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Ch 7 Atomic Structure
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Rutherford Model
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Figure 7.2
Classification of Electromagnetic
Radiation
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Electromagnetic Radiation
Radiant energy that exhibits
wavelength-like behavior and
travels through space at the speed
of light in a vacuum.
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Light is a wave
Light: energy that travels like a wave through
space
Wave properties:
Wavelength 
Distance between 2
similar points (meters)
Frequency 
Number of waves in a
second (frequency (s1)
or hertz (Hz))
Speed
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v (m/s)
6
Wavelength and Frequency
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Light is a wave
All light travels the same speed:
high , has short 
low , has long 
high  = high energy
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Light is a wave
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Created by movement of
electric charge
An electric field and
magnetic field
perpendicular to each
Self-propagating
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Wavelength and frequency can be interconverted.
= c/
(C = 

 = frequency (s1) or hertz (Hz)
 = wavelength (m)
c = speed of light (m s1)
(2.9979 x 108 m/s)
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Figure 7.2
Classification of Electromagnetic
Radiation
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Matter is not what it appears to
be.
Before 1900:
Matter particle
Light a wave
Max Planck:
not all energies were
emitted from objects
heated incandescence
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Planck’s Constant
Energy gained or lost only in whole number
multiples. Transfer of energy is quantized:
occur in discrete units, called quanta.
E = nh  nhc

n = 1, 2,3,..
E = change in energy, in J
h = Planck’s constant, 6.626  1034 J s
 = frequency, in s1
 = wavelength, in m
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Light is a particle
Einstein: theorizes that
light made of photons.
Gets Nobel Prize
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Photoelectric Effect
Experiment: light of different frequency shone on
metal
Results: e- ejected only at minimum 
No e- ejected if  too low,
EVEN IF light intensity is increased.
WHY? Why some frequency and not others?
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Light made of photons
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- High  photon =
high E photon
- One photon hits one
electron
- If photon E not
= to e- E
nothing happens
(even if bright)
-Giant example
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Photoelectric Effect
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Light made of photons
Einstein: electromagnetic radiation is
quantized:
Ephoton = h = hc

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Energy and Mass
Einstein’s special theory of relativity:
(1905)
Energy has mass:
E = mc2
Or
m=E/c2
E = energy
m = mass
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Energy and Mass
Does a photon have mass?
for a photon with wavelength 
m = E = hc/  m h
c2
c2
c
Ephoton= hc/
(Hence the dual nature of light.)
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Figure 7.4
Dual Nature of light
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If light can be a particle can a
particle (e ) be a wave?
Louis de Broglie’s
Equation 1920
m = h / v
 h / m v
 = wavelength, in m
h = Planck’s constant, 6.626  1034 J s =
kg m2 s1
m = mass, in kg
v = velocity in m/s
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Wave interference
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Water wave interference
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Interference in water waves
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Water interference
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Interference patterns
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Light interference
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Interference Pattern
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Electron interference
Diffraction patterns caused by interference
X-rays passing through NaCl crystal are diffracted.
Electrons passing through NaCl crystal are diffracted
x-rays
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Diffraction using NaCl
crystal
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So, Debroglie was right: all
matter show both wave like and
particle like behavior
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How does all this stuff relate to
the e- and the atom?
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Bohr model: electrons
are at set distances from
the nucleus (energy
levels)
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How do we know this?:
we can use Einstein’s ideas to
explain bright-line spectra
Observe the light coming from the hydrogen emission tube.
“Excited” atoms only emit certain frequencies (colors) of light
Why not all frequencies of light?
Look at hydrogen emission. Each line is one frequency of light.
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Spectrum (a) and
A Hydrogen Line
Spectrum (b)
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Absorption, emission, and energy
Ground State: electrons in lowest
energy state.
photon
Absorption
Excited State: when one electron
absorbs one photon and jumps to
higher energy level.
When electron falls back to G.S. it
emits one photon.
photon
electron can only absorb photons or
emit photons of just the right energy
because levels are fixed.
Emission
The emitted photons are seen
as light of specific frequency
(i.e. colors).
What color emitted if electron could go
anyway?
A: the contiuous spectrum
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What is the energy difference
between levels?
Must be equal to the energy of the photon
emitted. Energy levels are quantized:
∆E = h = h c

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A Change between Two Discrete
Energy Levels
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Figure 7.8
Electronic Transitions in the
Bohr Model for the
Hydrogen Atom
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The Bohr Model
The electron in a hydrogen atom moves around the
nucleus only in certain allowed circular orbits. The
energy of each level is given by this equation:
E =  2.178  10
18
2
2
J (z / n )
E = energy of the levels in the H-atom
z = nuclear charge (for H, z = 1)
n = an integer
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The Bohr Model
Ground State: The lowest
possible energy state for an atom
(n = 1).
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Energy Changes in the Hydrogen
Atom
E = Efinal state  Einitial state
hc
 =
E
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Standing waves in spring:
- exist only at specific
 and  ( .5, 1,
1.5, 2.0, 2.5, ect.)
- are quantized.
Electron waves:
- exist only at certain 
and  (and energy)
- only form certain
distances from nucleus.
- are quantized
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Figure 7.10
The Hydrogen Electron
Visualized as a Standing Wave
Around the Nucleus
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Quantum Mechanics
Schrodinger’s equations: Based on the wave
properties of the atom
H  = E
 = wave function
H = mathematical operator
E = total energy of the atom
A specific wave function is often called an orbital.
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What is the orbital for H when n=1?
1s orbital
Orbitals are not the Bohr
orbits.
Where is the electron in
the orbital?
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Heisenberg
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There is a limit to how
precisely we can both the
position and momentum
of a an electron at a
given time
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Heisenberg Uncertainty
Principle
h
x   mv 
4
•
•
•
•
x = position
mv = momentum
h = Planck’s constant
The more accurately we know a
particle’s position, the less accurately we
can know its momentum.
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What is an orbital?
If we don’t know the motion
of an electron, what is an
orbital?
square
of the wave function
gives the probability of
finding an electron at a given
position.
--> (a) probability distrib.
for H 1s orbital
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Figure 7.12
Radial Probability Distribution
(the probability distribution in each spherical shell.)
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Bohr radius
Turns out that for H 1s orbital, the max radial
probability is 5.29x10-2 nm,
= Bohr’s innermost “orbit”.
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How big is the 1s?
Probab. Decreases with radius, but never goes
to zero.
Size definition: Size of the orbital is the radius
of the sphere that encloses 90% of the
electron’s probability.
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Quantum Numbers (QN)
Schrodinger equation has many wave function solutions.
Each are described by quantum numbers
1.
Principal QN (n = 1, 2, 3, . . .) - related to size and energy of the
(this gives the”rings” or “shells)
orbital.
2.
Angular Momentum QN (l = 0 to n  1) - relates to the shape of the orbital.
(ex. s p d f also called subshells)
3.
Magnetic QN (ml = l to  l ) - relates to orientation of the orbital in space
relative to other orbitals. Gives you the number of each type of orbital.
(ex.: px py pz)
4.
Electron Spin QN (ms = +1/2, 1/2) - relates to the spin states of the
electrons.
(see pg 310, table 7.2)
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Orbitals when n=1 (principal quantum #):
(l = 0 to n-1)
(ml = l to - l ):
l =0
ml=0
1s orbital only.
Only 1 or 2 electrons are
described by an orbital.
Total electrons at n=1? : .
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When n= 2,
(l = 0 to n-1)
ml = l to - l ):
l = 0 and l = 1
When l = 0, ml = 0
2s orbital
When l = 1 ml = -1, 0, 1
giving three 2p orbitals:
2px 2py 2pz
Total orbitals:
Total electrons:
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The P orbitals (energy level 2 and
up)
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When n=3 (l = 0 to n-1)
ml = l to - l ):
l=0,1,2
When l = 0, ml = 0
giving one 3s orbital
When l = 1 ml = -1, 0, 1
giving three 3p orbitals:
3px 3py 3pz
When l = 2 ml = -2,-1, 0, 1,2
giving 5 d orbitals
orbitals
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level 3:
Total electrons level 3:
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When n=4
l = 0 , 1 , 2, 3
when l = 0 , 1 , 2 : 4s 4px 4py 4pz 4 d’s (5 of them)
when l =3 4f orbitals (7 of them)
Total orbitals:
Total electrons:
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electron spin
quantum
number:
ms = +1/2 or
- 1/2
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Pauli Exclusion Principle
In a given atom, no two electrons can have
the same set of four quantum numbers (n, l,
ml, ms).
Therefore, an orbital can hold only two
electrons, and they must have opposite
spins.
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Quantum Model
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Hydrogen orbitals are degenerate
All H orbitals with the
same n have the same
energy.
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Polyelectronic Atoms
•
•
•
•
•
•
How does it work after Hydrogen?
Shielding (e- repel, feel less attraction to +)
Hydrogen orbitals: degenerate
Hydrogenlike orbitals: NOT degenerate!
Ens < Enp < End < Enf …etc.
Why?
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Figure 7.20
A Comparison of the Radial Probability
Distributions of the 2s and 2p Orbitals
2p appears closer to nucleus? Less energy?
No, look at small 2s hump.
“2s penetrates to the nucleus”
penetration effect
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Figure 7.21
The Radial
Probability
Distribution for the
3s, 3p, and 3d
Orbitals
so,
E3s<E3p<E3d
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History of the Periodic Table
7.10
Dobereiner: triads
Newlands: octaves
Meyer / Mendeleev: arrangements by atomic
masses.
Theory → prediction
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Figure 7.23
Mendeleev’s Early Periodic Table,
Published in 1872
Prediction: Ga, Ge (see table 7.3, pg318 to see how
cool Mendeleev was)
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Aufbau Principle
As protons are added one by one to the
nucleus to build up the elements,
electrons are similarly added to these
hydrogen-like orbitals.
“ an electron occupies the lowest
energy orbital that can receive it”
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Hund’s Rule
The lowest energy configuration for an
atom is the one having the maximum
number of unpaired electrons allowed
by the Pauli principle in a particular set
of degenerate orbitals.
“in the p, d, f, orbitals, spread out
before you pair up”
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The Electron Configurations in the Type of
Orbital Occupied Last for the First 18
Elements
Note: elements in same group have number valence electron
(valence: The electrons in the outermost principle quantum level of
an atom.
(Core electron: other than valence)
Ex:Cl
1s22s22p63s23p5 or [Ne] 3s23p5
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# valence = ?
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Figure 7.25
Electron Configurations for
Potassium Through Krypton
Why doe the 4s fill before the 3d? Penetration effect
Notice Cr, Cu columns.
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Figure 7.26
The Orbitals Being Filled for Elements
in Various Parts of the Periodic Table
After lathanum [Xe] 6s25d1 , go to lathanide series,
4fs (fig 7.27: note anomalies)
fill the
After Actinium [Rn]7s26d1, fill actinide series with 5fs.
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Figure 7.27 The Periodic Table With Atomic
Symbols, Atomic Numbers, and Partial Electron
Configurations
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Figure 7.36 Special Names for
Groups in the Periodic Table
Main-group elements or
representative elements:
1A 2A 3A 4A 5A 6A
7A 8A or 1,2, 13-18
M-g e: each group has
same #valence
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Broad Periodic Table
Classifications
Representative Elements (main group):
filling s and p orbitals (Na, Al, Ne, O)
Transition Elements: filling d orbitals (Fe,
Co, Ni)
Lanthanide and Actinide Series (inner
transition elements): filling 4f and 5f
orbitals (Eu, Am, Es)
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Figure 7.30
The Positions of the Elements
Considered in Sample Exercise 7.7
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Ionization Energy
The quantity of energy required
to remove one electron from the
gaseous atom or ion.
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Periodic Trends
First ionization energy:
increases from left to right across a
period (increase +, no shielding)
decreases going down a group.
( increase n, more shielding)
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Figure 7.31
The Values of First Ionization Energy for the
Elements in the First Six Periods
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Trends in Ionization Energies for the
Representative Elements
Who has the highest IE? Who has the lowest?
Do metals or nonmetals have higher IE?
What does IE tell us about metal reactivity
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Periodic Trends: IE
Al(g) --> Al+ + eAl+(g) --> Al2+ + eAl2+(g) --> Al3+ + eAl3+(g) --> Al4+ + e-
I1= 580 kJ/mol
I2= 1815 kJ/mol
I3= 2740 kJ/mol
I4= 11,600 kJ/mol
Why the differences? Indicates something
about the electron structure.
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Practice:
2
2
6
1
1s 2s 2p 3s
1s22s22p6
1s22s22p63s2
Which atom has the largest first I.E.?
Which one has the smallest second I.E.?
Explain.
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More Practice
IE increases across the period.
Check IE of P and S on pg 329. Explain
the anomaly
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Electron Affinity
The energy change associated with
the addition of an electron to a
gaseous atom or ion.
X(g) + e  X(g)
If change is exothermic, then E.A. is
negative.
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Figure 7.33
The Electronic Affinity Values for Atoms
Among the First 20 Elements that Form
Stable, Isolated X- Ions
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Electron Affinity
Left to right mostly more neg kJ, more energy released. (note
the missing elements: why C, but not N? write o.d. for both
and t.t.y.n)
Down a group, usually less neg kJ, less energy released.
(see table 7.7. Notice anomaly. T.t.y.n.)
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E.A.
Who has more neg EA, metals or nonmetals?
What does this say about the reactivity of
nonmetals?
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Radius trend
Data usually from distance between nuclei in
a compound.
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Radius trend
Data usually from distance between nuclei in
a compound.
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Figure 7.35
Atomic Radii
for Selected
Atoms
Period trend?
Group trend?
Why?
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Periodic Trends
Atomic Radii:
decrease going from left to right across
a period: increase + draws in valence.
increase going down a group: increase
in orbital sizes with n.
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Alkali Metals Trends
look at pg 335.
Note trends down the group in:
1. IE, radius
2. Density. Why?
3. mp/bp
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Alkali metals
Reaction with water. Write the balanced equation for
Na(s) with water.
List alkali from most to least reactive (think I.E.)
Cs > Rb > K > Na > Li
But in water,
Li > K > Na , even though K loses electrons the
easiest. why? Hydration energy (see table 7.9)
Li is small, higher charge density, better at
attracting water.
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Alkali metals
What we observe when they react with water:
K > Na > Li
Look at mp. K and Na melt, increasing reaction
rate.
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Waves
Waves have 3 primary characteristics:
1.
Wavelength: distance between two
peaks in a wave.
2.
Frequency: number of waves per
second that pass a given point in space.
3.
Speed: speed of light is 3.00  108
m/s.
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Figure 7.1 The Nature of Waves
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