Chapter 7
Atomic Structure
Dr. S. M. Condren
ELECTROMAGNETIC RADIATION
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Electromagnetic Spectrum
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Electromagnetic Radiation
Electromagnetic wave
• A wave of energy having a frequency
within the electromagnetic spectrum and
propagated as a periodic disturbance of
the electromagnetic field when an electric
charge oscillates or accelerates.
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Electromagnetic Radiation
Electromagnetic wave
• wavelength
• frequency
• amplitude
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Electromagnetic Radiation
Figure 7.1
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Wave motion: wave length and nodes
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Wave Nature of the Electron
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Electromagnetic Radiation
• Waves have a frequency
• Use the Greek letter “nu”,
units are “cycles per sec”
, for frequency, and
l
• Use the Greek letter “lambda”, , for
wavelength, and units are “meters”
• All radiation:
l• = c
• c = velocity of light = 3.00 x 108 m/sec
• Long wavelength --> small frequency
• Short wavelength --> high frequency
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Electromagnetic Radiation
Long wavelength --> small frequency
Short wavelength --> high frequency
increasing
frequency
increasing
wavelength
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Fireworks
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Flame Tests
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The Electric Pickle
• Excited atoms can emit light.
• Here the solution in a pickle is
excited electrically. The Na+ ions
in the pickle juice give off light
characteristic of that element.
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Line Emission Spectrum
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Electromagnetic Radiation
Example: Calculate the frequency, , of
red light that has a wavelength, l, of
700. nm.
 = (1/700. nm)(109nm/1m)(3.00x108m/sec)
= 4.29x1014 s-1
= 4.29x1014 cycles/s
= 4.29x1014 hertz
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Electromagnetic Radiation
Short wavelength -->
high frequency
high energy
Long wavelength -->
small frequency
low energy
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Black Body Radiation
http://www.cbu.edu/~mcondren/C11599/BBvis.mov
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Photoelectric Effect
Experiment demonstrates the particle nature of light.
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Energy of Radiation
Energy of 1.00 mol of photons of red light.
E = h•
= (6.63 x 10-34 J•s)(4.29 x 1014 s-1)
= 2.85 x 10-19 J per photon
E per mol =
(2.85 x 10-19 J/ph)(6.02 x 1023 ph/mol)
= 171.6 kJ/mol
This is in the range of energies that can break
bonds.
Dr. S. M. Condren
Spectra
Line Spectrum
• A spectrum produced by a luminous gas or
vapor and appearing as distinct lines
characteristic of the various elements
constituting the gas.
Emission Spectrum
• The spectrum of bright lines, bands, or
continuous radiation characteristic of and
determined by a specific emitting substance
subjected to a specific kind of excitation.
Absorption Spectrum
• Wavelengths of light that are removed from
transmitted light.
Dr. S. M. Condren
Atomic Line Emission Spectra
and Niels Bohr
Bohr’s greatest contribution
to science was in building
a simple model of the
atom. It was based on an
understanding of the
Niels Bohr
(1885-1962)
SHARP LINE
EMISSION SPECTRA
of excited atoms.
Dr. S. M. Condren
Atomic Spectra and Bohr
Bohr said classical view is wrong.
e- can only exist in certain discrete
orbits — called stationary states.
e- is restricted to QUANTIZED energy
states.
Energy of state = - C/n2
where n = quantum no. = 1, 2, 3, 4, ....
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Bohr Atom
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Energy States
Ground State
• The state of least possible energy in a
physical system, as of elementary
particles. Also called ground level.
Excited States
• Being at an energy level higher than the
ground state.
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Energy Adsorption/Emission
Active Figure 7.11
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Atomic
Spectra and
Bohr
∆E = -(3/4)C
C has been found from experiment (and is now
called R, the Rydberg constant)
R (= C) = 1312 kJ/mol or 3.29 x 1015 cycles/sec
so, E of emitted light
= (3/4)R = 2.47 x 1015 sec-1
and l = c/ = 121.6 nm
This is exactly in agreement with experiment!
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Line Emission Spectra
of Excited Atoms
High E
Short l
High 
Low E
Long l
Low 
Visible lines in H atom spectrum are
called the BALMER series.
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Origin of Line Spectra
Paschen series
Balmer series
Active Figure 7.12
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Atomic Line Spectra
and Niels Bohr
Niels Bohr
(1885-1962)
Bohr’s theory was a great
accomplishment.
Rec’d Nobel Prize, 1922
Problems with theory —
• theory only successful for H.
• introduced quantum idea
artificially.
• So, we go on to QUANTUM or
WAVE MECHANICS
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Quantum or Wave Mechanics
Schrodinger applied idea of ebehaving as a wave to the problem
of electrons in atoms.
He developed the WAVE
EQUATION
Solution gives set of math
expressions called WAVE
E. Schrodinger
FUNCTIONS, 
1887-1961
Each describes an allowed energy
state of an eQuantization introduced naturally.
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WAVE FUNCTIONS, 
• is a function of distance and two
angles.
• Each  corresponds to an
ORBITAL — the region of space
within which an electron is found.
•  does NOT describe the exact
location of the electron.
• 2 is proportional to the probability
of finding an e- at a given point.
Dr. S. M. Condren
Uncertainty Principle
W. Heisenberg
1901-1976
•Problem of defining nature
of electrons solved by W.
Heisenberg.
•Cannot simultaneously
define the position and
momentum (=m*v) of an
electron.
•We define e- energy
exactly but accept limitation
that we do not know exact
position.
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Types of Orbitals
s orbital
p orbital
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d orbital
Orbitals
• No more than 2 e- assigned to an orbital
• Orbitals grouped in s, p, d (and f) subshells
s orbitals
also
p orbitals
d orbitals
f orbitals
Dr. S. M. Condren
s orbitals
p orbitals
d orbitals
f orbitals
p orbitals
d orbitals
f orbitals
1
3
5
7
2
6
10
14
s orbitals
No.
orbs.
No. e-
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QUANTUM NUMBERS
The shape, size, and energy of each orbital is
a function of 3 quantum numbers:
n (principal) =>
l (angular) =>
ml (magnetic) =>
shell
subshell
designates an orbital
within a subshell
s (spin) => designates the direction of spin
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QUANTUM NUMBERS
Symbol
ValuesDescription
n (principal)
1, 2, 3, ..
l (angular)
ml (magnetic)
s (spin)
Orbital size and energy
where E = -R(1/n2)
0, 1, 2, .. n-1 Orbital shape or type
(subshell)
-l..0..+l
Orbital orientation
# of orbitals in
subshell = 2 l + 1
-1/2 or +1/2 Direction of spin of electron
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Types of
Atomic
Orbitals
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Atomic Orbitals
• Types of orbitals found in the known
elements: s, p, d, and f
• schools play defensive football
• Packer version: secondary pass defense
fails
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S Orbitals
1s
2s
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3s
p Orbitals
The three p orbitals lie 90o apart
in space
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2px Orbital
3px Orbital
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d Orbitals
3dxy Orbital
3dxz Orbital
3dx2- y2 Orbital
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3dyz Orbital
3dz2 Orbital