Download Lecture Outline Chp 11

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
?
http://abyss.uoregon.edu/~js/glossary/wave_particle.html
Properties of Light (Electromagnetic Radiation)
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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
GAMMA RAYS HURT!
HULK THIRSTY!
HULK WANT PEPSI!
Electromagnetic Spectrum
High Energy
gamma rays > x-rays > ultraviolet light > visible light >
Low Energy
infrared light > microwaves > radio waves
Visible Light
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Chemistry with Light
UV light –
Vis light –
IR light –
Microwaves –
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B. Emission of Light
What happens when atoms gain energy?
What happens when atoms lose energy?
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Information from light emission
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What is really going on?
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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
 2  d 2 d 2 d 2 

 2  2  2   V ( x, y, z )  E
2m  dx
dy
dz 
m = mass of particle
ђ = planck’s constant / 2
Me
H-atom
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Locations Where There is High Probability of Finding Matt
Orbitals
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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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Shape of Orbitals
1. s-orbitals
1s
2s
2. p-orbitals
2p
3. d-orbitals
3d
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1s
2s
3s
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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Homes for the electron in Hydrogen
N=3
N=2
N=1
1s
2s
2p
3s
3p
3d
Homes for the electron in Helium
N=3
N=2
N=1
1s
2s
2p
3s
3p
3d
Homes for the electron in Lithium
N=3
N=2
N=1
1s
2s
2p
3s
3p
3d
Homes for the electron in Nitrogen
N=3
N=2
N=1
1s
2s
2p
3s
3p
3d
Electron configurations for elements
1H
 1 e1s
2s
2p
1s
2s
2p
1s
2s
2p
1s
2s
2p
1s
2s
2p
He

2
e
2
Li

3e
3
Be

4e
4
5B
 5e-
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6C
 6e-
7N
 7e-
8O
 8e-
9F
 9e-
1s
2s
2p
1s
2s
2p
1s
2s
2p
1s
2s
2p
1s
2s
2p
Ne

10e
10
11Na
 11e1s
2s
2p
3s
3p
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The Periodic Table and Electron Configurations
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
Electron Configurations (Period 3)
Period 4 and Below
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The Periodic Table
Period 5
Period 6
Period 7
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“Easily” Explained Exceptions
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s1
s2
p1
d1
exception
p2 p3 p4 p5
d2 d3 d4 d5 d6 d7 d8 d9 d10
Lots of exceptions
p6
Sum Up (Categories of Electrons)
a) Valence electrons
-
-
b) Inner core electrons -
reactive electrons in the
outermost energy level
(highest n)
usually the electrons in the
highest or outermost s & p
orbitals
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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III. Atomic Properties and the Periodic Table
A. Atomic Size
Down a Group
Across a Period
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Comparing the electron configuration and size of Li and Na
Li
Na
Comparing the electron configuration and size of Mg and Al
Mg
Al
B. Ionization Energy
Ionization energy -
Down a Group
Across a Period
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Formation of Binary Ionic Compounds
Metal + Non-Metal  Ionic Compound
Metals
Non-metals
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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
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