
A. General Information
› 1. electron behavior has been studied
through light
› 2. remember, light IS radiant energy
› 3. originally considered to be wave energy
alone
› 4. in the 1900’s, scientists determined light
behaved like a particle
 So it is both!
I. Radiant Energy

B. Waves
1. Light waves are electromagnetic waves
• Called electromagnetic radiation (ER)
• X-rays, gamma rays, and radio waves are,
also, a part of ER
2. Electromagnetic waves consist of electric
and magnetic fields oscillating at right
angles.
a) All waves are described by 4 characteristics
Electromagnetic Radiation
I. Radiant Energy

B. Waves
2. Electromagnetic waves consist (CONT.)
a. all waves are described by 4
characteristics
1. Amplitude – the height of the wave measured
from its origin to its peak or crest
• The brightness or intensity of light is dependent on
this part of the wave.
2. Wavelength – the distance between successive
crests
• The distance traveled in a full cycle
• Visual light has a range between 400 to 750 nm
(10-9)
I. Radiant Energy

B. Waves
2. Electromagnetic waves consist (CONT.)
a. all waves are described by 4
characteristics
3. Frequency – how fast the wave
oscillates up and down (during a given
time, usually 1 second)
 The unit is cycle/s or hertz (Hz)
 1 Hz = 1 cycle/s
 Ex. FM Radio 93.1 MHz = 93.1 x 106 cycle/s
 Ex. Visual light is between 4 x 104 cycle/s and 7 x
1014 cycle/s
Wavelength and Amplitude
I. Radiant Energy

B. Waves
2. Electromagnetic waves consist (CONT.)
a. all waves are described by 4
characteristics
4. Speed – which is a constant value
• Called “the speed of light” = 3/00 x 108 M/S.
• This creates a direct relationship between
wavelength and frequency.
• The shorter the distance, the greater the
oscillations.
• The longer the distance, the fewer the osscilations.
I. Radiant Energy

B. Waves
2. Electromagnetic waves consist (CONT.)
b. The relationship between frequency and
wavelength is a mathematical expression.
 λ (lambda) = wavelength, V (nu) = frequency,
c = speed of light
 λ=v*c
 Ex. Helium – neon laser has a wavelength of
633 nm…v = ?
 4.74 x 10-14 s-1
I. Radiant Energy

C. Electromagnetic Spectrum
1. Prisms separate light into the different wavelengths
a. A rainbow is all of the light in the visible spectrum
(ROY G BIV)
• Violet has the shortest wavelength, Red has
the longest wavelength
b. Visible light constitutes of a very small portion of
the electromagnetic spectrum.
c. The rest of the electromagnetic spectrum is
invisible to the eye.
• Consists of Gamma rays, X-rays, UV, visible, infra red
(IR), Microwaves, TV waves, and Radio
• From smallest wave to longest wave
Visible Part of the Spectrum
II. Quantum Theory

A. General Information
› 1. Hot objects emit electromagnetic waves
(why?)
 a. first emits heat (IR energy/light)
 b. begins to glow (Red to yellow to white for
metal)
 Electric stove tops
› 2. Barium and Strontium emit green and red
colors (why?)
› 3. Gases give off specific colors of light when
heated
II. Quantum Theory

B. Planck’s Theory
1. Max Planck theorized the spectrum of radiation
emitted changes with temperature
2. Theorized energy emitted or absorbed is restricted
to “pieces” of particular size”
3. Proposed – There is a fundamental restriction on the
amounts of energy that an object emits or absorbs,
which are called quantum ( meaning fixed amount)
 Derived from the concept of the relationship between
frequency (v) and the energy (E) with which it is
associated
 Plank determined the energy constant, known as
Plank’s constant (h) with a value of 6.6262 x 10-34 J/s.
II. Quantum Theory

B. Planck’s Theory
4. Plank’s equation is E = H * v
5. Quantum of energy of extremely small, so it
looks like a continuous climb
II. Quantum Theory

C. Photoelectric Effect
1. Electrons are ejected from metal when light is
shined on it.
2. A minimum frequency of light is needed to
release electron
• Ex. Sodium metal wont release electrons with
red light, but will with violet light
3. Light consists of quanta of energy that
behave like tiny particles of light.
• Called photons
• Photon energy is equal to Plank’s energy
Photoelectric Effect
Photoelectric Effect
III. Atoms: A Second Look

A. Line Spectra
1. A line spectrum is a spectrum of colors
created from a prism
2. Elements emit light when they are vaporized
3. Each element has a unique line spectrum
•
An atomic fingerprint
4. Each element when placed under a flame
appears as a color
•
•
•
•
Salt – yellow, because of sodium
Lithium – red
Potassium – blue w/ red
Neon – red
•
Nitrogen - orange
Line Spectra for H, Ne, and Fe
III. Atoms: A Second Look

B. Bohr Model (Neils Bohr)
1. Used Rutherford’s planetary model to help
explain element emissions
2. Orbitals around the nucleus were based on
quanta and given a quantum number
• Lowest level, n = 1 (Ground state)
3. When the electron absorbs enough energy,
it will jump to the next energy level (excited
state)
• n = 2, 3, 4, etc.
4. Light is emitted as the electron “falls” back
to the ground state
Bohr Model
III. Atoms: A Second Look

C. Matter Waves
1. Louis De Broglie theorized that matter has a
dual nature.
•
•
•
•
•
Believed matter should have wave like behavior
and exhibit wavelengths
Called it matter waves
Came up with a mathematical formula relating
the mass and velocity of a moving particle and
its possible wavelength.
This finding was used to create the electron
microscope.
For waves to be seen from objects the mass must
be very small.
III. Atoms: A Second Look

D. Heisenberg’s Uncertainty Principle
1. Werner Heisenberg proposed that the
position and the momentum of a moving
object cannot be simultaneously measured
and known exactly.
• Uncertainty Principle
2. Hard to predict where a particle will be in
the future.
IV. New Approach to the Atom

A. General Information
1. 3 known concepts shape the new look
a. Energy is quantized
b. Electrons exhibit wavelike behavior
c. Impossible to know exactly where an
electron is in space
2. These 3 concepts lead to the quantummechanical model
IV. New Approach to the Atom

B. Probability and Orbitals
1. General Information
a. Consider an electrons place around the
nucleus as a blurry cloud
b. The cloud’s density is greater where there is
a higher probability that the electron is
present
• Called electron density
IV. New Approach to the Atom

B. Probability and Orbitals
1. General Information
c. An atomic orbital is a region around the
nucleus of an atom where an electron with a
given energy is likely to be found
• Orbitals have characteristic shapes (not energy)
• Draw orbitals based on where they are likely to be
located 90% of the time
• Different orbitals are designated by different letters
•
•
•
•
S, p, d and f
S = spherical shaped
P = dumbbell shaped
D & f = complex
Orbitals S, P, and D
IV. New Approach to the Atom

B. Probability and Orbitals
2. Orbitals and Energy
a. The Principle Energy Levels in an atom are
designated by the quantum number (n)
•
n is the principle quantum number
IV. New Approach to the Atom

B. Probability and Orbitals
2. Orbitals and Energy
b. The energy of the electron increases as n increases
(1, 2, 3, 4, 5, 6)
• Each energy level is divided into one of more
sublevels
• The number of sublevels in each principle energy level
= quantum number
• For example 1 = 1 sub, 2 = 2 sub, 3 = 3 sub, etc
• The sublevels are indicated by a letter. For example
• n = 1; 1s
• n = 2; 2s 2p
• n = 3; 3s 3p 3d
• N = 4; 4s 4p 4d 4f
IV. New Approach to the Atom

B. Probability and Orbitals
2. Orbitals and Energy
b. The energy of the electron increases as n increases
(1, 2, 3, 4, 5, 6)
 The number of orbitals in each sublevel is always
equals the quantum number
 n=1; 1s : 1 spherical orbital
 n=2; 2s : 1 spherical orbital (larger than 1s)
2p: 3 bell shaped orbitals
 n=3; 3s I spherical orbital (larger)
3p: 3 bell shaped orbitals (larger than 2p)
3d: 5 complex (d) orbitals
 n=4; 4s: 1 spherical orbital (larger()
4p: 3 bell shaped orbitals (larger)
4d: 5 complex orbitals (larger than 3d)
4f: 7 complex orbitals
Orbitals S, P, D, and F
IV. New Approach to the Atom

B. Probability and Orbitals
3. Electron Spin
a. Electrons spin on their access (2 ways only)
• Can spin clockwise and counterclockwise
b. Spinning charges create magnetic fields
• Clockwise in N ↑
• Counterclockwise is N ↓
c. Can have parallel spins or opposite spins
• If opposite spins, cancel the magnetic pull
• If parallel, they create magnetic effect
d. Pauli exclusion principle
• Each orbital in an atom can hold only 2 atoms with
opposite spins.
Opposite Spins w/ Mag fields
V. Electron Configurations

A. General Information
1. Electron configuration is the distribution of
electrons among the orbitals of an atom
2. Electron configurations describe where the
electrons are found and what energies
they possess
3. Electron configurations of atoms are
determined by distributing the atom’s
electrons among levels, sublevels, and
orbitals based on a set of stated principles.
V. Electron Configurations

B. Determining Electron
Configurations
1. Easy once you learn the energy levels of
the orbitals within each principle energy
level, the s-sublevel is the lowest level.
2. When electrons populate the lowest
energy orbitals, they are in the ground
state.
Example of Electron Configuration
V. Electron Configurations

B. Determining Electron Configurations
3. The electron locations can be predicted by
using the Aufbau principle, the Pauli Exclusion
principle, and Hund’s Rule
a. Aufbau – electrons are added one at a time, to the
lowest energy orbitals available
• Until all electrons are accountable
b. Pauli Exclusion principle – An orbital can hold a maximum
of 2 electrons
• To occupy the same orbital, the electrons must have
opposite spins (called paired electrons)
c. Hund’s – Electrons occupy equal-energy orbitals so that a
maximum number of unpaired electrons result.
Aufbau & Hund’s
Aufbau Principle
Short-cut to Electron
Configurations
V. Electron Configurations

B. Determining Electron Configurations
› 4. Arrows represent electrons, boxes represent
orbitals called orbital diagrams (show electrons
in orbitals)
 Ex. 6C : 1s
↑↓
2s
↑↓
2p
↑
↑
↑↓
 ↑ = counterclockwise spin
 ↓ = clockwise spin
e. electron configurations are created from electron
diagrams (and vice versa)
ex. 6C
= 1s22s22p2
• Exponents give you the number of electrons in each
energy level.
Orbital Diagram– C
Correct and Incorrect Orbital
Diagrams
Orbital Diagram– P
V. Electron Configurations

C. Exceptions to Aufbau’s Principle
1. Some elements don’t follow the rule
a. “They are interesting”
• Ex. Cr-23, Cu-29
• Expected: Cr-23: 1s22s22p63s23p64s23d4
Actual: 1s22s22p63s23p64s13d5
• Expected: Cu-29: 1s22s22p63s23p64s23d9
Actual: 1s22s22p63s23p64s13d10
b. Cause by interactions of electrons in
orbitals with very similar energies
Periodic Table and S, P, D, and
F Orbitals