Chapter 4
ARRANGEMENT OF
ELECTRONS IN
ATOMS
Visible Light
We are all familiar with
light but what is “visible”
is just a very, very small
portion of the
electromagnetic spectrum
What colors make up the
rainbow?
Red, Orange, Yellow,
Green, Blue, Indigo,
Violet (ROYGBIV)
The E-M
Spectrum
X-rays
Ultraviolet
( not
Radio
Rays
Microwaves
Gamma
Infrared
Rays
(Cancerous
as Radio,
harmful
in-large
(TV,
other
(Cooking,
(Very
(Heat,
Harmful
/
doses;
sunburn;
small doses
communications)
communications)
Cancerous)
communication)
medical
black lights)
scanning)
The Development of a New Atomic Model
 Wavelength () - length of one complete wave
 Frequency () - # of waves that pass a point during a
certain time period

hertz (Hz) = 1/s
 Amplitude (A) - distance from the origin to the
trough or crest
Waves

crest
A

greater
origin
amplitude
(intensity)
A
trough
greater
frequency
(color)
The Electromagnetic Spectrum
Decreasing wavelength
Increasing frequency
Increasing photon energy
AM radio
V
i
s
i
b
l
e
Television channels
Short wave
radio
FM
radio
Radar
Microwave
Radio Waves
L
i
g
h
t
infrared
R
O
Red
Orange
Y
Yellow
G
Green
UV
Rays
Gamma Rays
X- Rays
B
I
V
Blue
Indigo
Violet
Electromagnetic Spectrum
 Frequency & wavelength are inversely
proportional
c = 
c: speed of light (3.00  108 m/s)
: wavelength (m, nm, etc.)
: frequency (Hz)
Electromagnetic Spectrum
 EX: Find the frequency of a photon with a
wavelength of 434 nm.
GIVEN:
WORK:
=?
=c
 = 434 nm

= 4.34  10-7 m
8 m/s

=
3.00

10
c = 3.00  108 m/s
4.34  10-7 m
 = 6.91  1014 Hz
So why is the electromagnetic spectrum so
important to chemistry?
• Why is the steel emitting light when it is heated?
• We take it for granted that when things get hot
they turn red then orange and finally white; but
that isn’t good enough any more
Black Body Radiation Colors
Black Body Radiation Colors
°C
Subjective color
480
faint red glow
580
dark red
730
bright red, slightly orange
930
bright orange
1100
pale yellowish orange
1300
yellowish white
>
1400
white (yellowish if seen from a
distance through atmosphere)
So why is the electromagnetic spectrum so
important to chemistry?
 Incandescence is heat made visible – the
process of turning heat energy into light
energy.
 Our usage of terms like "red hot," "white
hot," and so on, is part of the color
sequence black, red, orange, yellow, white,
and bluish white, seen as an object is
heated to successively higher
temperatures.
So why is the electromagnetic spectrum so
important to chemistry?
 The light produced consists of
photons emitted when atoms and
molecules release part of their thermal
vibration energy.
 For increasing temperatures, the
sequence of radiated colors is: black,
red, orange, yellow-white, bluishwhite.
Heat and Light
 Planck (1900)

Observed - emission of light from hot objects

Concluded - energy is
emitted in small, specific
amounts (quanta)

Quantum - minimum amount of energy change
Energy and Light
The energy of a photon is proportional to
its frequency.
E = h
E: energy (J, joules)
h: Planck’s constant (6.6262  10-34 J·s)
: frequency (Hz)
Energy and Light
 EX: Find the energy of a red photon with a
frequency of 4.57  1014 Hz.
GIVEN:
WORK:
E=?
E = h
 = 4.57  1014 Hz
-34 J·s)
E
=
(6.6262

10
h = 6.6262  10-34 J·s
(4.57  1014 Hz)
E = 3.03  10-19 J
Niels Bohr and the Bohr model of
the atom
Bohr hypothesized that
instead of haphazardly
orbiting the nucleus,
electrons had clearly
defined orbits – very
similar to the planetary
orbits circling our sun
His model is (cleverly)
named the Planetary Model
Niels Bohr
Bohr Model (1913)
Bohr’s Proof
 Bohr said this: If you assume that the electrons
have clearly defined orbits that are congruent to
the energy levels…
Bohr’s Proof
 … then when an electron gets “excited” it jumps to
a higher energy level. When it “relaxes” it emits a
certain wavelength of light.
• Bohr showed the
energy of an
electron in an atom
is quantized, which
means it has a
particular numerical
value, not some
arbitrary number.
Excitation of Hydrogen Atoms
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 328
Return to Ground State
Bohr’s Proof
1
λ
= 1.097373 x 107 m-1
1
nr2
-
1
ne2
n=7
n=6
n=5
n=4
Paschen Series (ir)
n=3
Balmer Series (vis and uv)
n=2
Lyman Series (uv)
n=1
Emission Spectrum of an Element
410 nm 434 nm
486 nm
656 nm
1 nm = 1 x 10-9 m = “a billionth of a meter”
1 nm = 1 x 10-9 m = “a billionth of a meter”
Continuous and Line Spectra
Hydrogen to Steel
 If Hydrogen emits 4 distinct wavelengths of light
when its one electron is excited what can we
extrapolate to that of steel which is made mostly of
iron?
 http://jersey.uoregon.edu/vlab/elements/Element
s.html
Flame Emission Spectra
Photographs of flame tests of burning wooden splints soaked in different salts.
methane gas
wooden splint
sodium ion
calcium ion
copper ion
strontium ion
Include link to web page
http://www.unit5.org/christjs/flame%20tests.htm
Fireworks
Composition of Fireworks
 Gunpowder
 Sulfur, charcoal, potassium nitrate (saltpeter)
 Salts (to give color)
 Red = lithium
 Green = copper
Good News Bad News
 Good News
 Bad (Frustrating) News
 Bohr’s Model works and
 Lots of Math
moves us along in the
development of the
Atomic Theory
 End of this little unit
 Everything I taught you
only works for Hydrogen
and therefore is
completely wrong and
obsolete.
Check for Understanding
c= λν E=hν
c=3.0 x108 m/s h=6.626 x 10-34 J s
 What is the frequency of a radar photon with an
energy of 7.2 x 10-24 J?
 What is the frequency of light having a wavelength
of 6.20x10-7m?
Models of the Atom
e
e
+
e
+
e
+
+
e
+e
+
e
e
+ e + e
Dalton’s
Greek model
model
(400
(1803)
B.C.)
Thomson’s plum-pudding
model (1897)
Bohr’s model
(1913)
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125
-
- +
Rutherford’s model
(1909)
Your Current View of the Atom
nucleus
electrons
Again… so why is it so important to
chemistry?
 Einstein (1905)

Observed - photoelectric effect
Again… so why is it so important to
chemistry?
 Einstein (1905)

Concluded - light has properties of both waves and
particles
“wave-particle duality”

Photon - particle of light that carries a quantum of energy
Quantum Mechanical Model
Modern atomic theory describes the
electronic structure of the atom as the
probability of finding electrons within certain
regions of space (orbitals).
Modern View
 The atom is mostly empty space
 Two regions
 Nucleus


protons and neutrons
Electron cloud

region where you might find an electron
 Also called the electron cloud
model
Modern View of Atom
Electrons can only be at
specific energy levels,
NOT between levels.
Excited state
e-
Ground state
Models of the Atom
e
e
+
e
+
e
+
+
e
+e
+
e
e
+ e + e
Dalton’s
Greek model
model
(400
(1803)
B.C.)
Thomson’s plum-pudding
model (1897)
Bohr’s model
(1913)
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125
-
- +
Rutherford’s model
(1909)
Charge-cloud model
(present)
Electrons as Waves
 Louis de Broglie (1924)

Applied wave-particle theory to e-

e- exhibit wave properties
QUANTIZED WAVELENGTHS
C. Johannesson
Quantum Mechanics
 Heisenberg Uncertainty Principle

Impossible to know both the velocity and position of an
electron at the same time
C. Johannesson
Quantum Mechanics
 Schrödinger Wave Equation (1926)

finite # of solutions  quantized energy levels

defines probability of finding an e-
Ψ 1s 
C. Johannesson

1 Z 3/2 σ
π a0
e
Quantum Theory
 quantum theory Describes





mathematically the wave properties of
electrons and other small particles
orbital- a region of an atom in which there is a high
probability of finding electrons
Today’s atomic model predicts quantized, or
particular energy levels for electrons.
does not describe the exact path or location electrons
take or can be found around the nucleus
concerned with the probability, or likelihood, of
finding an electron in a certain position
Two electrons can occupy each orbital, also called an
electron cloud.
Quantum Numbers
 Four Quantum Numbers:

Specify the “address” or “seat” of each electron in an
atom
UPPER LEVEL
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Quantum Numbers
1. Principal Quantum Number ( n )

Energy level (ladder rungs)
1s

Size of the orbital

Positive integer
2s
3s
Quantum Numbers
1. Principal Quantum Number

> number, further away from the nucleus
1-
right next to the nucleus
3- further away from nucleus

1s
> number, higher the energy level
n
= 2 greater energy level than n = 1
these electrons have more energy than
electrons in the n = 1 level
2s
3s
Quantum Numbers
2. Angular Momentum Quantum # ( l )

Energy sublevel (orbital)

Shape of the orbital

Often represented by letters than numbers
s
p
d
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
f
Quantum Numbers
y
y
z
x
z
x
px
y
z
x
pz
py
d-orbitals
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 336
Quantum Numbers
 Orbitals combine to form a spherical shape.
2s
2px
2py
2pz
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Quantum Numbers
3. Magnetic Quantum Number ( ml )

Orientation of orbital

Specifies the exact orbital within each sublevel
Shapes of s, p, and d-Orbitals
Quantum Numbers
4. Spin Quantum Number ( ms )


Electron spin  +½ or -½
An orbital can hold 2 electrons that spin in opposite
directions.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Maximum Capacities of Subshells
and Principal Shells
n
1
2
l
0
0
1
0
1
2
0
1
2
3
Subshell
designation
s
s
p
s
p
d
s
p
d
f
Orbitals in
subshell
1
1
3
1
3
5
1
3
5
7
Subshell
capacity
2
2
6
2
6
10
2
6
10 14
Principal shell
capacity
2
8
Hill, Petrucci, General Chemistry An Integrated Approach 1999, page 320
3
18
4
...n
32
...2n2
Filling Rules for Electron Orbitals
Aufbau Principle: Electrons are added one at a time to the lowest
energy orbitals available until all the electrons of the atom
have been accounted for.
Pauli Exclusion Principle: An orbital can hold a maximum of two electrons.
To occupy the same orbital, two electrons must spin in opposite
directions.
Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum
number of unpaired electrons results.
*Aufbau is German for “building up”
Diagonal Rule
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
General Rules
 Pauli Exclusion Principle

Each orbital can hold TWO electrons with opposite spins.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
General Rules
 Hund’s Rule

Within a sublevel, place one electron per orbital
before pairing them.
WRONG
RIGHT
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Orbital Diagrams and Electron Configurations
 Orbital diagrams
 Show how electrons are distributed within sublevels
 Electrons represented by an arrow
 Orbital is represented by a box
 Direction of spin represented by direction of arrow
 Electron configuration
 Abbreviated form of orbital diagram
Orbital Diagrams and Electron Configurations
H
1 e-
Orbital diagram
↑
1s
Electron configuration
1s1
Orbital Diagrams and Electron Configurations
 Orbital Diagram
O
8e-
1s
2s
• Electron Configuration
2
2
4
1s 2s 2p
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
2p
Orbital Diagrams and Electron Configurations
 Orbital Diagram
Ne
10e-
1s
2s
2p
• Electron Configuration
2
2
6
1s 2s 2p
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Noble Gas Configuration
A neon's electron configuration (1s22s22p6)
B
third energy level
[Ne] 3s1
C
D
one electron in the s orbital
orbital shape
Na = [1s22s22p6] 3s1
electron configuration
Electron Configuration
16
S
32.066
 Longhand Configuration
S 16e- 1s2 2s2 2p6 3s2 3p4
Core Electrons
Valence Electrons
• Shorthand Configuration
S
16e
2
4
[Ne] 3s 3p
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