Chapter 4 - Electrons
Properties of Light
What is light?
• A form of electromagnetic radiation:
• energy that exhibits wavelike behavior as it travels
through space
Electromagnetic radiation is classified
into two types:
• Non-ionizing Radiation:
• Transfers energy causing vibrations, electron excitation,
and heat
• Parts of the spectrum: radio, microwaves, infrared, and
visible
• Ionizing Radiation:
• High energy that ejects electrons and transforms
molecules into reactive unstable fragments
• Parts of the spectrum: UV, X-ray, Gamma Ray
Radiation is organized on the
Electromagnetic Spectrum
Wave Properties
Wavelength (λ) :
• distance between
corresponding points on
adjacent waves
• Unit: meters or nanometers
Frequency (ν):
• number of waves that pass a
given point per unit time
(usually 1 sec)
• Unit: 1/s = s-1 = Hertz (Hz)
Properties of Light
• How are frequency and wavelength related?
c = λν
c : speed of light (m/s)
c = 3.00 x 108 m/s
λ : wavelength (m)
ν : frequency of wave (s−1 = 1/s = Hz)
Max Planck proposed …
Energy needed for electrons to move was quantized,
or a quantum of energy is needed to move an
electron.
E = hν
E : energy of light emitted (J)
h : Planck’s constant (J·s)
h = 6.626 × 10−34 J·s
ν : frequency of wave (s−1 = 1/s = Hz)
E = hn
E = hn
Electrons can
only move with a
quantum of
energy.
Every color
represents a
different amount
of energy
released.
Louis de Broglie proposed …
• that electrons be considered waves
confined to the space around an atomic
nucleus.
• If electrons exist as waves, then they must
have specific frequencies
• If they have specific frequencies, they also
have specific energies
So energy is quantized!
ELECTRON BEHAVIOR
Who made this model of the atom?
Electrons are always moving!
• The farther the electron is from the nucleus,
the more energy it has.
• Electrons can change energy levels and
emit a photon of light.
• Photon: packet of energy
• Ephoton = hν
Electron Movement Vocabulary
• Ground State: electrons in the lowest possible energy
level
• Excited State: electrons absorb energy and move to a
higher energy state
• Emission : when an electron falls to a lower energy level, a
photon is released
• Absorption : energy must be added to an atom in order to
move an electron from a lower energy level to a higher
energy level
• How much energy must it absorb?
• Quantum: amount of energy needed to move between energy
levels
Electron Movement
• When electrons move from the excited state
to the ground state they emit a photon of
light
• The light emitted by an element is viewed as
a bright line emission spectrum
• Each band of light represents the energy
released by an electron when it moves from
higher to lower energy
Line Emission Spectrum of Hydrogen
Atoms
Because electrons move so much…
Heisenberg Uncertainty Principle:
We cannot know an electrons exact
speed and position at the same time
Schrodinger Wave Equation
• Equation that describes the wave & particle nature of an
electron
• Don’t know the exact position…
• Predict where electron will be most of the time
Where 90% of the
e- density is found
for the 1s orbital
Electron Location
DO NOW
• How would you describe the location of your home?
• Like giving your house an address, each electron in an
atom is assigned an address or location inside the
atom.
• Remember: Do electrons ever stop moving?
We do not know where the electrons are,
but we can describe where they might be
• The address of all the electrons in an atom is called an
Electron Configuration
Sodium #11
1s22s22p63s1
First Location: ENERGY LEVEL
n = 1, 2, 3, 4 …
ENERGY
LEVEL
SUBLEVEL
Second Location: SUBLEVEL
These are the probability regions called: s , p , d , f
SUBLEVEL
ORBITAL
Third Location: ORBITAL
s=1
p=3
d=5
5s
4p
Energy
3d
4s
3p
3s
2p
2s
1s
f=7
Sublevels on the Periodic Table
Examples:
• What is the electron configuration of Mg?
• What is the electron configuration of Cl?
How to fill in an
Orbital Diagram
Step 1 – Aufbau Principle:
Electrons are represented as arrows and drawn one at a time.
Electrons start at the lowest possible energy.
5s
4p
3d
Energy
4s
3p
3s
2p
2s
1s
Element:
Number of
Electrons:
Electron
Configuration:
Hydrogen
Step 2 – Pauli Exclusion Principle:
Electrons have opposite spins.
Draw one up and one down.
5s
4p
3d
Energy
4s
3p
3s
2p
Element:
Number of
Electrons:
Electron
Configuration:
2s
 Paired electrons
1s
Helium
Step 3 – Hund’s Rule:
Electrons occupy equal energy orbitals so that a maximum number of unpaired
electrons results
Draw electrons in each orbital first then pair them.
5s
4p
3d
Energy
4s
3p
3s
2p
2s
1s
Element:
Number of
Electrons:
Electron
Configuration:
 Unpaired electrons
Oxygen
What is the electron configuration of S?
5s
4p
3d
Energy
4s
3p
3s
2p
2s
1s
Element:
Number of
Electrons:
Electron
Configuration:
Sulfur
What is the electron configuration of Sr?
5s
4p
3d
Energy
4s
3p
3s
2p
2s
1s
Element:
Number of
Electrons:
Electron
Configuration:
Strontium
Valence Electrons
Why is location important?
location influences behavior
Valence Electrons
• valence electrons:
• electrons in the outermost energy level
• can be lost, gained, or shared
• electrons that are responsible for reactions
• Elements in a group have similar properties
because they have the same number of
valence electrons
How many valence electrons in each
group?
Only the main block elements!
1
8
2
3 4 5 6 7
Inner Core Electrons
• All electrons under the outer most level