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Chapter 11: Modern Atomic Theory or
Quantum Mechanics ROCKS!!!
I. Interaction of Light and Matter
A. Properties of Light
B. Emission of Light
C. Bohr Model of the Atom
II. Quantum Mechanical Model of the Atom
A. The Hydrogen Atom
1. Atomic Orbitals and Quantum #’s
2. Shapes of Orbitals
B. Electron Configurations of Elements
III. Atomic Properties and the Periodic Table
A. Atomic Size
B. Ionization Energy
C. Metals and Non-Metals
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I. Interaction of Light and Matter
A. Properties of Light
Properties of Light (Electromagnetic Radiation)
?
http://abyss.uoregon.edu/~js/glossary/wave_particle.html
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Terms Used to Describe Waves
1. Wavelength (λ) – distance between successive
peaks in a wave
2. Frequency (ν) – the number of wavelengths that
pass a given point per second
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Terms Used to Describe Waves Cont’d
3. Amplitude – measure of the intensity of light
4. Energy (E)
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Range of Wavelengths
Electromagnetic Spectrum
High Energy
gamma rays > x-rays > ultraviolet light > visible light >
Low Energy
infrared light > microwaves > radio waves
Visible Light
GAMMA RAYS HURT!
HULK THIRSTY!
HULK WANT PEPSI!
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C. Bohr Model of the Atom
1. Electrons move in circular orbits around the nucleus.
2. There are only certain allowed orbitals.
3. In order for an electron to move between orbitals it must
gain/lose the right magnitude of energy.
Absorption and Emission of Light from Atoms
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II. Quantum Mechanical Model of the Atom
A. The Hydrogen Atom
De Broglie – If light is particle-like and wave-like
then perhaps all matter has both
types of properties
Electrons – have both wave-like and particle-like
properties!
Schrödinger Equation
Orbital –
3-dimensional – 3 variables
EzyxVdz
d
dy
d
dx
d
m
),,(
2 2
2
2
2
2
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m = mass of particle
ђ = planck’s constant / 2
Me H-atom
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A Closer Look at Orbitals - All atoms have the same general pattern of “living
spaces for electrons.
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B. Electron Configurations of Elements
Atomic Structure and the Periodic Table
a) Aufbau Principle (building up) - electrons are added to an atom starting with the
lowest energy orbitals first
b) Pauli Exclusion Principle - 2 electrons can fit in each orbital
- electrons have spin
c) Hund’s Rule - the electron configuration with the lowest energy
has the maximum number of unpaired electrons
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Electron configurations for elements
1H 1 e-
2He 2 e-
3Li 3e-
4Be 4e-
5B 5e-
1s 2s 2p
1s 2s 2p
1s 2s 2p
1s 2s 2p
1s 2s 2p
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6C 6e-
7N 7e-
8O 8e-
9F 9e-
10Ne 10e-
11Na 11e-
3s 3p 1s 2s 2p
1s 2s 2p
1s 2s 2p
1s 2s 2p
1s 2s 2p
1s 2s 2p
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Sum Up (Categories of Electrons) a) Valence electrons - reactive electrons in the
outermost energy level
(highest n)
- usually the electrons in the
highest or outermost s & p
orbitals
b) Inner core electrons - unreactive lower energy level
electrons
Helpful Simplification
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Possible Questions
What is the full electron configuration for
Calcium?
What is the abbreviated electron configuration
for Br?
What is the abbreviated electron configuration
for Zr?
How many d electrons does Mo have?
How many valence electrons does Mo have?
How many unpaired electrons?
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Electron Configurations and Ions
Trend - When they react to form ions, atoms lose
electrons until they have 0 valence electrons, or
they gain electrons until they have 8 valence
electrons
Magnesium
Fluorine
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Metals with Electrons in d-orbitals
Trend - Metals tend to lose electrons from the
outermost s & p orbitlas to form ions (sometimes
d-electrons can also be lost)
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Example Problems
Which element in group 1 most easily loses
electrons? Why? Which element in group 1
is the most reactive?
Which element in group 7 will most easily
lose electrons? Why? Which element in
group 7 is the most reactive?
Arrange each of the following elements in
order of increasing atomic size.
Sn, Xe, Rb, Sr