Electron Configurations
Chapter 4
Chemistry-CP
Radiant Energy
• The understanding of how electrons behave
comes from studies of how light interacts
with matter.
• Light carries energy through space in the
form of waves and also in the form of
extremely tiny, fast moving particles.
– Light has the properties of waves & particles.
Light as Waves
• Light waves are a form of electromagnetic
radiation
Waves
• Light travels at the speed of light.
– The speed of light is constant, which means it is always
the same value: 3.00 × 108 m/s
• Because light moves at a constant speed, wavelength &
frequency are inversely proportional as per the following
equation.
c=×
c = speed of light
 = wavelength (lambda)
 = frequency (measured in 1/s or s-1 or Hertz (Hz))
c=×
• What is the frequency of a wave having a wavelength of
8.12 x 102 m?
• A helium neon laser produces red light whose wavelength
is 633 nm. What is the frequency of this radiation?
• Calculate the wavelength of a radio wave with a frequency
of 9.31 × 106 s-1.
Gamma Rays
• Generated by radioactive atoms, nuclear
explosions and supernova explosions
• Can kill living cells—used for cancer
treatment
• Used to sterilize medical equipment
•
http://www.youtube.com/watch?v=NZF3_e6_xj4
X-Rays
• Discovered by accident in 1895, when W.C. Roentgen
shielded a cathode ray tube with black paper and found
that a fluorescent light could be seen on a screen a few
feet from the tube (first bone x-ray was of his wife’s
hand!)
• Electrons shot at an element (such as tungsten or
molybdenum) with high energy can knock an electron
out of that atom, producing x-rays
• Used for radiography, crystallography, astronomy, airport
security
Ultraviolet Radiation
• Gets its name from the fact that it consists of
waves with frequencies higher than what
humans associate with violet light
• Emitted from the sun, from black lights
• UV-B produces Vitamin D, too much = DNA
damage & collagen fibers, can cause sunburn,
may lead to cataracts
• Some animals, insects, birds and reptiles can
see the near ultraviolet making certain flowers,
etc. brigher to them.
• Portion of the
electromagnetic
spectrum that is
visible to the human
eye
• ROYGBIV—Violet
has the highest
frequency
Infrared
• “Below red”
• “Heat radiation”
• Emitted from humans at normal body
temperature
• Military use (surveillance, night vision, homing)
• Short ranged wireless communication, weather
forecasting, remote temperature sensing
– Purple white light get on cheaper digital cameras
(poor infrared filters)
Microwaves
•
•
•
•
Wireless LAN & Bluetooth
Radar Detectors, Air Traffic Control
GPS
The frequency of the waves used in microwave
ovens, 2500 megahertz, targets water, sugar &
fat molecules
– Thin, sharp metals can not handle the electric current
passing through them and may spark or start a fire
• Has never been conclusively shown that
microwaves have biological effects
• http://www.youtube.com/watch?v=Ug8hSqkFUXY
•
http://www.youtube.com/watch?v=PIrd4172Czw
Radio Waves
• Transport information
through the
atmosphere or space
without wires
• AM & FM Radio, TV
transmission, mobile
phones, military
communications,
wireless computer
networks
Visible Spectrum
Part of the electromagnetic spectrum
•Continuous Spectrum: One color
fades gradually into the next.
•Different colors have different
wavelengths.
•The color of visible light with the
largest wavelength and lowest
frequency is:
•The color of visible light with the
shortest wavelength and highest
frequency is:
•The brightness of visible light is
determined by:
• Radiation with the largest
wavelength and lowest frequency is:
Aircraft & Shipping Bands, Radio Waves
• Radiation with the shortest
wavelength and highest frequency is:
Gamma Rays
• Radiation with frequencies greater
than visible light can pose health
hazards because:
Have high enough energy to be capable of
damaging organisms
• Radiation with frequencies lower
than visible light are less harmful
because:
Do not have enough energy and pose no
health hazards
What puzzled scientists about
electromagnetic radiation?
• Why do objects at different temperatures
give off different color light?
• Why do different elements emit different
colors when heated?
Planck’s Theory
• Suggested that the energy emitted or
absorbed by an object is restricted to
“pieces” of particular sizes called quanta.
– Substances can emit or absorb only certain
amounts of energy (so only certain
wavelengths)
– Showed that frequency and energy are
directly proportional
Planck’s
Planck’sTheory
Theory
• E=h×
– h = Planck’s constant = 6.626 × 10-34 J●s
• Joule (J) = S.I. Unit for Energy
What
the
Ultraviolet
Light?
What
ismuch
the approximate
frequency
wave of
with
an aenergy
ofwith
2.90ax
Howis
energy of
is aenergy
contained
in
wave
22 J?
10
frequency of 2 x 1016 Hz?
What is the energy of radiation with a wavelength of 290
nm?
Why can’t you see quantum effects
in the everyday world around you?
• Planck’s constant is very small, therefore,
each quantum of energy is very small
– Quanta are too small to see in the everyday
world
– Atoms are very small, so in relation to the
atom, quanta are significant
The Photoelectric Effect
• Proposed by Albert
Einstein
– In the photoelectric effect,
electrons are ejected
from the surface of a
metal when the metal
absorbs photons
– Photon: Quantum of light
(a tiny particle of light)
The Photoelectric Effect
• When a photon strikes
the surface of a metal, it
transfers its energy to an
electron in a metal atom
– If the energy of the photon
is too small for the
electron, the electron stays
put
– If the photon has enough
energy, the electron will
escape the surface of the
metal.
The Photoelectric Effect
• Why is it easier for violet light (vs. red light) to
cause the photoelectric effect?
Violet light has a higher frequency and, therefore,
more energy than red light.
Light has a Dual Nature
• A photon behaves like a
particle but always travels
at the speed of light and
has an associated
frequency and wavelength
– In 1923, Arthur Compton
showed that a photon could
collide with an electron
– Light possesses the
properties of both particles
and waves
How can atoms gain or lose energy?
• Atoms can only gain or lose energy in a
quantum
• Take a look through your spectral tube at
the emission tube at the front of the room.
– How does what you’re looking at demonstrate
the idea above?
Line Spectrum
• A spectrum that contains only certain
colors, or wavelengths
•When heat or electricity is passed through
an atom, the atom absorbs the energy and
then gives off that energy in the form of light
•The emitted light is unique for every
element
•Atomic Emission Spectrum: An atomic
fingerprint showing the emission line
spectrum of that atom
•Useful in identifying an element
NIELS BOHR—1913
PLANETARY MODEL OF THE ATOM
Electrons move
in defined
orbits around
the nucleus—
just as the
planets move
around the sun.
Bohr’s Postulate
• Was applicable only to
hydrogen
• Able to show that
electrons move to higher
energy levels (excited
states) when they absorb
radiation.
• Electrons will immediately
return to the lower energy
levels (ground state) by
emitting energy of a
specific wavelength
Light has a Dual Nature
– When light travels through space,
it acts like a wave
– When light interacts with matter, it
acts like a particle
– De Broglie predicted matter
waves--that matter should
behave like waves and exhibit a
wavelength
– Clinton Davisson & Lester
Germer proved that electrons
(believed to be particles) were
reflected from a matter like waves
• Mass of an object must be very
small in order to observe its
wavelength
Heisenberg Uncertainty Principle
• An electron is located by
striking that electron with a
photon which bounces back to
a detection device
• The electron is so small in
mass that the electron is
moved by the collision
• Proved a problem with Bohr’s
model: You cannot think of
electrons moving in defined
paths because there is no way
to prove the electrons follow
defined paths
MODELS OVER TIME
Quantum Mechanical Model
a.k.a: Wave Model
•Explains the
properties of atoms by
treating the electron as
a wave that has
quantized its energy
•Does not describe
exact positions of the
electrons; instead
describes the
probability that
electrons will be found
in certain locations
around the nucleus
Electron Cloud
An illustration that uses a blurry cloud to illustrate the
probability of finding an electron in various locations around
the nucleus.
(Determined by wave functions electron density charts)
Atomic Orbitals
•Region of space
where the electron is
located
•Have characteristic
shapes, sizes and
energies
•Do not describe how
the electron actually
moves
•The orbital occupied
is determined by the
amount of energy of
an electron
s-Orbital
The s-orbital consists of
1 orbital on all 3 axes
1 orbital has a maximum
of 2 electrons
p-Orbital
The p-orbital can exist on 3
different axes (x, y and z).
Therefore there are 3 p
orbitals.
A p-orbital has a
dumbbell shape
The p-sublevel’s 3 orbitals can
hold a maximum of 6
electrons (2 on each of the 3
orbitals).
d-Orbital
There are 5 different
orientations of a d-orbital.
A d-orbital has a
cloverleaf shape
The d-sublevel’s 5 orbitals can
hold a maximum of 10 electrons
(2 electrons on each orbital).
f-Orbital
An f-orbital has a complex shape
There are 7 different
orientations of the f-orbital.
The f-sublevel can hold a
maximum of 14 electrons (2
for each orbital).
Energy & Orbitals
•Energy of electrons are quantized (exact)
•Principal Energy Levels or Principal Quantum
Number designates the distance of the electron from
the nucleus
Principal energy levels
are divided into sublevels
Sublevels
Sublevels of the atom are
designated:
s, p, d & f
The number of the energy level tells you how many sublevels are
present within that sublevel. Another words:
1
Energy Level 1 has __________
Sublevel
2
Energy Level 2 has __________
Sublevels
Energy Level 3 has __________
Sublevels
3
4
Energy Level 4 has __________
Sublevels
The electrons address consists of its principal energy level, its
sublevel, and its electrons within that sublevel
SUBLEVEL s
Orbital
Shape
Max # of
electrons
Region on
Periodic
Table
1 orbital
Sphere
2
Groups 1
&2
s
(1st tower)
Orbital
Models
SUBLEVEL p
Orbital
Shape
3 orbitals dumbbell
px
py
pz
Max # of
electrons
Region on
Periodic
Table
6
Groups
13-18
(2nd tower)
Orbital
Models
SUBLEVEL d
Orbital
Shape
5 orbitals cloverleaf
dxy
dxz
dyz
dx2-y2
dz2
Max # of
electrons
Region on
Periodic
Table
10
Groups 312
(transition
metals)
Orbital
Models
SUBLEVEL f
Orbital
Shape
7 orbitals complex
Max # of
electrons
Region on
Periodic
Table
14
Bottom 2
rows
(innertransition
metals)
Orbital
Models
Some Atomic Models
More Models
Example
Beryllium: ______ protons, ______ electrons
E- Configuration: 1s22s2
Example
Oxygen: ______ protons, ______ electrons
E- Configuration: 1s22s22p4
PRACTICE PROBLEMS
Electron configurations for 3 different elements are given below.
Draw the atomic model of each element and then identify the element.
Examples: 1s22s1
1s22s22p3
1s22s22p63s23p4
1) 1s22s22p1
2) 1s2
3) 1s22s22p63s1
Example
Boron: ______ protons, ______ electrons
E- Configuration: 1s22s22p1
Examples
Helium: ______ protons, ______ electrons
E- Configuration: 1s2
Examples
Sodium: ______ protons, ______ electrons
E- Configuration: 1s22s22p63s1
Electron Spin
•Electrons spin either clockwise or counterclockwise
•The spinning creates a magnetic field
•Clockwise is like a magnet whose north pole is pointing up
•Counterclockwise behaves like a magnet whose north pole is
pointing down
•Parallel Spins result in a net magnetic effect
•Opposite Spins cancel each other out
Pauli Exclusion Principle
-1925, Austrian physicist-Wolfgang Pauli
-States that each orbital in an atom can hold at most 2
electrons and that these electrons must have
opposite spins (or be paired).
Sublevels
Orbitals
Max # of e-
s
1
2
p
3
6
d
5
10
f
7
14
Electron Configuration
• The addresses of an atom’s electrons
• Determined by distributing the atom’s electrons
among levels, sublevels and orbitals based on a set
of principles
• Orbitals from lowest to highest energy:
s p  d  f
• Ground State: The electrons are in the lowest
energy levels available
How do electrons occupy energy
levels?
• Aufbau Principle: Electrons are added one at a
time to the lowest energy orbitals available until
all the electrons are accounted for
• Pauli Exclusion Principle: An orbital can hold a
maximum of 2 electrons that must spin in opposite
directions
• Hund’s Rule: Electrons occupy equal-energy
orbitals so that a maximum number of unpaired
electrons results
Orbital Diagrams
4p
3d
4s
3p
3s
2p
2s
1s
____
____
____
____
____
____
____
____
____ ____
____ ____ ____ ____
____ ____
____ ____
What happens when an element in its
ground state is supplied with electricity
or heat?
– Electrons may move to the excited state.
– Excited State: Energy level attained when an electron
absorbs energy and jumps to a higher energy level
Ground State
Excited State
For each pair of orbital diagrams below,
which represents the ground state and which
represents the excited state of that atom?
4p
3d
4s
3p
3s
2p
2s
1s
____
____
____
____
____
____
____
____
____ ____
____ ____ ____ ____
4p
3d
4s
3p
3s
2p
2s
1s
____ ____
____ ____
Magnesium
4p
3d
4s
3p
3s
2p
2s
1s
____
____
____
____
____
____
____
____
____
____
____
____
____
____
____
____
____ ____
____ ____ ____ ____
Scandium
____ ____
____ ____
____ ____
____ ____ ____ ____
____ ____
____ ____
4p
3d
4s
3p
3s
2p
2s
1s
____
____
____
____
____
____
____
____
____ ____
____ ____ ____ ____
____ ____
____ ____
What happens to the excited electron?
http://www.meta-synthesis.com/webbook/11_five/five04.jpg
Exceptions to the Aufbau Rule
• A half-full or full d sublevel will increase an
atom’s stability
– An electron may be removed from the s
sublevel to create a full or half full d sublevel
4p
3d
4s
3p
3s
2p
2s
1s
____
____
____
____
____
____
____
____
____ ____
____ ____ ____ ____
____ ____
____ ____
Cr
4p
3d
4s
3p
3s
2p
2s
1s
____
____
____
____
____
____
____
____
____ ____
____ ____ ____ ____
____ ____
____ ____
Cu
Groups
(also called Families)
• The vertical columns on the periodic table
• There are 18 groups, labeled with the numbers 1-18.
1
18
2
13 14 15 16 17
3 4
5
6
7
8 9
10 11 12
Corresponding Regions on the
Periodic Table
He
N
Ti
I
Fr
Ce
Noble Gas Configuration: Uses the symbol of the noble gas
in brackets to represent the inner level electrons of an atom.
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
6s
5d
Ba
6p
7s
6d
7p
Cd
4f
5f
U
5p
VALENCE ELECTRONS
The electrons in the outermost energy level.
Remember, the number in front of the
sublevel indicates the energy level:
1s22s22p6
So…find the highest energy level and add up
all the electrons in that level.
EXAMPLES
•
•
•
•
•
Calcium
Aluminum
Iodine
Oxygen
Iron
ENERGY
Electrons with the most
energy are located farthest
from the nucleus
Electrons with the
lowest energies are
located close to the
nucleus.
Quantum
A quantum is the
specific amount of
energy needed for an
electron to move
between energy levels.