Atomic Physics
Late 1800’s
Scientists thought that the understanding of classical
physics was nearly complete.
(Yeah! We’re almost done!)
In Newton’s time, light was thought of as a particle,
but Huygens, Young, and Maxwell seemed to verify
that light was a wave once and for all.
One of a few remaining
questions…
Scientists were concerned about understanding the
glow of objects when they reach high temperatures.
All objects emit electromagnetic radiation – how much
depends on their temperature and other properties of
the object.
Typically, objects emit a continuous distribution of
wavelengths from infrared, visible, and ultraviolet parts
of the EM spectrum.
The intensity varies with temperature.
Blackbody Radiation
Physicists describe the radiation from a blackbody, a
perfect radiator and absorber that emits radiation
based only on its temperature.
Low temps  Warmer  Higher temps Even hotter
Infrared
Yellow glow White glow
(not visible) (molten metal) (light bulb)
Bluish glow
λ Peak Shifts with
Temperature
PROBLEM
As temperature increases, the peak of intensity shifts to shorter
wavelengths.
As temperature increases, the total energy given off also
increases.
CLASSICAL PHYSICS COULD NOT EXPLAIN THE LAB
RESULTS !!!!!
Enter Max Plank
In 1900, came up with a formula that was in agreement
with all of the data collected.
En = nhf
where h = 6.63 x 10-34 Js (Plank’s constant)
The n is an integer number.
BIG IDEA: ENERGY COMES IN
DISCRETE UNITS  IT IS QUANTIZED
Quantum Mechanics
This is the beginning of quantum
mechanics
Quantum = “chunk” of energy
The plural of quantum is “quanta”.
Plank believed that light existed as
waves, but emission and absorption
occurred in quantum chunks.
Quanta
Matter is quantized – whole number multiple of mass
of one atom.
Charge is quantized – whole number multiple of
charge on one electron.
Albert Einstein went further and proposed that light
itself is composed of quanta.
…introduced the idea of ‘photons’: No rest energy,
travel at one speed only – the speed of light!
Photons – particles of light!
Light is emitted, not continuously, but as a stream of
photons, each with energy hf. E = hf
Energy can be expressed in Joules.
When dealing with parts of atoms, energy is usually
expressed in eV (electron volts)
1 eV = energy that an electron or proton gains when
accelerated through a potential difference of 1 volt.
1 eV = 1.60 x 10-19 Joules
The Photoelectric Effect
Electrons from certain metals are
ejected when certain frequencies
of light falls upon them.
Low frequencies of light (red)
never caused electrons to eject,
even if very bright.
Higher frequencies of light
(greater energy) always did, no
matter the intensity (brightness).
Einstein’s Explanation
Einstein explained the photoelectric
effect by thinking of light in terms of
photons.
One photon is absorbed by each
emitted electron. The energy
absorbed by an atom is an all or
nothing process.
The NUMBER of photons striking
the metal is not important. But if the
energy in the photon is large enough,
then an electron is emitted.
Energy of a Photon
Remember E = hf for each photon. If frequency is
high, then energy of photon is high.
Blue, violet light – high frequency, so photons have
enough energy.
Red, orange light – no matter how bright, photons
don’t have enough energy.
Radical Idea Accepted
This idea was verified 11 years later by Robert Millikan.
Einstein received the Nobel Prize for this idea in 1921.
Electrons are emitted almost instantaneously.
Quantization of energy must be considered a real
description of the physical world, not just a
mathematical one.
True nature of energy seen in submicroscopic level of
atoms and molecules where quantum effects become
more important and measurable.
Early Models of the Atom
1897 – J. J. Thompson –
electrons imbedded in a
spherical volume of
positive charge (like seeds
of a watermelon or
“plumb pudding”).
Conclusion: Atom has
positive and negative
charges.
Early Models of the Atom
1911 – Ernest Rutherford – performed important
experiment
Most of the alpha particles passed through like empty
space, some were deflected at large angles, a small
amount were deflected backwards!
Conclusion of Rutherford’s
Gold Foil Experiment
“It was almost as incredible as if you fired a 15 inch
shell at a piece of tissue paper and it came back and hit
you!”
Conclusion: All the positive charge and nearly all of
an atom’s mass is concentrated at the center of the
atom. Rutherford called this the nucleus.
Any electrons in the atom were assumed to be in the
relatively large volume around the nucleus, like planets
orbiting around the sun.
Atomic Spectra
When made to emit light, every element has its own
characteristic color. (Neon gas in glass tube with
voltage applied!)
When light given off by an atomic gas is passed
through a prism, a series of distinct bright lines is seen.
Each line corresponds to a different wavelength, or
color of light.
These lines are known as an emission spectrum.
Emission and Absorption
Spectra
The wavelengths contained in a given spectrum are
characteristic of the element giving of the light.
(‘fingerprints’ to identify the element)
Similarly, an absorption spectrum can be obtained by
passing light containing all wavelengths of light
through a vapor of the element being analyzed.
Niels Bohr
In 1913, Niels Bohr came up with a model of the atom
that explained atomic spectra.
Electrons move in circular paths around nucleus, held
in orbit by the positive charge of the nucleus.
Only certain orbits are allowed…electrons are never
found in between these orbits, or energy levels.
When an electron absorbs external energy, the electron
is boosted to a higher energy level. We say it is in an
excited state.
An electron only radiates energy when it jumps from
an outer orbit to an inner one (from high energy level
to a lower one).
The frequency of radiation is related to the change in
the atom’s energy
Einitial – Efinal = hf
Atomic Spectra Explained
In Bohr’s model,
transitions between
stable orbits with
different energy levels
account for the distinct
spectral lines.
The frequency of the
emitted photon (its
color) is proportional to
the energy transition of
the electron.
The Nature of Light
Light is BOTH a wave and a particle.
It behaves like waves when it travels in empty space.
It behaves like particles when it interacts with matter
Most scientists accept both models and believe that the
true nature of light cannot be explained in a simple
classical picture.
Logical Question
If waves can have particle
properties, can particles have wave
properties?
Louis de Broglie asked this
question in his 1924 PhD
dissertation that earned him the
Nobel Prize in Physics
De Broglie Wavelength
De Broglie proposed that all matter could be viewed as
having wave properties: electrons, protons, atoms,
marbles, humans!
De Broglie wavelength = Plank’s constant
Momentum
λ= h = h
p
mv
An object with a large mass would have a large
momentum, and so a small wavelength.
Particles With Wave Properties
A tiny particle, like an electron would have a detectible
wavelength.
Years later, a beam of electrons was found to have
diffraction and interference patterns, just like a beam
of light! De Broglie was right!
Explaining Bohr’s Model
With Bohr’s atomic model, no one could explain why
electrons could exist only in specific orbits.
De Broglie saw a connection between his theory of
wave characteristics of matter and the stable orbits of
the Bohr model.
An orbit was only stable if it contained an integral
(whole) number of electron wavelengths! 1st orbit had
1 λ, 2nd orbit had 2 λ, 3rd orbit had 3 λ, etc.
Only integers of wavelengths were possible!
Only an integer number of λ
Certainty
In classical mechanics there is no limit to the accuracy
of measurements.
You can make more precise measurements by using a
more finely detailed meter stick, for example
In quantum mechanics, this idea doesn’t hold
This is NOT due to limits of instruments or to us
perturbing the system when we make measurements.
There is a fundamental limit in nature due to the wave
nature of particles.
Heisenberg Uncertainty Principle
The more we learn
about a particle’s
momentum, the less we
know about its position.
The reverse is also true.
It is impossible to
simultaneously
determine a particle’s
position and
momentum with
infinite accuracy.
An electron’s location is described by a probability
distribution – electron cloud