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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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1
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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2
Figure 7.1 The Nature of Waves
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Wavelength and frequency can be interconverted.
 = c/
 = frequency (s1)
 = wavelength (m)
c = speed of light (m s1)
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4
Planck’s Constant
Transfer of energy is quantized, and can only
occur in discrete units, called quanta.
E = h =
hc

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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Energy and Mass
Energy has mass, so does yo mama
E = mc2
E = energy
m = mass
c = speed of light
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Energy and Mass
Ephoton =
mphoton
hc

h
=
c
(Hence the dual nature of light.)
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Wavelength and Mass
de Broglie’s Equation
h
 =
m
 = wavelength, in m
h = Planck’s constant, 6.626  1034 J s =
kg m2 s1
m = mass, in kg
 = velocity in m/s
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8
Atomic Spectrum of Hydrogen
Continuous spectrum: Contains all the
wavelengths of light.
Line (discrete) spectrum: Contains only
some of the wavelengths of light.
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9
The Bohr Model
The electron in a hydrogen atom moves around the
nucleus only in certain allowed circular orbits.
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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10
The Bohr Model
Ground State: The lowest
possible energy state for an atom
(n = 1).
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11
Energy Changes in the Hydrogen
Atom
E = Efinal state  Einitial state
hc
 =
E
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12
Quantum Mechanics
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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13
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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Probability Distribution
 square
of the wave function
 probability of finding an electron at a given
position
Radial probability distribution is the
probability distribution in each spherical
shell.
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15
Figure 7.12
Radial Probability Distribution
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16
Quantum Numbers (QN)
1.
Principal QN (n = 1, 2, 3, . . .) - related to size and
energy of the orbital.
2.
Angular Momentum QN (l = 0 to n  1) - relates to
shape of the orbital.
3.
Magnetic QN (ml = l to l) - relates to orientation
of the orbital in space relative to other orbitals.
4.
Electron Spin QN (ms = +1/2, 1/2) - relates to the
spin states of the electrons.
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Figure 7.13
Two
Representations
of the
Hydrogen 1s,
2s, and 3s
Orbitals
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18
Figure 7.14
The Boundary Surface
Representations of All Three 2p
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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Figure 7.18 Orbital Energy Levels for the
Hydrogen Atom
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21
Figure 7.20
A Comparison of the Radial Probability
Distributions of the 2s and 2p Orbitals
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22
Polyelectronic Atoms
•
•
•
•
•
Shielding
Hydrogen orbitals: degenerate
Hydrogenlike orbitals: NOT degenerate!
Ens < Enp < End < Enf …etc.
Penetration: e- tunneling
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23
Figure 7.21
The Radial
Probability
Distribution for the
3s, 3p, and 3d
Orbitals
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24
Figure 7.23
Mendeleev’s Early Periodic
Table, Published in 1872
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25
History of the Periodic Table
Dobereiner: triads
Newlands: octaves
Meyer / Mendeleev: arrangements by atomic
masses.
Theory → prediction
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26
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.
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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.
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28
Figure 7.36Special Names for
Groups in the Periodic Table
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Valence Electrons
The electrons in the outermost principle
quantum level of an atom.
Atom
Valence Electrons
Ca
2
N
5
Br
7
Inner electrons are called core electrons.
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Figure 7.24
The Electron Configurations in the Type of
Orbital Occupied Last for the First 18
Elements
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31
Figure 7.25
Electron Configurations for
Potassium Through Krypton
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Figure 7.26
The Orbitals Being Filled for Elements in
Various Parts of the Periodic Table
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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 7.27 The Periodic Table With Atomic
Symbols, Atomic Numbers, and Partial Electron
Configurations
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35
Figure 7.30
The Positions of the Elements
Considered in Sample Exercise 7.7
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36
Ionization Energy
The quantity of energy required
to remove one electron from the
gaseous atom or ion.
X(g) → X+(g) + eCopyright©2000 by Houghton
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Periodic Trends
First ionization energy:
increases from left to right across a
period;
decreases going down a group.
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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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1s22s22p63s1
1s22s22p6
1s22s22p63s2
Which atom has the largest first I.E.?
Which one has the smallest second I.E.?
Explain.
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40
Electron Affinity
The energy change associated with
the addition of an electron to a
gaseous atom or ion.
X(g) + e  X(g)
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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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Figure 7.35
Atomic Radii
for Selected
Atoms
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43
Periodic Trends
Atomic Radii:
decrease going from left to right
across a period;
increase going down a group.
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44
Figure 7.34
The Radius of an Atom
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Alkali Metals
Expected trend in reducing ability:
Cs > Rb > K > Na > Li
But in water,
Li > K > Na
Hydration energy
Charge density
What we observe in water:
K > Na > Li
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