07/12/2010 1 Visible light is a form of electromagnetic radiation Electromagnetic radiation is characterized by its wave nature 6.1 The wave nature of light The electromagnetic spectrum or “radiant energy” Chapter 6 Electronic structure of atoms c = nl 6.2 Quantized Energy and Photons 1. Blackbody radiation 2. The photoelectric effect 3. Emission spectra emission of light from hot objects emission of electrons from metal surfaces on which light shines emission of light from electronically excited gas atoms Some phenomena cannot be explained using a wave model of light: A quantum is the smallest amount of energy that can be emitted or absorbed as electromagnetic radiation The relationship between energy, E, and frequency is: E = hn where h is Planck’s constant = 6.626 × 10 -34 joule-seconds (J.s) Energy of one photon = E = hn = hc/l
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07/12/2010
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Visible light is a form of electromagnetic radiation
Electromagnetic radiation is characterized by its wave nature
6.1 The wave nature of light
The electromagnetic
spectrum
or “radiant energy”
Chapter 6 Electronic structure of atoms
c = nl
6.2 Quantized Energy and Photons
1. Blackbody radiation
2. The photoelectric effect
3. Emission spectra
emission of light from hot objects
emission of electrons from metal surfaceson which light shines
emission of light from electronically excited gas atoms
Some phenomena cannot be explained using a wave model of light:
A quantum is the smallest amount of energy that can be emitted or absorbed as electromagnetic radiation
The relationship between energy, E, and frequency is:
E = hnwhere h is Planck’s constant = 6.626 × 10-34 joule-seconds (J.s)
Energy of one photon = E = hn = hc/l
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Photoelectric effect
Chem 101 3
Chem 101 4
Bohr’s model predicts:
= -2.18 x 10-18 J {(1/nf2)-(1/ni2)} = E photon = h
Allowed energy states
Ground state
Electron escapes the nucleus -ionization
E = -hcRH (1/n2)= -2.18 x 10-18 J (1/n2), where n is an integer between 1 and ∞
Electron microscope imagehttp://intranet.dalton.org/departments/science/Science5/microscopy.html
Chem 1016
www.grayfieldoptical.com/
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A. No matter waves are produced.B. No, because the mass of the baseball is too large.C. Yes; but too small to allow any way of observing them.D. Yes; and they can be observed.E. Let me ask YOU; what is the sound of one hand clapping?
A baseball pitcher throws a fastball at 150 km/h. Does that moving baseball generate matter waves? If so, can we observe them?
C. Yes; but too small to allow any way of observing them.
A baseball pitcher throws a fastball at 150 km/h. Does that moving baseball generate matter waves? If so, can we observe them?
λ = h/mv= (6.63 x 10-34 J s)/{(150 g)(150 km/hr)}
λ = {6.63 x 10-34 (kg m2 /s2) s}/{(0.150 kg)(150x103m/3600s)}
λ = 1.08x10-34m
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Chem 101 9
Δx • Δ(mv) ≥ h / 4π
Δx = uncertainty in position (m)Δ(mv) = uncertainty in momentum (kgms-1)h =Plank’s constant (s-1)
4π = 4π
Probability function (Ψ2)
analogy: compare probability of dart landing herevs. there
Chem 101 10
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11
For interest only: do not need to know
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Orbital shapes (www.quimica3d.com)
Chem 101 13
To memorise
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To be aware of (ie: draw a d orbital)
Orbital shapes
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Chem 101 17
s orbitals (ℓ = 0)
n = 1, ℓ = 0
1s orbital
n = 2, ℓ = 0
2s orbital
node
[4πr2Ψ(r)2]
pg 230-231 (a closer look)
Chem 101 18
describes the main energy level; specifies electron
shell
describes the shape; specifies subshell
designates specific orbital; specifies orientation
mℓ = (-ℓ),…,0,…,(+ ℓ)
must be a positive integer n = 1,2,3,4,…
maximum value is (n-1), i.e. ℓ = 0,1,2,3…(n-1)use letters for ℓ (s, p, d and f for ℓ = 0, 1, 2, and 3).
maximum value depends on ℓ, can take integral values from – ℓ to + ℓ
ΔH is negative (rxn. is exothermic) whenweak bonds are broken and strong bonds are formed
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Bond lengths
Chem 101 61
also depend on nature of atom and type of bond
trends : shorter bonds are stronger
1.47 163 kJ/mol
1.24 418 kJ/mol
N N
N N
N N 1.10 941kJ/mol
also calculated as averages
9.1 Molecular Shapes
AB2AB3
linear bent trigonal planar
trigonal pyramidal
T-shaped
Chapter 9 Molecular Geometry and Bonding Theories
For molecules of the general form ABn there are 5 fundamental shapes:
180° 109.5°
linear trigonal planar
120°
tetrahedral
90°
120°
90°
90°
trigonal bipyramidal
octahedral
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1. Draw Lewis structure, count electron domains
We use the electron-domain geometry to help us predict the molecular geometry.
2. Arrange electron domains to minimize repulsion
3. Inspect arrangement of atoms to determine molecular geometry
3
tetrahedral
trigonal pyramidal
2
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2 0
3 0
2 1
2
3
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4 0
3 1
4
5 05
4 1
PCl5
SF4lone pair always in least crowded position
Molecules with Expanded Valence Shells
Atoms that have expanded octets have five electron domains (trigonal bipyramidal) or six electron domains (octahedral) electron-domain geometries.
2 2
Effect of nonbonding electrons and multiple bonds on bond angles
bonding pair
non-bonding pair
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3 25
2 3
ClF3
XeF2
6 06
5 1
SF6
BrF5
4 2 XeF4
note position of lone pairs
lone pairs as farapart as possible
9.3 Molecular Shape and Molecular Polarity
Dipoles are a vector quantity
overall dipole moment
= 0
electron density models
(red = high, blue = low)
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9.4 Covalent Bonding and Orbital Overlap
HCl2
HHCl
overlap regions
The change in potential energy as two hydrogen atoms combine to form the H2 molecule:
H2
9.5 Hybrid Orbitals
s p2 × sp
large lobes of sp hybrid orbitals
overlap regions
F 2porbital
F 2porbital
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D. The unhybridized p-orbital is perpendicular to the plane of the sp2 orbitals.
In an sp2 hybridized atom, what is the orientation of the unhybridized p orbital relative to the three sp2 hybrid orbitals?
one sorbital
two porbitals
hybridize
three sp2
hybrid orbitals
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9.5 Hybrid Orbitals
s p2 × sp
large lobes of sp hybrid orbitals
overlap regions
F 2porbital
F 2porbital
one sorbital
two porbitals
hybridize
three sp2
hybrid orbitals
sp2 hybrid orbitals shown together
(large lobes only)
one sorbital three p orbitals
four sp3 hybrid orbitals
sp3 hybrid orbitals shown together
(large lobes only)
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For geometries involving expanded octets on the central atom, we use d orbitals in our hybrids:
octahedrals, p, p, p, d
five sp 3d
s, p, p, p, d, d
six sp 3d 2
trigonal bipyramidal
9.5 Hybrid Orbitals
one π bond
9.6 Multiple Bonds
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Delocalized π Bonding
When writing Lewis structures for species like the nitrate ion, we draw resonance structures to more accurately reflect the structure of the molecule or ion
In reality, each of the four atoms in the nitrate ion has a p orbital
the π electrons are delocalized throughout the ion
The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen
The organic molecule benzene has six σ bonds and a p orbital on each carbon atom
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Molecular orbital (MO) theory (only section 9.7)
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MO diagram (energy level diagram)
Chem 101 83
bonding MO lowers energy
antibonding MO raises energy
bonding electrons
Bond order
Bond order = ½ {no. bonding electrons – no. antibonding electrons}
antibonding electrons
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Metal bonding (Sections 23.5 and 12.2)
MO model
an infinite chain of atoms
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Metals, insulators and semiconductors
11.1 A molecular comparison of gases, liquids and solidsThe fundamental difference between states of matter is the distance between particles.
gasliquid
solid
incompressible
Chapter 11 Intermolecular forces, liquids, and solids
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Intermolecular forces: a summary
interacting molecules or ions
are ions involved?
are polar molecules and ions
both present?
are polar molecules involved?
are H atoms bonded to N, O or
F atoms?A) ionic bondinge.g. NH4NO3
YES
NO
B) ion-dipole forcese.g. NaCl in H2O
YES NO
E) dispersion forcesonly (induced dipoles)
e.g. Ar(l), I2(s)
D) dipole-dipole forces
e.g. H2S, CH2Cl2
C) hydrogen bondinge.g. H2O, NH3, HF
YES
NO
NO
YES
van der Waals forcesA > B > C > D > E
11.2 Intermolecular forces
covalent bond (strong)
intermolecular attraction (weak)
Dipole-Dipole Interactions
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Hydrogen bonding
H2O
CH4
H2SeH2Te
SiH4GeH4
SnH4H2S
London Dispersion Forces
e ―
e ―
He atom
momentarily polar
δ– δ+
e ―
e ―e ―
e ―
He atom 1 δ– δ+He atom 2 δ– δ+
n-pentanebp 309 K
neopentanebp 283 K
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Ion-dipole forces
Ion-dipole interactions, are an important force in solutions of ions.
cation-dipole
anion-dipole
11.8 Bonding in solidsMolecular Solids
Consist of atoms or molecules held together by intermolecular forces.
London dispersiondipole-dipole
hydrogen bonding
benzene toluene phenol 5 –95 43 80 111 182
high mp due to efficient packing
high bp due to larger intermolecular forces
high mp & bp due to hydrogen
bonding
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Covalent-Network Solids
Consist of atoms held together, in large networks or chains, with covalent bonds.
1.42 Å
3.41 Å
Ionic Solids
Consist of ions held together by ionic bonds (i.e. by electrostatic forces of attraction).
CsCl ZnS“zinc blende”
CaF2“fluorite”
Cs
Cl
Zn
SCa
F
Metallic Solids
Consist entirely of metal atoms;
not covalently bonded.
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Hydrocarbons
Chem 101 97
contain only C and H
four different types: defined by kind of carbon to carbon bonds
largest possible no. of H atoms saturated hydrocarbons