Atomic
Structure and
Periodicity
Advanced Chemistry
Ms. Grobsky
Food For Thought

Rutherford’s model became
known as the “planetary model”

The “sun” was the positivelycharged dense nucleus and the
negatively-charged
electrons
were the “planets”
The Planetary Model is
Doomed!

The classical laws of motion and gravitation
could easily be applied to neutral bodies like
planets, but NOT to charged bodies such as
protons and electrons


According to classical physics, an electron in
orbit around an atomic nucleus should emit
energy in the form of light continuously because
it is continually accelerating in a curved path
Resulting loss of energy implies that the electron
would necessarily have to move close to the
nucleus due to loss of potential energy

Eventually, it would crash into the nucleus and
the atom would collapse!
The Planetary Model is Doomed!
Electron crashes into the nucleus!?
Since this does not happen, the Rutherford
model could not be accepted!
The Bridge Between the
Planetary Model and the Bohr
Model
 Atomic
structure was often elucidated by
interaction of matter with light

Classical wave theory of light described
most observed phenomenon until about
1900
But What Exactly is Light?
 Light
is a form of ELECTROMAGNETIC
RADIATION
 A form of energy that exhibits wavelike
behavior as it travels through space
 Does
 In
not require a medium to travel through
a vacuum, every electromagnetic
wave has a velocity (speed) of 3.00 x 108
m/s, which is symbolized by the letter “c”
Some Properties of Waves

Wavelength (λ)



Frequency (ν)



Number of waves that
pass a given point per
second
Measured in hertz (sec-1)
Speed ( c )


Distance between two
consecutive peaks or
troughs in a wave
Measured in meters (SI
system)
Measured in meters/sec
Amplitude (A)

Distance from maximum
height of a crest to the
undisturbed position
Relationships of EM Wave
Properties

The wavelength and frequency of light are
inversely proportional to each other



Wavelength and frequency are related via
the speed of light in a vacuum (c)


As wavelength increases, frequency decreases
As wavelength decreases, frequency increases
c = 3.00 x 108 m/s
Speed of light in a vacuum is a constant
c = λ· ν
The Electromagnetic
Spectrum
 Electromagnetic
spectrum is the range of all
possible frequencies of electromagnetic radiation
 The highest energy form of electromagnetic
waves is gamma rays and the lowest energy form
is radio waves
Relationship of EM Wave
Properties
c = λ· ν
Max Planck
Around the year 1900, a physicist
named Max Planck introduced his
hypothesis of the quantum behavior
of radiation
•
•
•
A major turning point in physics!
But what is quantum???
Max Planck’s Obituary
Take
a few minutes to survey the reading
(skim/scan the text)
 Turn headings and subheadings into
questions related to who, what, where,
when, why, or how?
Write these questions on the LEFT side of the
reading
 Write down key ideas found during reading on
RIGHT side of reading (answers to questions,
significant information, reflections)

Summarize key ideas at BOTTOM of
reading

Planck and Quanta
 Planck
studied the energy given off by
heated objects until they glow
 He made a wild assumption that there is a
fundamental restriction on the amounts of
energy that an object emits or absorbs

He called these pieces “quanta”
 To
understand quantization, consider
walking up a ramp versus walking up the
stairs

For the ramp, there is a continuous change
in height whereas up stairs, there is a
quantized change in height
More on the Idea of Quanta
 The
energy possessed by the wave is only
related to the frequency of the wave
 The frequency of an electromagnetic
wave can be converted directly to
energy by:
h
= Planck’s constant = 6.626 x 10-34 J· s
Planck’s Constant
 Planck’s


constant, h, is just like a penny
Planck determined that all amounts of
energy are a multiple of a specific value, h
This is the same as saying that all currency in
the US is a multiple of the penny
Practice, Practice, Practice
 Calculate
the energies of one photon of
UV (λ = 1 x 10-8 m), visible (λ = 5 x 10-7 m),
and IR (λ = 1 x 10-4 m).
And then there was a problem…
• In the early 20th century, several effects were observed
which could not be understood using the wave theory of
light
• Every element emits light when energized either by
heating the element or by passing electric current
through it
• Elements in solid form glow when they are heated
• Elements in gaseous form emit light when electricity
passes through them
Einstein and the Photoelectric
Effect

Another observation that could not be explained via
the wave theory of light: The Photoelectric Effect

Electrons are attracted to the (positively charged)
nucleus by the electrical force

In metals, the outermost electrons are not tightly bound,
and can be easily “liberated” from the shackles of its
atom


It just takes sufficient energy
If light was really a wave, it was thought that if one
shined light of a fixed wavelength on a metal surface
and varied the intensity (made it brighter and hence
classically, a more energetic wave), eventually,
electrons should be emitted from the surface
Photoelectric Effect
“Classical” Method
Increase energy by
increasing amplitude
What if we try this ?
Vary wavelength, fixed amplitude
electrons
emitted ?
No
No
No
No
electrons
emitted ?
No
Yes, with
low KE
Yes, with
high KE
• No electrons were emitted until the frequency of the light
exceeded a critical frequency, at which point electrons were
emitted from the surface!
(Recall: small l  large n)
Einstein’s Theory






Einstein proposed an alternative theory to
the classical wave theory of light
He used Planck’s idea of energy quanta to
understand the photoelectric effect
Light exists as ‘quanta’ of energy (specific
amounts)
These quanta behave like particles
Light ‘particles’ are known as photons
Each photon carries an amount of energy
that is given by Planck’s equation
Einstein’s Photons and the
Photoelectric Effect

The light particle must have sufficient energy to “free” the electron
from the atom

Increasing the Amplitude is simply increasing the number of light
particles, but its NOT increasing the energy of each one!

However, if the energy of these “light particle” is related to their
frequency, this would explain why higher frequency light can knock
the electrons out of their atoms, but low frequency light cannot
“Light particle”
Before Collision
After Collision
The Dual Nature of Light
A “Waveicle”
Light travels through space as a
wave
 Light transmits energy as a particle
 Each photon carries an amount of
energy that is given by Planck’s
equation
ℎ𝑐
𝐸𝑝ℎ𝑜𝑡𝑜𝑛 = ℎν =
λ

So is Light a
Wave or a
Particle ?
• On macroscopic scales, we can treat a large number of
photons as a wave
• When dealing with subatomic phenomenon, we are often
dealing
with a single photon, or a few
• In this case, you cannot use the wave description of light
• It doesn’t work!
The Dualism of Light
 Dualism
is not such a strange concept
 Consider the following picture

Are the swirls moving, or not, or both?
But How is This Related to
the Atom?
Light and the Dilemma of
Atomic Spectral Lines
Experiments
show that
when white light is passed
through a prism, a
continuous spectrum
results
Contain
all
wavelengths of light
When
a hydrogen
emission spectrum in
visible region is passed
through a prism, a line
spectrum results
Only
a few
wavelengths of visible
light pass through
Seeing Atomic Spectral Lines
 Use
your diffraction grating to observe the
atomic spectra of:



Hydrogen
Oxygen
Neon
hydrogen (H)
mercury (Hg)
neon (Ne)
Planck’s Quanta and
Atomic Spectra


To produce a line spectrum, the electrons
in an atom move between energy levels
Electrons typically have the lowest energy
possible (ground state), but upon
absorbing energy via heat or electricity:


Electrons jump to a higher energy level,
producing an excited and unstable state
Those electrons can’t stay away from the
nucleus in those high energy levels forever so
electrons would then fall back to a lower
energy level
Just a Thought…
But
if electrons are going from
high-energy state to a lowenergy state, where is all this
extra energy going?
Connecting Planck’s Quanta
to the Atomic Model
 Energy does not disappear
 First Law of Thermodynamics!
 Electrons
re-emit the absorbed
energy as photons of light

Difference in energy would correspond
with a specific wavelength line in the
atomic emission spectrum

Larger the transition the electron makes, the
higher the energy the photon will have
Just for Thought

How is energy related to
wavelength?
Summary of Quanta and
Atomic Spectra

Atoms must somehow absorb energy and then give
the energy off in the form of light


Excited electrons in an atom return to lower energy
states
Each element has a unique emission spectrum

Electron movements create the specific colors that we
witness
Only certain energies are possible



Electron energy levels are quantized!
Thus, the electron arrangement in every element is
unique!
∆𝐸 = ℎν =
ℎ𝑐
λ
Just for Thought…
 Can
we map the electrons by using these
energy relationships from the emission
spectrum?
Neils Bohr and the Atomic
Model


The answer is YES!
Neils Bohr was one of the first to see some
connection between the wavelengths an
element emits and its atomic structure

Related Planck’s idea of quantized energies
to Rutherford’s atomic model
Bohr and the Atomic Model

Bohr discovered that as the electrons in the
hydrogen atoms were getting excited and
then releasing energy, only four different color
bands of visible light were being emitted: red,
bluish-green, and two violet-colored lines


If electrons were randomly situated, as
depicted in Rutherford’s atomic model, then
they would be able to absorb and release
energy of random colors of light
Bohr concluded that electrons were not
randomly situated

Instead, they are located in very specific
locations that we now call energy levels
Bohr model of the Hydrogen
Atom
• Protons and neutrons compose
the nucleus
• Electrons orbit the nucleus in certain
well-defined ‘energy levels’
Niels Bohr
nucleus
Many Electron Atoms
 Recall
that because each element has a
different electron configuration and a
slightly different structure, the colors that
are given off by each element are going
to be different

Thus, each element is going to have its own
distinct color when its electrons are excited
(or its own atomic spectra)
 Flame
Test Lab!