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Chapter 7
Atomic Structure
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7.1 Types of EM waves
They have different  and 
 Radio waves, microwaves, infra red,
ultraviolet, x-rays and gamma rays are
all examples.
 Light is only the part our eyes can
detect.
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Wavelength increases
Gamma
Rays
Frequency decreases
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Radio
waves
Parts of a wave
Wavelength

Frequency = number of cycles in one second
Measured in hertz 1 hertz = 1 cycle/second
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Frequency = 
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The speed of light
c = in a vacuum is 2.998 x 108 m/s c = 
 What is the wavelength of light with a
frequency 5.89 x 105 Hz?
 What is the frequency of blue light with a
wavelength of 484 nm?
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(YDVD) EMwave
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7.2 Nature of Matter
Matter and energy were seen as different
from each other in fundamental ways.
 Matter was particles.
 Energy could come in waves, with any
frequency.
 Max Planck found that as the cooling of
hot objects couldn’t be explained by
viewing energy as a wave.
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Energy is Quantized
Planck found E came in chunks with
size h
 E = nh
 where n is an integer.
 and h is Planck’s constant
 h = 6.626 x 10-34 J•s
 These packets of h are called quantum.
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Einstein is next
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Said electromagnetic radiation is quantized in
particles called photons.
Each photon has energy = h = hc/
Blue color in fireworks has a wavelength of
450 nm. How much energy does a photon of
blue light have? A mole of these photons?
Combine this with E = mc2
You get the apparent mass of a photon.
m = h/(c)
Which is it?
Is energy a wave like light, or a
particle?
 Yes
 Concept is called the Wave - Particle
duality. (Flatman)
 What about the other way, is matter a
wave? Yes
 Photoelectric Effect (YDVD)
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(BDVD) Blue & Yellow
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Matter as a wave
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All matter exhibits both particle and wave
properties.
De Broglie’s equation  = h/mv
We can calculate the wavelength of an object.
The laser light of a CD is 7.80 x 102 m. What is
the frequency of this light?
What is the energy of a photon of this light?
What is the apparent mass of a photon of this
light?
What is the energy of a mole of these
photons?
What is the wavelength?
m = h/(c)
 Of an electron with a mass of
9.11 x 10-31 kg traveling at 1.0 x 107 m/s?
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Of a softball with a mass of 0.10 kg moving
at 35 mi/hr?
7.3 Hydrogen spectrum
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Emission spectrum because these are the
colors it gives off or emits.
Called a line spectrum.
There are just a few discrete lines showing
only certain energies are possible.
656 nm
434 nm
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410 nm
486 nm
What this means
Only certain energies are allowed for
the hydrogen atom.
 Can only give off certain energies.
 Use E = h = hc/
 Energy in the in the atom is quantized.
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(YDVD) Refraction
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7.4 Niels Bohr
Developed the quantum model of the
hydrogen atom.
 He said the atom was like a solar
system.
 The electrons were attracted to the
nucleus because of opposite charges.
 Didn’t fall in to the nucleus because it
was moving around.
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The Bohr Ring Atom
He didn’t know why but only certain
energies were allowed.
 He called these allowed energies energy
levels.
 Putting Energy into the atom moved
the electron away from the nucleus.
 From ground state to excited state.
 When it returns to ground state it gives
off light of a certain energy.
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The Bohr Ring Atom
n=4
n=3
n=2
n=1
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The Bohr Model
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n is the energy level (an integer).
Z is the nuclear charge, which is +1 for
hydrogen.
E = -2.178 x 10-18 J (Z2/n2 )
n = 1 is called the ground state.
When the electron moves from one energy
level to another.
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E = Efinal - Einitial
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E = -2.178 x 10-18 J Z2 (1/ nf2 - 1/ ni2)
Examples
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E = -2.178 x 10-18J Z2 (1/ nf2 - 1/ ni2)
Calculate the energy need to move an electron
from level n=1 to level n=2.
What is the wavelength of light that must be
absorbed by a hydrogen atom in its ground
state to reach this excited state.
Calculate the energy released when an
electron moves from n= 5 to n=3 in a He+1 ion
The Bohr Model
Doesn’t work.
 Only works for hydrogen atoms.
 Electrons don’t move in circles.
 The quantization of energy is right, but
not because they are circling like
planets.
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7.5 The Quantum Mechanical Model
A totally new approach.
 De Broglie said matter could be like a
wave.
 De Broglie said they were like standing
waves.
 The vibrations of a stringed instrument.
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What’s possible?
You can only have a standing wave if
you have complete waves.
 There are only certain allowed waves.
 In the atom there are certain allowed
waves called electrons.
 1925 Erwin Schroedinger described the
wave function of the electron.
 A lot of math, but what is important are
the solutions.
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Schroedinger’s Equation
The wave function ( is a (x, y, z)
function.
 Solutions to the equation are called
orbitals.
 These are not Bohr orbits.
 Each solution is tied to a certain energy.
 These are the energy levels.
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The Heisenberg Uncertainty
Principle.
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There is a fundamental limitation on how
precisely we can know both the position and
momentum of a particle at a given time.
MATHEMATICALLY:
x · (mv) > h/4
x is the uncertainty in the position.
(mv) is the uncertainty in the momentum.
the minimum uncertainty is h/4
What does the wave Function
mean?
Nothing. It is not possible to visually
map it.
 The square of the function is the
probability of finding an electron near a
particular spot.
 Best way to visualize it is by mapping
the places where the electron is likely to
be found.
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Defining the size
The nodal surface.
 The size that encloses 90% to the total
electron probability.
 NOT at a certain distance, but a most
likely distance.
 For the first solution it is a a sphere.
This is of course and s orbital.
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7.6 Quantum Numbers
There are many solutions to
Schroedinger’s equation
 Each solution can be described with
quantum numbers that describe some
aspect of the solution.
 Principal quantum number (n) - size
and energy of of an orbital.
 Has integer values > 0
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Quantum numbers
Angular momentum quantum number l .
 shape of the orbital.
 integer values from 0 to n - 1
 l = 0 is called s
 l = 1 is called p
 l =2 is called d
 l =3 is called f
 l =4 is called g
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Quantum numbers
Magnetic quantum number (ml)
 integer values between - l and + l
 tells direction in each shape.
 Electron spin quantum number (ms)
 Can have 2 values.
 either +1/2 or -1/2
 Q: What are the possible quantum
numbers for the last electron of a sulfur
atom? Iron?
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Example:
Describe the electrons defined by the following
quantum numbers:
nl m
30 0
3s electron
21 1
2p electron
4 2 -1
4d electron
33 2
not allowed (l must be < n)
31 2
not allowed (ml must be between
-l and l)
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1. What are the shapes of s, p, and d orbitals respectively?
2. How many 1s orbitals are there in an atom? 4p orbitals?
4d orbitals?
3. What is the maximum number of orbitals with:
n=4l=1
n=2l=2
n=3l=2
n = 5 l = 1 ml = -1
4. Which orbitals cannot exist?
2p 3p 4d 3f 6s 2d
5. Write a set of quantum numbers for a 4f orbital.
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1. What are the shapes of s, p, and d orbitals respectively?
s= spherical p = dumbbell d = cloverleaf
2. How many 1s orbitals are there in an atom? 4p orbitals? 4d orbitals?
1s: 1 4p: 3 4d: 5
3. What is the maximum number of orbitals with:
n = 4 l = 1 3 (the 4p orbitals)
n = 2 l = 2 none (l must be < n)
n = 3 l = 2 5 (the 3d orbitals)
n = 5 l = 1 ml = -1 1 (3 q.n. define a unique orbital)
4. Which orbitals cannot exist?
2p 3p 4d 3f 6s 2d
3f and 2d
5. Write a set of quantum numbers for a 4f orbital.
n = 4 l = 3 ml = 3, 2, 1, 0, -1, -2, -3
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S orbitals
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(BDVD) orbitals
P orbitals
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D orbitals
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F orbitals
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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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7.9 Polyelectronic Atoms
More than one electron.
 Three energy contributions.
 The kinetic energy of moving electrons.
 The potential energy of the attraction
between the nucleus and the electrons.
 The potential energy from repulsion of
electrons.
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Polyelectronic atoms
Electron shielding occurs. This affects the
properties of these atoms.
 Electrons are attracted to the nucleus.
 Electrons are repulsed by other electrons.
 Electrons would be bound more tightly if
other electrons were not present.
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7.10 The Periodic Table
Developed independently by German
Julius Lothar Meyer and Russian Dmitri
Mendeleev (1870”s).
 Didn’t know much about atom.
 Put in columns by similar properties.
 Predicted properties of missing
elements.
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7.11 Aufbau Principle
Aufbau is German for building up.
 As the protons are added one by one,
the electrons fill up hydrogen-like
orbitals.
 Fill up in order of energy levels.
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Increasing energy
7s
6s
5s
5p
4p
4s
6d
5d
4d
3d
3p
3s
2p
2s
1s
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7p
6p
He with 2
electrons
5f
4f
Details
Valence electrons - the electrons in the
outermost energy levels (not d).
 Core electrons - the inner electrons.
 Hund’s Rule - The lowest energy
configuration for an atom is the one
have the maximum number of of
unpaired electrons in the orbital.
2 2
2
 C 1s 2s 2p
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Details
Elements in the same column have the
same electron configuration.
 Put in columns because of similar
properties.
 Similar properties because of electron
configuration.
 Noble gases have filled energy levels.
 Transition metals are filling the d
orbitals
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Exceptions
Ti = [Ar] 4s2 3d2
 V = [Ar] 4s2 3d3
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Cr = [Ar] 4s1 3d5
 Mn = [Ar] 4s2 3d5
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Cu = [Ar] 4s1 3d10
 Half filled orbitals.
 Scientists aren’t sure of why it happens
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7.12 Periodic Trends in Properties
Ionization Energy - The energy
necessary to remove an electron from a
gaseous atom.
 The ionization energy for a 1s electron
from sodium is 1.39 x 105 kJ/mol .
 The ionization energy for a 3s electron
from sodium is 4.95 x 102 kJ/mol .
 Demonstrates shielding.
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Shielding
Electrons on the higher energy levels
tend to be farther out.
 Have to “look” through the other
electrons to “see” the nucleus.
 They are less effected by the nucleus.
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How orbitals differ
The more positive the nucleus, the
smaller the orbital.
 A sodium 1s orbital is the same shape
as a hydrogen 1s orbital, but it is
smaller because the electron is more
strongly attracted to the nucleus.
 The helium 1s is smaller as well.
 This provides for better shielding.
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Periodic Trends
Highest energy electron removed first.
 First ionization energy (I1) is that
required to remove the first electron.
 Second ionization energy (I2) - the
second electron
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Trends in ionization energy
for Mg
•I1 = 735 kJ/mole
•I2 = 1445 kJ/mole
•I3 = 7730 kJ/mole
 The effective nuclear charge increases as
you remove electrons.
 It takes much more energy to remove a
core electron than a valence electron
because there is less shielding.
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Explain this trend
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For Al
•I1 = 580 kJ/mole
•I2 = 1815 kJ/mole
•I3 = 2740 kJ/mole
•I4 = 11,600 kJ/mole
Across a Period
Generally from left to right, I1 increases
because there is a greater nuclear
charge with the same shielding.
 As you go down a group I1 decreases
because electrons are farther away.
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It is not that simple
Zeff changes as you go across a period,
so will I1
 Half filled and filled orbitals are harder
to remove electrons from.
 Here’s what it looks like.
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Atomic number
First Ionization energy
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Atomic number
First Ionization energy
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Atomic number
First Ionization energy
The information it hides
Know the special groups
 It is the number and type of valence
electrons that determine an atoms
chemistry.
 You can get the electron configuration from
it.
 Metals lose electrons have the lowest IE
 Non metals- gain electrons most negative
electron affinities.
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Electron Affinity & Atomic Radii
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Electron affinity - The energy change
associated with the addition of an electron to
a gaseous atom.
X(g) + e-  X-(g)
Affinity tends to increase across a period and
decrease as you go down a group.
Atomic Radii:
Decrease going from left to right across a
period and increase going down a group.
7.13 The Alkali Metals
Doesn’t include hydrogen - it behaves
as a non-metal.
 Decrease in IE.
 Increase in radius.
 Decrease in density.
 Decrease in melting point.
 Behave as reducing agents.
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Reducing ability
Lower IE = better reducing agents.
 Cs>Rb>K>Na>Li
 Works for solids, but not in aqueous
solutions.
 In solution Li>K>Na
 Why?
 It’s the water - there is an energy
change associated with dissolving.
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Hydration Energy
It is exothermic.
 for Li+ -510 kJ/mol
 for Na+ -402 kJ/mol
 for K+ -314 kJ/mol
 Li is so big because of it has a high
charge density, a lot of charge on a
small atom.
 Li loses its electron more easily because
of this in aqueous solutions.
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The reaction with water
Na and K react explosively with water.
 Li doesn’t.
 Even though the reaction of Li has a more
negative H than that of Na and K.
 Na and K melt.
 H does not tell you speed of reaction.
 More in Chapter 12.
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