as updated by A.L. Allred in 1961. (cont’d)
Arbitrarily set F as 4
molecular structure
Skeletal Structure
• Hydrogen atoms are always terminal atoms.
• Central atoms are generally those with the
lowest electronegativity.
• Carbon atoms are always central atoms.
• Generally structures are compact and
symmetrical.
molecular structure
Exceptions to the Octet Rule
• Molecules with an odd number of electrons.
• Molecules in which an atom has less than an
octet of electrons.
• Molecules in which an atom
has more than an octet of
electrons.
molecular structure
Resonance Forms
• Lewis structures that differ only in the placement
of electrons are resonance forms. For O3:




O
O
O
=



 O





O=O

• Experimentally, it is found that both bonds are
0.128 nm long.
• The Lewis structure of O3 must show both
resonance forms.
molecular structure
Molecular Shapes
AB2
Linear
AB3
Trigonal planar
AB5
Trigonal bipyramidal
AB3E
Angular or Bent
AB4
Tetrahedral
AB4E
Irregular tetrahedral
(see saw)
AB6
Octahedral
AB3E
Trigonal
pyramidal
AB3E2
T-shaped
AB6E
Square pyramidal
AB3E2
Angular
or Bent
AB2E3
Linear
AB5E2
Square planar
Dipole Moment
Bond dipoles
C
O
In H2O the bond dipoles are also equal in
magnitude but do not exactly oppose each
other. The molecule has a nonzero overall
dipole moment.
O
Overall dipole moment = 0
O
Nonpolar
The overall dipole moment of a molecule
is the sum of its bond dipoles. In CO2 the
bond dipoles are equal in magnitude but
exactly opposite each other. The overall
dipole moment is zero.
F 
k  q1  q 2
d
m=Qr
2
Coulomb’s law
Dipole moment, m
Bond dipoles
H
H
Overall dipole moment
Polar
Bond Enthalpies and Bond Lengths
As bond order increases, the bond enthalpy
increases and the bond length decreases.
D(C-C) = 348 kJ
0.154 nm
D(C=C) = 614 kJ
0.134 nm
D(CC) = 839 kJ
0.120 nm
D(C-O) = 358 kJ
0.143 nm
D(C=O) = 799 kJ
0.123 nm
D(CO) = 1072 kJ
0.113 nm
Molecular orbitals
Hydrogen, H2
Hydrogen fluoride, HF
Fluorine, F2
Molecular orbitals
(a) Lewis structure of the methane molecule (b) the tetrahedral molecular
geometry
of the methane molecule.
Molecular orbitals
Hybrid Orbitals
Types of Hybrid Orbitals
sp
Shapes: linear
# orbitals: 2
sp2
sp3
sp3d
sp3d2
triangular tetrahedral trig. bipyram. Octahedral
3
4
5
6
Molecular orbitals
The relationship among the number
of effective pairs, their spatial arrangement,
and the hybrid orbital set required
Molecular orbitals
(a) Orbitals predicted by the LE model to describe (b)
The Lewis structure for carbon dioxide
Molecular orbitals
energies
The combination of hydrogen 1s atomic orbitals to form MOs
Molecular orbitals
energies
(a) The MO energy-level diagram for
the H2 molecule (b) The shapes of the Mos are obtained
by squaring the wave functions for MO1 and MO2.
Molecular orbitals
energies
The expected MO energy-level diagram for the combustion of
the 2P orbitals on two boron atoms.
The MO energy-level diagrams, bond
orders, bond energies, and bond lengths for the
diatomic molecules, B2 through F2.
Molecular orbitals
energies
Intermolecular Forces
Intermolecular forces
• The covalent bond holding a molecule together is an
intramolecular force.
• The attraction between molecules is an intermolecular
force.
• Intermolecular forces are much weaker than
intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for
HCl).
• When a substance melts or boils the intermolecular
forces are broken (not the covalent bonds).
• When a substance condenses intermolecular forces are
formed.
16a–26
Intermolecular forces
Larger INTERmolecular forces →
• Higher melting point
• Higher boiling point
• Larger enthalpy of fusion
16a–27
Intermolecular forces
Larger INTERmolecular forces →
• Higher melting point
• Higher boiling point
• Larger enthalpy of fusion
•Larger viscosity
•Higher surface tension
•Smaller vapor pressure
16a–28
Intermp;ecular forces
Intermolecular
Table of Force Energies
Type of Force
Energy (kJ/mol)
Ionic Bond
300-600
Covalent
200-400
Hydrogen Bonding
20-40
Ion-Dipole
10-20
Dipole-Dipole
1-5
Instantaneous Dipole/
Induced Dipole
0.05-2
Intermolecular Forces
Intermolecular forces
London Dispersion Forces
•
•
•
•
•
London dispersion forces increase as molecular weight increases.
London dispersion forces exist between all molecules.
London dispersion forces depend on the shape of the molecule.
The greater the surface area available for contact, the greater the dispersion forces.
London dispersion forces between spherical molecules are lower than between sausagelike molecules.
Intermolecular forces
H-Bonding
Occurs when Hydrogen is attached to a highly
electronegative atom (O, N, F).
d+
N-H… N-
O-H… N-
F-H… N-
N-H… O-
O-H… O-
F-H… O-
N-H… F-
O-H… F-
F-H… F-
dRequires Unshared Electron Pairs of Highly Electronegative Elements
Intermp;ecular forces
Intermolecular
Intermolecular Forces Summary
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16a–32
Intermolecular
Intramolecular
Intermolecular forces
Which forces?
London Dipole H-bond
Xe
CH4
CO2
CO
HBr
HF
CH3OH
NaCl
CaCl2
ionic
X
X
X
X
X
X
X
X
X
X
X
X
X
16a–33
Intermolecular forces
Relative forces
Larger
London
I2
>
<
H2O
H-bond
CH3OCH3
<
CH3CH2OH
H-bond
CsBr
>
Br2
CO2
<
CO
SF2
>
SF6
H2S
ionic
polar
Cl2
16a–34
polar
SOLIDS
Bonding in Solids
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16a–35
SOLIDS
Examples of Three Types of Crystalline Solids
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16a–36
crystals
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rights reserved.
16a–37
crystals
Figure 16.11: Reflection of X rays of
wavelength
n λ = 2 d sin θ
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16a–38
crystals
Atoms in unit cell
1
½
¼
1/8
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rights reserved.
16a–39
crystals
Cubic Unit Cells of Metals
Simple
cubic (SC)
1 atom/unit cell
Bodycentered
cubic (BCC)
2 atoms/unit cell
Facecentered
cubic (FCC)
4 atoms/unit cell
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16a–40
crystals
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rights reserved.
16a–41
crystals
Your eyes “see” 14 Cl- ions and 13 Na+
ions in the figure
Ion Count for the Unit Cell: 4 Na+ and 4 Cl- Na4Cl4 = NaCl
Can you see how this formula comes from the unit cell?
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rights reserved.
16a–42
crystals
fcc
V o lu m e :
D en sity :
R 8
4 x m ass
R 8
4 x
fra ctio n :
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16a–43
3
4
3
R
3
R 8
3
3
 0 .7 4
crystals
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16a–44
Concentration
Molarity =
Moles of solute/Liters of Solution (M)
Molality =
Moles of solute/Kg of Solvent (m)
Mole Fraction=
Moles solute/total number of moles
Mass %=
Mass solute/total mass x 100
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17a–45
solutions
solutions
phase cahnges
Figure 16.55: The phase diagram for water
phase cahnges
Qtotal = q1 + q2 + q3 + q4 + q5
Thermodynamics of Phase Changes
B
phasesolutions
cahnges
A
Why does a liquid at A form a solid when the temperature is lowered to B
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17a–49
phasesolutions
cahnges
Gases:
Large
Entropy
Liquid:
Smaller
Entropy
Solids:
Smallest
Entropy
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17a–50
Thermodynamics for Phase Change
∆G =
Negative for
spontaneous process
•
•
•
•
•
•
•
∆H
Negative for
liquid to solid
phasesolutions
cahnges
- T∆S
Positive for
liquid to solid
liquid→solid
∆H is negative (stronger intramolecular forces)
∆S is negative (more order)
-T∆S is positive
As T decreases, -T∆S becomes smaller
∆G goes to zero when ∆H = T∆S (at T = Tfusion)
For T less than Tfusion, ∆G is negative, solid is stable.
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17a–51
Factors Affecting Solubility
solutions
Gas – solvent: Pressure Effects
Henry’s Law:
C g  kP g
Cg is the solubility of gas, Pg the partial pressure, k
= Henry’s law constant.
Carbonated beverages are bottled under > 1 atm.
As the bottle is opened, Pg decreases and the
solubility of CO2 decreases. Therefore, bubbles of
CO2 escape from solution.
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17a–52
Raoult’s Law
• Raoult’s Law:
PA is the vapor pressure of A with solute
PA is the vapor pressure of A alone
A is the mole fraction of A
o
PA = XA PA
o
PTotal = XA PA
+
o
X B PB
solutions
solutions
Figure 16.44: Behavior of a liquid in a
closed container
solutions
Colligative properties
•
•
•
•
Vapor pressure –
Freezing point depression –
Boiling point elevation –
Osmosis -
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17a–55
Mole fraction
molality
molality
Molarity
semiconductors
Figure 16.24: A representation of the energy
levels (bands) in a magnesium crystal
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16a–56
semiconductors
Band structure of Semiconductors
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16a–57
semiconductors
Silicon Crystal Doped
with (a) Arsenic and (b)
Boron
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16a–58
semiconductors
Figure 16.34: The p-n junction involves the
contact of a p-type and an n-type semiconductor.
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16a–59
semiconductors
Semiconductors – key points to remember
• Band structure:
Valence band – gap – conduction band
•DOPING:
Group V  n type,
Group III  p type
•n-p junctions
•Devices:
(LED, laser, transistor, solar cell)
transition metal
complexes
What is a transition metal?
“an element with valance d- or f-electrons”
ie. a d-block or f-block metal
d-block: transition elements
3d
4d
5d
l=2
ml =
-2,-1,0,1,2
6d
l=3
ml =
-3,-2,-1,0, 4f
1,2,3
5f
f-block:
inner transition elements
transition metal
complexes
What is a coordination complex?
charge on complex
n+/-
ligands
X+/n
metal ion
counterion
•Central metal ion or atom surrounded by a set of ligands
•The ligand donates two electrons to the d-orbitals around the
metal forming a dative or coordinate bond
transition metal
complexes
-1
2+
transition metal
complexes
Common
Coordination
Numbers of
Transition
Metal
Complexes
Classes of isomers
transition metal
complexes
transition metal
complexes
Isomers I and
II
Energy of 3d orbitals
eg
t2g
transition metal
complexes
transition metal
complexes
Strong/weak fields, d6 Configuration
Paramagnetic – 4 Unpaired
Electron Spins
Diamagnetic – No Unpaired
Electron Spins
Correlation of High and Low Spin Complexes
With Spectrochemical Series
t2g4eg2
t2g3eg3
t2g6 t2g5eg1
transition metal
complexes

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