Download Electron Configurations

Ch. 4 “Electron Configurations
Quantum Mechanics Made Simple!
In chapter 3, we began our
historical journey though the
development of atomic theory.
Rutherford’s ‘Nuclear Atom” was
more useful than Dalton’s or
Thomson’s models because it was
able to explain the results of the
“alpha particle scattering
experiment”.
As more evidence was accumulated,
it, too, was replaced by a better
model!
What we know so far…
An atom consists of a
 nucleus
(of
protons and neutrons)
 electrons
in space outside the
nucleus.
Electron cloud
Nucleus
Much of our understanding of how
electrons behave in atoms comes from
studies of how light interacts with
matter.
As you know, light travels through space
& is a form of radiant energy.
This is what makes you feel warm as you
stand in sunlight!
How light travels through space has
been a major source of debate for
centuries!
1600’s, Isaac Newton suggested that
light was made of tiny particles.
Newton used a glass prism to refract
(bend) sunlight (white light) into a
continuous spectrum.
Continuous Spectrum – a complete array
of colors from red to violet. (a rainbow!)
ROYGBIV (or VIBGYOR)
This process is called Diffraction –
passing white light through a diffraction
grating to produce a continuous
spectrum.
1600’s, Christian Huygens (Dutch)
suggested that light consists of waves
(rather than particles) - “Wave Model of
Light”
He thought light travels away from its source
the way water waves travel away from a stone
dropped in a pond.
This “Wave Model of Light”
Survived into the 1900’s!
In the early 1900’s scientists were still
using cathode ray tubes to study light !
When they passed electricity through
gases, the electrons in the gas atoms
would absorb the extra energy.
The atom is then said to be “excited”!
However, the electrons don’t keep this
extra energy for long.
They immediately give it back off in the
form of Electromagnetic Radiation –
Energy that travels through space as
waves.
Light (E-M Radiation)
• All types travel at light speed (c)
• 3.00x108 m/s
• All types have wave characteristics
(wavelength, frequency)
One cycle
(Frequency is # of cycles per second)
wavelength (l -lambda) - distance between
successive peaks (m)
Frequency (n - nu) - # cycles passing a
given point each second (1/s or Hz)
9
Electromagnetic Radiation covers a broad spectrum:
Types of Electromagnetic Radiation
(only long wave on list!)
(decreasing l)
Radio waves ~ 103m
Microwaves ~ 10-3m
Infrared light ~ 10-5m
Visible light
~ 10-6m
Ultraviolet light ~ 10-8m
X-rays
~ 10-10m
Gamma rays ~ 10-12m
red
orange
yellow
green
blue
indigo
Violet
750 nm
400 nm
Link to FCC Radio Frequency Chart
Because all EM radiation moves
at the same speed, wavelength (l)
and frequency () are inversely
proportional:
Speed of light!
c=l
What is the wavelength of radiation
whose frequency is 6.24 x l014 sec-1?
c = l 3 x 108 m/s
-7 m
=
4.81
x
10

6.24 x 1014 s
Is this visible light? If so, what color?
4.81 x 10-7m x 109nm = 481 nm YES! Blue
1m
2.
what is the frequency of radiation whose
wavelength is 2.20 x l0-6 nm? (1 m = 109 nm)
2.20 x l0-6 nm x
•
c
=
l
1m
109nm
3 x 108 m/s
2.20 x l0-15 m
= 2.20 x 10-15m
= 1.36 x 1023 s-1
Is this visible light? If so, what color?
No! Gamma or cosmic radiation
Remember - When heat or electricity is
passed through a gas, the electrons in
the gas atoms absorb the extra energy.
The atom is then said to be “excited”!
But, the electrons don’t keep this extra
energy for long.
They immediately give it back off in the
form of Electromagnetic Radiation
(visible light)
One way to demonstrate the emission of
light from excited atoms is by using a
Flame Test.
Flame Tests
You heat a metallic salt & it burns with a
colored flame!
 This is the “characteristic glow” of the
excited metal ions!

Fireworks
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Flame Emission Spectra
methane gas
wooden splint
sodium ion
calcium ion
copper ion
strontium ion
“Neon” signs
Bent up cathode ray tube!
NOT!!!
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.

Bright-Line Spectra


Passing the light
from excited atoms
through a prism
does something
different The spectrum
contains lines of
only a few colors
or wavelengths.
Bright-Line Emission Spectrum
excited state
Wavelength (nm)
410 nm
486 nm
434 nm
Slits
ENERGY IN
PHOTON OUT
ground state
Prism
656 nm

Each element has a unique bright-line
spectrum.
i.e. an element’s “fingerprint”
Helium
This is how we know what stars are made of!
Spectrum of White
Light
Spectrum of
Excited Hydrogen Gas
Emission Spectrum
of Hydrogen
1 nm = 1 x 10-9 m = “a billionth of a meter”
410 nm 434 nm
486 nm
1 nm = 1 x 10-9 m = “a billionth of a meter”
656 nm
Continuous and Line Spectra
Visible
spectrum
light
l (nm)
400
450
500
550
600
650
700
Na
H
Ca
Hg
o
4000 A
5000
6000
7000
750 nm
At the beginning of the 20th century
the accepted theory of light was still
the wave model.
(light & other forms of electromagnetic
radiation travel as waves)
Scientists found that only a certain
minimum energy could excite atoms &
get them to emit light.
So they knew that energy had to be
related to the fundamental properties
of frequency & wavelength
The temperature of a Pahoehoe lava
flow can be estimated by observing its
color. The result agrees well with the
measured temperatures of lava flows at
about 1,000 to 1,200 °C.
1900, Max Planck (Germany)
accurately predicted how the
spectrum of radiation emitted by an
object changes with its temperature.
Max Planck
The color (wavelength) of light
depends on the temperature –
“Red hot” objects are cooler
than “white hot” objects
Planck suggested that the energy
absorbed or emitted by an object is
restricted to ‘pieces’ of particular size.
•He named each small “chunk”
of energy a quantum
(meaning “fixed amount)
A quantum is the smallest unit of energy
Although small, quanta are significant
amounts of energy on the atomic level.
Planck
said that the energy of a
light is directly proportional to
its frequency
c = ln so n = c/l
E = hn
E:
h:
n:
E=hc
l
energy (J, joules)
Planck’s constant (6.6262  10-34 J·s)
frequency (Hz)
really small!
(so inversely proportional to wavelength!)
Small wavelength
Large frequency
Large energy
Large wavelength
Small frequency
Small energy
Quantum Theory

Example: Find the energy of a red photon
with a frequency of 4.57  1014 1/s.
GIVEN:
E=?
n = 4.57  1014 1/s
h = 6.6262  10-34 J· s
WORK:
E = hn
E = (6.6262  10-34 J· s) (4.57  1014 1/s)
E = 3.03  10-19 J
Examples:
1. If a certain light has 7.18 x l0-19 J of energy,
what is the frequency of this light?
A: 1.08X1015 s-1or Hz
b. what is the wavelength of this light?
A: 2.78X10-7 m
2. If the frequency of a certain light is 3.8 x l014 Hz,
what is the energy of this light?
A: 2.5X10-19 J
3. The energy of a certain light is 3.9 x l0-19 J.
What is the wavelength of this light? Is it
visible?
A: 510 nm – Yes visible light.
What if the energy of a car was ‘quantized’?
The car would only be able to move at
certain speeds!
Let’s say a car’s fundamental quantum of
energy was equal to 10 mph.
If it had 7 quanta, how fast would it be
going?
Yeppers! 70 mph
If it had 3 quanta?
30 mph
The car can gain or lose energy only in
multiples of its fundamental quantum – 10
No gradual acceleration or deceleration!
It couldn’t go 25 mph or 67 mph…
So, why aren’t we aware of quantum
effects in the world around us?
Remember the size of Planck’s Constant?
It is very small (10-34)
To us, energy seem continuous because
the quanta are too small to be noticed.
However, for atoms, which are also very
small, quanta are of tremendous
significance!
Albert Einstein saw Planck’s idea
of quantized energy as a new way
to think about light.
In 1905, Einstein used Planck’s
equation to explain another
puzzling phenomenon –
The Photoelectric Effect.
The Photoelectric Effect – refers to the
emission of electrons from a metal when
light shines on the metal.
The wave theory of light (early 1900) could not
explain this phenomenon. For a given metal, no
electrons were emitted if the light’s frequency was
below a certain minimum – regardless of how long
the light was shone. Light was known to be a form
of energy, capable of knocking loose an electron
from a metal. But the wave theory of light predicted
that light of any frequency could supply enough
energy to eject an electron. Scientists couldn’t
explain why the light had to be of a minimum
FREQUENCY in order for the photoelectric effect to
occur.
Solar Calculator
Solar Panel
Albert Einstein expanded on Planck’s
theory by explaining that electromagnetic
radiation has a dual wave-particle nature.
While light exhibits many wavelike
properties, it can also be thought of as a
stream of particles.
Each particle of light carries a
quantum of energy directly
proportional to the frequency.
Einstein called these particles photons.
A PHOTON is a packet of light
carrying a quantum of energy.
Photon
• Is light a wave or a particle?
• Macroscopically it behaves as a wave!
• On the atomic level, we observe
particle properties!
Seen on the door to a light-wave lab:
"Do not look into laser with remaining good eye."
Dual Nature of Light – light exhibits
wave properties & particle properties
Einstein explained the photoelectric effect by
proposing that electromagnetic radiation is
absorbed by matter only in whole numbers of
photons.
In order for an electron to be ejected from a metal
surface, the electron must be struck by a single
photon possessing at least the minimum energy
(Ephoton = hv) required to knock the electron loose,
this minimum energy corresponds to a minimum
frequency. If a photon’s frequency is below the
minimum, then the electron remains bound to the
metal surface. Electrons in different metals are
bound more or less tightly, so different metals
require different minimum frequencies to exhibit
the photoelectric effect.
Photoelectric Effect
Electrons are emitted
No electrons are emitted
Bright
red light
or
Dim
blue light
or
infrared rays
ultraviolet rays
Metal plate
Metal plate
Quantized Energy and Photons
Phenomena not explained by wave nature
of light:
• 1) Black-body radiation – light coming from a
heated object (Planck)
2) Photoelectric effect – electrons emitted from
light illuminated surface (Einstein)
emission spectrum (top),
absorption spectrum (bottom)
• 3) Emission Spectra – light from electronically
excited gas atoms
50
Neils Bohr
1913
studied under Rutherford at
Victoria University in Manchester.
Niels Bohr
Bohr refined Rutherford's idea
by adding that the electrons
were in orbits. Rather like
planets orbiting the sun. With
each orbit only able to contain a
set number of electrons.
Bohr’s Model of Hydrogen (1913)
Niels Bohr
(1885-1962)
Neils Bohr incorporated Planck’s
quantum theory to explain bright-line
spectra.
Bohr said the absorptions and
emissions of light by hydrogen
corresponded to energy changes
within the atom.
The fact that only certain
frequencies are absorbed or
emitted by an atom
tells us that only certain energy
changes are possible in an atom.
Bohr’s Planetary Model of the Atom
•
•
•
electrons exist only in orbits with
specific amounts of energy called
energy levels
Therefore…
electrons can only gain or lose certain
amounts of energy
(quanta)
The orbit closest to
The nucleus is the
most stable & lowest
In energy.
Bohr’s
Planetary
Model
of the
Atom
Nucleus
Electron
Orbit
Energy Levels
Niels Bohr &Albert Einstein
The lowest energy state of an atom is
its ground state.
A state in which an atom has a higher
amount of energy is an excited
state.
When an excited atom returns to its
ground state, it gives off photons
of energy (light!)
Electrons can’t stop
between energy
levels so the ‘jumps’
involve definite
amounts of
energy.
(amount of energy
~ light color!)
Electrons can only be at
specific energy levels,
NOT between levels.
Electrons can ‘jump’
to a higher energy
level when the atom
absorbs energy.
Excited state
When the electron
drops back down to
a lower level, it gives
the extra energy off
as light.
e-
Ground state
An excited lithium atom emitting a
photon of red light as it drops to a
lower energy state.
Energy
Excited Li atom
Photon of
red light
emitted
Li atom in
lower energy state
Electron Energy Levels
3rd energy level
2nd energy level
Energy
absorbed
1st energy level
Energy
lost
nucleus
Bohr Model of Atom
Increasing energy
of orbits
n=3
e-
n=2
e-
n=1
ee-
e-
e-
ee-
e-
e-
eA photon is emitted
Bohr Model
6
5
4
 Energy
3
2
1
nucleus
of photon
depends on the
difference in energy
levels
 Bohr’s
calculated
energies matched the
bright-line spectrum for
the H atom
Bohr Model Limitations
Unfortunately, this model only works for
Hydrogen!
The success of Bohr’s model of the hydrogen
atom is explaining observed spectral lines led
many scientist to conclude that a similar
model could be applied to all atoms. It was
soon recognized, however, that Bohr’s
approach did not explain the spectra of atoms
with more than one electron. Nor did Bohr’s
theory explain the chemical behavior of
atoms.
62
MORE SCIENTIFIC ADVANCEMENTS!
With more sophisticated equipment, spectral
lines were found to consist of closely spaced
lines called ‘Doublets”
Hydrogen (pretty simple!)
Helium (not so basic!)
doublets
So there had to be more to Bohr’s energy
levels (orbits) than he realized.
In 1924, Louis DeBroglie suggested that
every moving particle has a wave nature
just like light!

De Broglie’s Hypothesis Duality of Matter

Louis de
Broglie
~1924
Since waves have particle characteristics
(Dual Nature of Light)
 Moving
particles have wave
characteristics
According to Isaac Newton, we can determine
both the position & momentum of a large body.
(like an airplane)
However, we CANNOT accurately predict
where an electron will be at some future time!
Heisenberg Uncertainty Principle (1926)
says that it is impossible to know both the
location and the momentum of an electron
simultaneously.
Werner Heisenberg
1901-1976
Heisenberg Uncertainty
Principle
“One cannot simultaneously
determine both the position
and momentum of an electron.”
You can find out where the
electron is, but not where it
is going.
Werner
Heisenberg
OR…
You can find out where the
electron is going, but not
where it is!
67
Heisenberg Uncertainty Principle
In order to observe an electron, one would need to
hit it with photons having a very short wavelength.
Short wavelength photons would have a high
frequency and a great deal of energy.
If one were to hit an electron, it would cause the
motion and the speed of the electron to change.
g
Microscope
Electron
heck
In the Bohr Model of the atom, the
electron is at a fixed distance from
the nucleus.
He assumed we knew both the position
& the momentum of electrons.
The Uncertainty Principle disproves
this!
A New Model! (The Last One!)
So… De Broglie and Heisenberg’s contributions
lead us to a new atomic model.
It will recognize the wave nature of the electron
and describe it in terms appropriate to waves.
The resulting model will precisely describe the
ENERGY of the electron, while describing its
location as a probability.
72
The Quantum Mechanical
Model of the Atom
Erwin Schrodinger
1887-1961
Erwin Schrodinger
Schrodinger applied DeBroglie’s idea of electrons
behaving as waves to the problem of electrons in
atoms.
Schrödinger’s Quantum Mechanical Model

Used to determine the PROBABILITY of
finding an electron at any given distance
from the nucleus

Describes the electron as a 3-dimensional
wave surrounding the nucleus.

(fan blades)
Quantum Mechanical Model
1926
The Quantum Mechanical Model of
the atom describes the electronic
structure of the atom as the probability
of finding electrons within certain
regions of space (orbitals).
Today we say that the electrons
are located in a region of space
outside the nucleus called the
electron cloud.
Quantum Mechanics
Orbital - densist, darkest region of the
“electron cloud”)

Region in space where there is a high
(90%) probability of finding an electron
90% probability of
finding the electron
Electron Probability vs. Distance
Electron Probability (%)

40
30
20
10
0
0
50
100
150
200
250
Distance from the Nucleus (pm)
Electron cloud
Erwin Schrodinger (1887-1961)
• Won Nobel Prize in 1933 for his equation.
• Came up with a paradoxical thought
experiment to show problems in observing
isolated systems (Schrodinger’s Cat)
Experiment: A cat is placed in a sealed box
containing a device that has a 50% chance of
killing the cat.
LOHS AP Chemistry Fall 2007
Dr. Schrempp
78
Unfortunately, Schrodinger’s cat could not cope
with a life of uncertainty…
79
Modern View


The atom is mostly empty space
Two regions


Nucleus
 protons and neutrons
Electron cloud
 region where you are likely to
find an electron
“I don't like it, and I'm sorry I ever had anything to
do with it.”
- Erwin Schrodinger talking about Quantum Physics
Feeling overwhelmed?
Just a little
more!
"Teacher, may I be excused? My brain is full."
4
Quantum Numbers
The QM model makes it possible to
describe the location of an electron
Using four quantum numbers.
4
4

4
Four Quantum Numbers:
4
 Specify
4
the “address” of each
electron in an atom
4
4
4
4
Quantum Numbers
Principal Quantum Number ( n )
Angular Momentum Quantum # ( l )
Magnetic Quantum Number ( m )
Spin Quantum Number ( s )
Quantum Numbers
1. Principal Quantum Number ( n )

Energy level

Size of the orbital cloud

n = 1, 2, 3, 4…

n2 = # of orbitals in
the energy level

1s
2n2 = # of electrons per energy level
2s
3s
Quantum Numbers
2. Angular Momentum Quantum # ( l )
Corresponds to:
Energy sublevel
Shape of the orbital
l = s, p, d, f
(in order of increasing energy)
s cloud is spherical,
p cloud is dumb-bell shaped
s
p
d
f
Sublevel ‘names’
Quantum Numbers
3. Magnetic Quantum Number ( m )


Orientation (direction in space) of orbital
Specifies the exact orbital within each
sublevel that the electron occupies.
Quantum Numbers
A “p” sublevel has 3 possible orbitals
oriented along the “x”, “y”, & “z” axes.
y
y
z
x
z
x
px
y
z
x
pz
py
px
py
pz
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
d-orbitals
A “d” sublevel has 5 possible orbital clouds
f Orbitals
An f sublevel has 7
possible orbitals
Quantum Numbers
4. Spin Quantum Number ( s )


Electron spin  +½ or -½
An orbital can hold 2 electrons that spin in
opposite directions – clockwise or counterclockwise.
+½
-½
A Cross Section of an Atom
# of sublevels per energy level = n
Rings of Saturn
n0
p+
1s
2s
2p
3s
3p
3d
The first energy level has only one sublevel (1s).
The second energy level has two sublevels (2s and 2p).
The third energy level has three sublevels (3s, 3p, & 3d).
*Although the diagram suggests that electrons travel in circular orbits,
this is a simplification and is not actually the case.
Quantum Numbers
Principal
level
n=1
Sublevel
s
Orbital
n=2
s
p
px
py pz
n=3
s
p
px
py pz
d
dxy
dxz
dyz
dz2
dx2- y2
Maximum Electron Capacities of Subshells
and Principal Shells
n
1
l
0
0
1
0
1
2
0
1
2
3
Sublevel
designation
s
s
p
s
p
d
s
p
d
f
Orbitals in
sublevel
1
1
3
1
3
5
1
3
5
7
Sublevel
capacity
2
2
6
2
6
10
2
6
10 14
Principal energy
Level capacity
2
2
8
3
18
4
...n
32
...2n 2
Quantum Numbers
 Each
electron has a unique
“address”:
Electron spin
+1/2
2px
Energy level
orbital
sublevel
1. Principal #
2. Ang. Mom. #
3. Magnetic #
4. Spin #
n
l
m
s




energy level
sublevel (s,p,d,f)
orbital
electron
ATOMIC STRUCTURE
There are 3 ways to represent the electron
arrangement of an atom:
1.
Electronic Configuration
2.
Orbital Filling Diagrams
3.
Lewis Dot Diagrams
3 ways to represent electron
arrangements in atoms:
Orbital Notation (orbital filling diagrams):
 An orbital is represented by a line or box.



The lines are labeled with the principal
quantum number and the sublevel letter.
Arrows represent the electrons.
An orbital containing one electron is
written as  , an orbital with two
electrons is written as .
Electron Configurations
A list of all the electrons in an atom (or ion)
4
2p
Number of
electrons in
the sublevel
Energy Level
Sublevel
1s22s22p4
Filling Rules for Electron Orbitals
 Pauli
Exclusion Principle
 No
two electrons in an atom can
have the same 4 quantum numbers.

Each orbital can only hold TWO electrons
with opposite spins.
Wolfgang Pauli
General Rules
6d
Aufbau Principle
7s
6p
5d

Electrons fill the
lowest energy
orbitals first.
“Lazy Tenant
Rule”
6s
4d
3p
5f
7s
6p
5d
6s
5p
5s
4p
4s
6d
4f
5p
Energy

5f
4d
5s
3d
4p
3d
4s
3p
3s
3s
2p
2p
2s
2s
1s
1s
*Aufbau is German for “building up”
4f
Examples:
(Remember that you must place
one electron into each orbital of a sublevel before a
second electron in placed into an orbital.)
1s1
Hydrogen
1s
1s22s2
1s2
Helium
1s
Lithium
1s
Carbon
Boron
2s
1s
2s
1s22s22p2
1s
Be
1s22s1
2px
1s
2py
2s
2pz
2px
2py
2pz
1s22s22p1
2s
1s22s22p2
Carbon
1s
2s
2px
2py
2pz
Hund’s Rule
Within a sublevel, place one electron per
orbital before pairing them.
“Empty Bus Seat Rule”
Hund’s Rule


Also called The “Monopoly Rule”
In Monopoly, you have to build
houses EVENLY. You can not
put 2 houses on a property until
all the properties has at least 1
house.
8
O
Orbital Notation

15.9994
Orbital Diagram
O
8e
1s
2s
Electron Configuration
2
2
4
1s 2s 2p
2p
Orbital Notations
Orbital Filling
Element
1s
2s
2px 2py 2pz
3s
Electron
Configuration
H
1s1
He
1s2
Li
1s22s1
C
1s22s22p2
N
1s22s22p3
O
1s22s22p4
F
1s22s22p5
Ne
1s22s22p6
Na
1s22s22p63s1
Write the electron config. for K
4f
Sc ?
4d
Energy
n=4
Filling order becomes
irregular after Ar
Because of overlapping
of electron clouds
(orbitals) in larger
atoms
n=3
4p
3d
4s
3p
3s
2p
n=2
2s
n=1
1s
Order in which orbitals are filled with
electrons
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p66s24f145d106p67s2
Diagonal Rule
Must be able to write it for the test!
Without it, you may not get correct
answers !
 The diagonal rule is a memory
device that helps you remember the
order of the filling of the orbitals from
lowest energy to highest energy

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p66s24f145d106p67s2
1s2
2s2
2p6
3s2
3p6
3d10
4s2
4p6
4d10
4f14
5s2
5p6
5d10
5f14
6s2
6p6
6d10
6f14
7s2
7p6
7d10
7f14
Electron Configurations Ws. #3
Symbol
# e-
Orbital Diagram and
Longhand Electron Configuration
Mg
12

1s

2s
  
2p

3s
P
15

1s

2s
  
2p

3s
  
3p
V
23

1s

2s
  
2p

3s
  
3p


1s

2s
  
2p

3s
  
3p
     
3d
4s
  ___
Ge
32
 
3d

4s
4p
Kr
36

1s

2s
  
2p

3s
  
3p
     
3d
4s
  
4p
O
8

1s

2s
  
2p
Part B – Rules of Electron Configurations
Which of the following “rules” is being violated in each electron configuration below?
Explain your answer for each. Hund’s Rule, Pauli Exclusion Principle, Aufbau Principle

1s

2s
 __ __ Hund’s Rule – should be 1 electron in each of
2p
the 1st 2 orbitals (not doubled up)

1s

2s
   ___
2p
3s

1s

2s
      _ Pauli Exclusion Principle
2p
3s
3p – 1 electron in 3s needs to
 _ _ Aufbau Principle – need
3p
to fill 3s before 3p
point down (opposite spins)

1s

2s
           
2p
3s
3p
3d
Aufbau – must fill 4s before 3d
Write out the complete electron configuration for the following:
1) An atom of nitrogen
1s22s22p3
2) An atom of silver 1s22s22p63s23p64s23d104p65s24d9
Fill in the orbital boxes for an atom of nickel (Ni)
1s
2s
2p
3s
3p
3d
4s
Shorthand Notation
A
way of abbreviating long
electron configurations
 Since we are only concerned
about the outermost electrons,
we can skip to places we know
are completely full (noble
gases), and then finish the
configuration
Shorthand Notation
 Step
1: It’s the Showcase
Showdown!
Find the closest noble gas to the
atom (or ion), WITHOUT GOING
OVER the number of electrons
in the atom (or ion). Write the
noble gas in brackets [ ].
 Step 2: Find where to resume by
finding the next energy level.
 Step 3: Resume the
configuration until it’s finished.
Shorthand Configuration for Na
A
B
neon's electron configuration (1s22s22p6)
third energy level
[Ne] 3s1
C
D
one electron in the s orbital
orbital shape
Na = [1s22s22p6] 3s1
electron configuration
Part B – Shorthand Electron Configuration
Use the Noble gas that comes before an element on the periodic table to
represent all inner electrons. Put the symbol of the Noble gas in
parentheses
Sy
Shorthand Electron
#e
mb
Configuration
ol
Ca 20
[Ar] 4s2
Pb 82
[Xe] 6s24f145d106p2
F
9
[He] 2s22p5
U
92
[Rn] 7s25f4
Shorthand Notation
Chlorine
 Longhand is 1s2 2s2 2p6 3s2 3p5
You can abbreviate the first 10
electrons with a noble gas,
Neon. [Ne] replaces 1s2 2s2 2p6
The next energy level after Neon
is 3
So you start at level 3 on the
diagonal rule (all levels start with
s) and finish the configuration by
adding 7 more electrons to bring
the total to 17
[Ne] 3s2 3p5
Boron is 1s22s22p1
The noble gas preceding Boron is He, so the
short way is [He]2s22p1.
Sulfur is 1s22s22p63s23p4
Short way: [Ne]3s23p4
Example: Titanium
[Ar]4s23d2
Practice Shorthand Notation
 Write
the shorthand notation for
each of the following atoms:
K
Ca
I
Bi
Shorthand Configuration
Element symbol
Electron configuration
Ca
[Ar] 4s2
V
[Ar] 4s2 3d3
F
[He] 2s2 2p5
Ag
[Kr] 5s2 4d9
I
[Kr] 5s2 4d10 5p5
Xe
[Kr] 5s2 4d10 5p6
Fe
Sg
[Ar] 4s23d6
[Rn] 7s2 5f14 6d4
Valence Electrons
Electrons are divided between core and
valence electrons –
electrons in the highest energy level
of an atom.
These are the electrons that take
part in reactions (so most important!)
Never more than 8 valence electrons
in an atom
(2 s & 6 p)
Lewis (Electron) Dot Structures
Shorter than configs. & only show
valence electrons
Element symbol surrounded by # dots =
to number of valence electrons.
No more than 2 electrons per side on
symbol.
Start at top of symbol, add dots
clockwise, one per side before
doubling them up.
No. of valence electrons = Group number
(for A groups)
B 1s2 2s2 2p1
B
Core = [He] , valence = 2s2 2p1
B is in Group 3A, so 3 valence electrons (dots)
Br [Ar] 4s2 3d10 4p5
Core = [Ar] 3d10
valence = 4s2 4p5
Br is in 7A, so 7 valence electrons
Br
It is very important to define “stable”
here.
STABLE means: (in order of stability)
1. all equal energy orbital’s are FULL
2.
all orbital’s are half-full
3.
all orbital’s are totally empty.
Some More Stuff!!
1. The highest energy electron is the LAST one
you write in the electron configuration.
1s22s22p63s23p5 -- the 3p5 electron is the
last written. *Remember Aufbau’s Principle,
electrons fill from the lowest to the highest
energy.
2.
The outermost electron is the one with the
LARGEST principle quantum number.
1s22s22p63s23p64s23d104p2. The 4 p2 is the
farthest from the nucleus. OR
(2) 1s22s22p63s23p64s23d10. Here, it is the 4s2
electron, because it has the largest principle
q.n.

Irregular Electron configurations –
sometimes the electron configuration is
NOT what we would predict it to be.
Sometimes electrons are moved
because (1) it will result in greater
stability for that atom or (2) for some
unknown reason??
Examples –
Predict the electron configuration for Cr
#24:
[Ar]4s23d4
However, the real E. C. is [Ar]4s13d5. The
4s2 electron has been moved to 3d to
achieve greater stability.
Exceptions, Con’t.
Cr: [Ar] 4s13d5 NOT Cr: [Ar] 4s23d4
Because lower energy results from halffilling 6 orbitals with spins aligned instead
of causing repulsion in one of the 3d
orbitals
Cu: [Ar] 4s13d10 NOT [Ar] 4s23d9
Because easier to add the electron to a
sublevel with four electrons that already
have the same spin than causing repulsion
in a different orbital
132
Electron Orbitals:
Electron
orbitals
Equivalent
electron
shells (Bohr)
Neon Ne-10:
1s, 2s and 2p
Relative Sizes 1s and 2s
1s
2s
s, p, and d-orbitals
A
s orbitals:
Hold 2 electrons
(outer orbitals of
Groups 1 and 2)
B
p orbitals:
Each of 3 pairs of
lobes holds 2 electrons
= 6 electrons
(outer orbitals of
Groups 13 to 18)
C
d orbitals:
Each of 5 sets of
lobes holds 2 electrons
= 10 electrons
(found in elements
with atomic no. of 21
and higher)
Principal Energy Levels 1 and 2
Arbitrary Energy Scale
Aufbau Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
CLICK ON ELEMENT TO FILL IN CHARTS
Fe La
Arbitrary Energy Scale
Aufbau Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Hydrogen
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
CLICK ON ELEMENT TO FILL IN CHARTS
Fe La
H = 1s1
Arbitrary Energy Scale
Aufbau Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Helium
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
CLICK ON ELEMENT TO FILL IN CHARTS
Fe La
He = 1s2
Arbitrary Energy Scale
Aufbau Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Lithium
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
CLICK ON ELEMENT TO FILL IN CHARTS
Fe La
Li = 1s22s1
Arbitrary Energy Scale
Aufbau Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Carbon
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
CLICK ON ELEMENT TO FILL IN CHARTS
Fe La
C = 1s22s22p2
Arbitrary Energy Scale
Aufbau Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Nitrogen
4f
Bohr Model
N
Hund’s Rule “maximum
number of unpaired
orbitals”.
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
CLICK ON ELEMENT TO FILL IN CHARTS
Fe La
N = 1s22s22p3
Arbitrary Energy Scale
Aufbau Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Fluorine
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
CLICK ON ELEMENT TO FILL IN CHARTS
Fe La
F = 1s22s22p5
Arbitrary Energy Scale
Aufbau Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Aluminum
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
CLICK ON ELEMENT TO FILL IN CHARTS
Fe La
Al = 1s22s22p63s23p1
Arbitrary Energy Scale
Aufbau Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Argon
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
CLICK ON ELEMENT TO FILL IN CHARTS
Fe La
Ar = 1s22s22p63s23p6
Arbitrary Energy Scale
Aufbau Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Iron
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
Fe = 1s22s22p63s23p64s23d6
NUCLEUS
H He Li C N Al Ar F
CLICK ON ELEMENT TO FILL IN CHARTS
Fe La
Arbitrary Energy Scale
Aufbau Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
4f
Lanthanum
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
CLICK ON ELEMENT TO FILL IN CHARTS
La = 1s22s22p63s23p64s23d10
Fe La 4s23d104p65s24d105p66s25d1
Electron
capacities
Copyright © 2006 Pearson Benjamin Cummings. All rights reserved.
So, where are the electrons of an atom
located?
Various Models of the Atom
Dalton’s Model
Thompson’s Plum Pudding Model
Rutherford’s Model
Bohr’s ‘Planetary’ Model – electrons rotate around the
nucleus
Quantum Mechanics Model – modern description of the
electron in atoms, derived from a mathematical equation
(Schrodinger’s wave equation)
Diagonal Rule
Steps:
1s
2s
3s
1.
Write the energy levels top to bottom.
2.
Write the orbitals in s, p, d, f order. Write
the same number of orbitals as the
energy level.
3.
Draw diagonal lines from the top right to
the bottom left.
4.
To get the correct order,
2p
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
6f
7s
7p
7d
7f
follow the arrows!
By this point, we are past
the current periodic table
so we can stop.
Atoms like to either empty or fill their outermost
level. Since the outer level contains two s
electrons and six p electrons (d & f are always in
lower levels), the optimum number of electrons
is eight. This is called the octet rule.