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Chapter 4
Development of the New Atomic
Model
Properties of Light

What is Light?
The Wave Description of Light

Electromagnetic Radiation (ER)- a form of
energy that exhibits wavelike behavior as it
travels through space.
 Forms of Electromagnetic Radiation include:

gamma rays
UV rays
infrared radiation
Radio waves
X rays
visible light
microwaves
Together these forms of radiation are called
the ELECTROMAGNETIC SPECTRUM.

ALL FORMS of ER
move at the same speed
through a vacuum. The
speed of light in a vacuum
is a constant.
3.0 x 108 m/s
This is also the speed of
light through air.
c = speed of light = 3.0 x 108 m/s
Since the speed of all ER is the same, what is the difference
between these different forms of Electromagnetic Radiation?
Wavelength ( )
and
Frequency ()
Wavelength is the distance between
corresponding points on adjacent
waves.
Frequency is the number of waves that
pass a given point in a specific time.

Frequency is expressed in waves per
second.
 One wave per second is called a
HERTZ (Hz).


Therefore Hertz measures
Frequency
Frequency and Wavelength are
mathematically related.
c=
Since c is constant @ 3.0 x 108m/s,
the product  is always
constant.
This also means that  is inversely
proportional to .
Determine the frequency of light with
a wavelength of 4.257 x 10–7 m

7.0 x 10^14 Hz.

Two experiments done in the early
1900’s could not be explained by the
wave theory of light:
1. The Photoelectric Effect
2. The Hydrogen line-emission spectrum
The Photoelectric Effect

The emission of electrons from metal after the
absorption of energy from electromagnetic
radiation, including visible light.
 But not just any light …the light had to have a
minimal frequency. photoelectric effect
Since wave theory suggested that light of any
frequency should be able to knock the eloose, there were questions to answer.

led to important steps in understanding
the nature of light and electrons.

Study of the photoelectric effect
influenced the formation of the concept
of wave-particle duality: the concept
that all matter and energy exhibits both
wave-like and particle-like properties.
Max Planck
Suggested that objects emit
energy in QUANTA- small,
specific amounts of energy
 A QUANTUM is the
minimum amount of
energy that can be lost or
gained by an atom.

Planck’s Constant
E=h
Energy = Planck’s constant x frequency.
The Planck constant relates the energy in
one quantum (photon) of electromagnetic
radiation to the frequency of that radiation.
Equal to approximately
6.626176 x 10-34 joule*seconds
a constant
Albert Einstein
Suggested a dual wave-particle nature
of electromagnetic radiation.
So… light can behave as a wave or as a
particle.

Each particle of light carries a quantum
of energy which Einstein called a
PHOTON.
Photon- a particle of
electromagnetic radiation that
has Zero Mass and a quantum
of energy.
Ephoton = hv
In order to be ejected from a metal,
an electron must be struck by a
photon that has at least the
minimum energy required to knock
the electron loose.

This minimum energy corresponds to
the minimum frequency of light required.
Different elements release their
electrons at different frequencies.
 Because
different elements hold
electrons more or less tightly the
amount of energy (thus the
frequency of the electromagnetic
radiation) required to remove the
electron varies.
This discovery led to the
quantum revolution in physics
and earned Einstein the Nobel
Prize in 1921.
"for his services to Theoretical
Physics, and especially for his
discovery of the law of the
photoelectric effect"
THE
HYDROGEN-ATOM
LINE EMISSION
SPECTRUM




Ground State- The lowest energy state of an atom.
Excited State- An energy state in which the atom has
more potential energy than it has in its ground state.
The electrons have absorbed energy and have
moved from a GROUND STATE to an EXCITED
STATE
The electrons that are close to the nucleus in the
hydrogen atoms gain energy and jump one, two or
more energy levels higher.
When they return to the lower energy levels, the
hydrogen spectrum (light) is emitted.

A hydrogen discharge tube is a slim tube
containing hydrogen gas at low pressure with
an electrode at each end.
 If you put a high voltage across the
electrodes, the tube lights up with a bright
pink glow.
– What causes this light to be produced?

If the light is passed through a prism or
diffraction grating, it is split into its various
colors; its line-emission spectrum.

What you would
see is a small
part of the
hydrogen
emission
spectrum, only
the visible light
portion,
however, the
spectrum
continues into
the UV and IR
regions of the
spectrum.
Scientists at the time thought that a continuous
spectrum would be observed; a continuous range
of frequencies of electromagnetic radiation.
The actual results suggested that energy
differences between the atoms ground state and its
excited state were fixed.
The electron in a hydrogen atom exists only in
very specific states.
Ephoton = E2 – E1 = hv
Ephoton = Eexcited state – Eground state = h
line emission spectrum
http://jersey.uoregon.edu/vlab/elements
/Elements.
The BOHR MODEL of an Atom


Niels Bohr in 1913- suggested the atom as a small,
positively charged nucleus surrounded by electrons
that travel in circular orbits around the nucleus, similar
in structure to the solar system.
The energy of the particles in the Bohr
atom is restricted to certain discrete values,
the energy is quantized.
This means that only certain orbits with
certain radii are allowed; orbits in between simply don't
exist.
HOME FUN
PG 97 Q 1-4
PG 118 Q 1,4,6,7

Quantum Theory is born to explain the
observation of the line-emission
spectrum and the photoelectric effect.
The QUANTUM MODEL of
the ATOM
If light had a dual wave-particle
behavior…perhaps electrons
could too!!!
Louis de Broglie- suggested that electrons be
considered waves confined around the nucleus
of the atom.
Experimentation found that electrons can be
diffracted and experience interference, much
like light waves.
Diffraction-the bending of a
wave as it passes by an object,
causing a redistribution of energy.
Interference- Overlapping of
waves that combine to reinforce
or cancel each other.
The bright lines indicate constructive interference and the
dark lines indicate destructive interference.
Nobel Prize in Physics in
1929 for development of
wave mechanics.
Yeah Louie!
Heisenberg Uncertainty Principle- It is
IMPOSSIBLE to know both the position
and velocity of an electron or any other
particle at a given instant.
Sooo, when the position of an atom (or
electron) is measured, the measurement process
will leave the momentum of the atom changed
by an uncertain amount
In the quantum model the electron cannot be found
precisely, but we can predict the probability, or likelihood,
of an electron being at some location in the atom.
The Schrodinger Wave Equation
Schrodinger analyzed
what an electron would
look like as a wave
around the nucleus of
the atom.
He also had/has a cat.
Check it out.
He treated everything as waves, with each
electron having its own unique wave function, its
own position so to speak.
These wave functions are described in Schrödinger's
equation by three quantum numbers:
1. Principal quantum number
2. Angular momentum quantum number
3. Magnetic quantum number
Now there are four quantum numbers that specify the
properties of atomic orbitals and the properties of the
electrons in orbitals.
4. Spin quantum number
We use quantum numbers to describe the likely location of
some object (like an electron) in three-dimensional space
Principal Quantum Number
Symbolized by n
Positive integer
Ex. n=1, n=2
Indicates the main energy level
occupied by an electron.
As n increases, the electron’s energy
and average distance from the
nucleus increases.
Angular Momentum Quantum Number
Symbolized by l (el).
Zero plus all positive integers less than or equal to
n-1.
Ex- if n =1 then l= n-1=0
If n =2 then l= 1 or l = 0
Describes the shape of the orbital.
The value of l indicates the letter value of the orbital.
L=0
s orbital
L=1
p orbital with the s
L=2
d orbital with the s and p
L=3
f orbital with the s, p, and d
AKA sublevels
Each atomic orbital is designated by the principal number
and the letter of the sublevel.
Ex: 1s, The first main energy level.
2p- The p orbitals in the second main energy level.
Magnetic Quantum Number
Describes the orientation of the
sublevels in space (x, y, z axis)
Designated by m
Values for m depend on the sublevel.
In the s sublevel m = 0
In the p sublevel m = -1, 0, or 1
In the d sublevel m = -2, -1, 0, 1 or 2
s sublevels are spherically symmetric
around the nucleus.
p sublevels are dumbbell shaped and
fall on the x, y and z axis
d and f sublevels are complex, with
more than one shape.
Orbitron
http://winter.group.shef.ac.uk/orbitron/AOs/4f/index.html
Spin Quantum Numbers
Electrons spin on an axis.
The value for spin numbers is
either ½ or - ½ depending
on the direction of the spin.
A single orbital can hold a
maximum of two electrons, which
must have opposite spin.
Chapter 4 Section 3
Electron Configurations
Electron configuration- The
arrangement of electrons in an
atom.
ORBITALS- (1,2,3, etc.) can be
thought of as houses where
electrons reside.
SUBORBITALS- (s,p,d,f)
The rooms in the houses where
the electrons spend their time.
Rules for electron
configurations.
Electrons tend to assume
positions that have the LOWEST
energies…the ground state
electron configurations.
Each element has a unique
ground state e- configuration.
Aufbau Principle- An electron
occupies the LOWEST energy
orbital that can receive it. (n)
Pauli Exclusion Principle- No
two electrons in the same atom
can have the same set of
quantum numbers.
Hund’s Rule- orbitals of equal
energies are occupied by one
electron before any orbital is
occupied by a second, and all
electrons in singly occupied
orbitals must have the same spin
Why does 4s fill before 3d?
Energy
The d-orbitals experience less
of the effect of the increasing
effective nuclear charge
because of the shape of the
orbitals.
The shape of the s orbital (a
sphere) means the electron is
closer to the nucleus than the dorbital and so is stabilized by the
nucleus to a greater degree than
the d-orbital
Look at palladium, gold, silver.
These elements “steal”
electrons from the s sublevel to
fill the d.
This stabilizes the atom
making them less reactive.