Download Electron Structure of Atoms Notes

Chapter 6
Electron structure of atoms
1
6.1 Electromagnetic Radiation

Light that we can see is
visible light which is a type
of electromagnetic
radiation.

Radiant energy is energy
that carries energy that acts
like a wave and travels
through space at the speed
of light.
Earth’s Radiant Energy
2
c = speed of light
3.0 x 108 m/s
3
Wave characteristics

Wavelength: λ, lambda
Distance between peaks
or troughs in a wave
Frequency: ν, nu
number of waves, per
second that pass a point
in one second.
Speed: you know this one.
4

Which color has the
highest frequency?

Lowest frequency?

Largest wave length?

Smallest wavelength?
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Electromagnetic Spectrum
6
Flame testing

http://www.sciencefriday.com/videos/watch/1
0227
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Relationship between λ and ν

Wavelength and frequency are inverses of each
other.

λv = c

λ = wavelength in meters (m)
ν = frequency in cycles per second (1/s or s-1 or Hertz)
c = speed of light 3.0 x 108 m/s


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Try one!

The red wavelength emitted from red
fireworks is around 650 nm and results when
strontium salts are heated. Calculate the
frequency of the red light of this wavlength.

λv=c
λ = (6.50 x 102 nm) = 6.50 x 10-7 m
v = 4.61 x1014 s-1 or Hz


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6.2 Planck’s Constant
Max Planck discovered
that energy could be
gained or lost in
multiples of a constant
(h) times its frequency
(ν).
h = 6.626 x 10-34 J *s
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Quantized Energy
Thus energy is
quantized or in steps
or packages.
Energy can only be
transferred as a whole
package or quantum.
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Solving equations with
Planck’s
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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Calculating energy lost

The blue color in fireworks is the result of
heated CuCl at 1200 °C. Then the compound
emits blue light with a wavelength of 450 nm.
What is the increment of energy (quantum)
that is emitted at 4.50 x 102 nm by CuCl?
13
Answer
ΔE = hν

v = 3.0 x 108 m/s
4.50 x 10-7 m
= 6.66 x 1014 s-1
v = c/λ
(6.626 x 10-34 J *s) x (6.66 x 1014 s-1)
= 4.41 x10-19J (quantum energy lost in this increment)
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photons

Einstein took Planck’s idea
a step further and proposed
that electromagnetic
radiation was quantized into
particles called photons
(light).

The energy of each photon
is given by the expression:
Ephoton = hν = hc/λ
15
Dual Nature of Light
Light can behave as if it
consists of both waves
and particles.
Thus light energy has
mass

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Old-ie but good-ie

Energy has mass
E



= mc2
E = energy
m = mass
c = speed of light
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6.4 The Behavior of the wave
De Broglie
We can calculate the
wavelength of an e-.
h
 =
m
 = wavelength, in m
h = Planck’s constant,
6.626  1034 J s
v = velocity
m = mass in kg
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Question

Compare the wavelength for an electron
(mass = 9.11 x10 -31 kg) traveling at a speed
of 1.0 x107 m/s with that of a ball (mass =
0.10 kg) traveling at 35 m/s
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Answer
h
 =
m
Electron wavelength = 7.27 x 10 -11 m
ball wavelength = 1.9 x 10 -34 m
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ν = nu = frequency
v = velocity
your book uses μ for velocity
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Homework
Chang: pg 303 #’s 1, 2, 7, 9, 15, 20
BL: Pg 230
 1, 3, 5, 7, 10, 13, 15,19
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Waves = Electrons

Planck and Einstein proved that electrons in
atoms act like waves of light

By understanding waves we can learn about
the properties of electrons

The study of the properties of electrons is
Quantum Mechanics
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Bohr Model

Proposed the first
theory on atom location
and movement

His proposal was a little
bit right and a lot
wrong…BUT we give
him props just the same
Nelis Bohr
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Bohr Model
Where he was right…
1.
2.
Electrons exist in
certain discrete energy
levels, which are
described by quantum
numbers
Energy is involved in
moving electrons from
one energy level to
another
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Heisenberg Uncertainty Principle

Blew the Bohr model
out of the water.

It states that we can
only know so much
about the exact position
and momentum of an
electron.
…And the electron cloud
is born
Werner Heisenberger
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Probability

Bohr Model



Probability distribution
Orbits
Electron Cloud


Radial probability
distribution
Orbitals
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6.5 Quantum model of an atom

Compared the
relationship between
the electron and the
nucleus of an atom to
that of a standing or
stationary wave.
The functions of these
waves tell us about the
electrons location and
energy.
Erwin SchrÖdinger
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Schrödinger's Cat
He proposed a scenario with a
cat in a sealed box, where
the cat's life or death was
dependent on the state of a
subatomic particle.
According to Schrödinger,
the Copenhagen
interpretation implies that
the cat remains both alive
and dead until the box is
opened.
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Quantum numbers!!!!!
Quantum numbers describe various properties of the
electrons in an atom.
There are 4 quantum numbers
Principal quantum number (n)
Azimuthual quantum number (angular momentum) (ℓ)
Magnetic quantum number (mℓ)
Electron spin quantum number (ms)
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Principal quantum number (n)

Integral values 1,2,3,4,5,6,7
Related to the size and energy of the orbital

Referred to as the shell or energy level

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Principal quantum number (n)

As n increases energy
increases and orbital
size increases
because the electrons are
farther away from the
nucleus and less tightly
bound to the positive
protons.
n=1
n=4
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Angular momentum quantum number
(ℓ)
Integral numbers with values
from 0 to n-1
if n = 3 possible ℓ values are
0,1,2
Sometimes referred to as the
“sub shell” number
Defines the shape of the
orbital.
ℓ
Orbital
shape
0
s
1
p
2
d
3
f
4
g
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Shape of orbitals
ℓ
Orbital
shape
0
s
1
p
2
d
3
f
4
g
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Magnetic quantum number (mℓ)
Integral values from ℓ to -ℓ
including zero
If ℓ = 2
Then mℓ = 2, 1, 0, -1, -2
Relates to the orientation
of the orbital in the
atom.
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Electron spin quantum number
(ms)

can only have one of
two values
+1/2 or -1/2
+½
-½
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Re cap

A collection of orbitals with the same n value
is called an electron shell.
EX: all orbitals that have n =3 are in the third
shell.
A collection of orbitals with the same n and ℓ
values are in the same sub shell
EX: 2s, 2p
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Sublevel
Shape
(ℓ)
Orbital
s
S
1
1 x 2 = 2e
0, 1
sp
p
3
3 x 2 = 6e
3
0, 1, 2
spd
d
5
5 x 2 = 10e
4
0, 1, 2, 3
spdf
f
7
7 x 2 = 14e
5
0, 1, 2, 3, 4
spdfg
g
9
9 x 2 = 18e
6
0, 1, 2, 3, 4, 5
spdfgh
h
11
11 x 2 = 22e
7
0, 1, 2, 3, 4, 5, 6
spdfghi
i
13
13 x 2 = 26e
Principle
Quantum #
n
# of possible l
values
1
0
2
number (ℓ)
Electron Capacity
Note: In order for the d orbital to be filled the s and p orbitals must
be filled.
Table 6.2 page 214
41
question
For the principle quantum level n = 5
Determine the number of allowed sub shells (ℓ)
and give the number and letter designation of
each

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Answer

Recall: Angular momentum quantum
Integral numbers with values from 0 to n-1
n=5
ℓ = 0 or s, 1 or p, 2 or d, 3 or f, 4 or g
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Nomenclature
n value
ℓ value
number of electrons in
orbital
Y
2p
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Sorting our the numbers

Orbitals with the same n value are in the same shell.
Ex: n = 3 is the third shell
One or more orbitals with the same set of n and ℓ values are
in the same sub shell
Ex: n = 3 ℓ= 2 3d sub shell
n=3ℓ=1
3p sub shell
45
Homework

Chang pg 305 #’s 43, 44. 46, 47, 48, 52,
53,56, 57, 63,

BL:Pg 232 41, 43, 45, 46
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Pauli exclusion principle

In a given atom no
electrons can have the
same 4 quantum number

So when we put more
than one electron in an
orbital
we must
alternate the spin. Thus
ms = +1/2 -1/2
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48

Example of Pauli Exclusion Principal
Quantum numbers for 2s2
2s
n
2
l
0
ml
0
ms
+1/2
2s
2
0
0
-1/2
When ever possible electrons will prefer to have a positive spin.
In this case this orbital will only hold 2 e- so one must be
negative
49
Question ?
What would the 4 quantum
3
numbers be for 3p ?
Note: all electrons have positive spin We will
get to why in a minute
50
Answer
n
l
ml
ms
3p
3
1
0
+1/2*
3p
3
1
1
+1/2*
3p
3
1
-1
+1/2*
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Homework
Page 232-33
#’s 52, 53, 54, 56

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Electron configuration

The order in which electrons are distributed
to orbitals

We need to have rules for how we distribute
electrons. Other wise all the electrons would
be in the 1s orbital because it has the lowest
energy
(e- ♥ ground state)

53
Rule 1:
Aufbau Principle “building up”

Shells fill based on their
energy level.

Lower energy shells fill
first followed by high
energy shells.
START
54
H: 1s1
Li: 1s2 2s1
He: 1s2
55
s
p
d
f
56
How to write EC?
3 electrons
Li
1s
Orbital Diagram
2s
2
1
1s 2s
electron configuration 57
Question ?
 What
is the electron
configuration for Carbon?
58
Answer
Carbon has
6 electrons
C
1s2
2s2
2p2
59
Hund’s Rule: “the grocery line rule”

Electrons are distributed among the orbitals
or a sub shell in a way that gives the
maximum number of unpaired electrons.
C
1s2
2s2
2p2
60
Question

Write the orbital diagrams and electron
configurations for the electron configurations
of each element.

Nitrogen
Oxygen
Fluorine
Potassium



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Answer
62
Valence Electrons
The electrons in the outermost principle
quantum level of an atom. Ve- = to group #
Atom
Ca
N
Br
Ve2
5
7
Location
4s
2s 2p
4p3d
Inner electrons are called core electrons.
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Short and Sweet!
Writing the EC for Carbon is one thing but
Xenon (54e-), Argon (18e-)?
To write the condensed EC look to the noble
gas BEFORE your element.
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Condensed Form Example



Cs = 55 eNoble gas before it is Xenon Xe= 54e[Xe]
We still need 1 more e- so we write it in
[Xe] 6s1
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Xe
Cs
69
Question?
 What
is the condensed electron
configuration for Selenium?
70
Answer
Se = 34 e[Ar] 4s2 3d10 4p4
71
Ar
Se
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EXCEPTION ALERT!!!



Memorize the EC of Copper and Chromium. They
are exceptions to our rules due to stability
Chromium [Ar] 4s13d5
Copper [Ar] 4s13d10
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Homework

Pg 233

#’s 59, 60, 61,62, 63, 65
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