Quantum mechanics has radically changed our
worldview.
By now all theoretical physics theories are
divided into two groups: classical and quantum.
Quantum mechanics forms the basis for all
modern theories of elementary particles and their interactions.
Some examples are:
Quantum electrodynamics, which describes electromagnetic
interactions of charged particles and photons.
Quantum chromodynamics, which studies the
interaction between quarks and gluons, the
building blocks of protons, neutrons and a
lot of other particles.
Promising theories like string theory are
also based on quantum mechanics.
A lot of other models and theories are based
on quantum mechanics.
Actually, there is no alternative for quantum
mechanics.
Its predictions confirmed with unprecedented
accuracy.
Not a single observation has ever been made
which would contradict the quantum theory.
But what made scientists look for a new theory
at the beginning of the 20th century?
Almost all observed phenomena could be explained
with Newtonian mechanics and Maxwell classical electrodynamics.
But unlike today, there were a few contradictions
between the predictions of classical theories
and actual observations.
The atomic stability problem is an example
of such a contradiction.
From the Rutherford experiments by the beginning
of the 20th century scientists knew that atoms
consist of positively charged particles (protons)
and negatively charged particles (electrons).
The most popular was the planetary model in
which the electrons go around the massive
nucleus just like the planets go around the
Sun.
But the calculations showed that such a system
would not be stable.
According to Maxwell's equations of classical
electrodynamics a charged particle moving
in a circle should emit electromagnetic waves.
But waves carry energy and according to the
energy conservation law it should be taken
from the electron kinetic energy.
So the electron should gradually lose its
velocity and finally fall down on the nucleus.
According to the calculations that would happen
within a tiny split of a second.
If that was true, the world as we know it
would not even exist.
By heating up a material and letting the light
go through the prism, one can see the so called emission spectrum.
This kind of experiments were popular in the
19th century and showed that atoms can emit
and absorb electromagnetic waves only of particular
frequencies.
For example, an atom can emit a photon of
frequency nu1 or nu2, but it cannot emit something
in between, like greater than nu1 and less
than nu2.
The emission spectrum of each atom is unique
just like our fingerprints.
This is how we know the chemical composition
of distant stars.
The light carries all the information.
It was found useful to draw the so-called
energy level diagram of an atom.
The frequencies of emitted and absorbed photons
can be found as energy differences of any two levels.
However the origin of these discrete energy
levels remained a mystery.
The existing theories could not give numerical
values for those quantized energy levels as they are called today.
Another discrepancy between the theory and
observations is known as the ultraviolet catastrophe.
The vibration of atoms in the crystal lattice
causes the emission of electromagnetic waves called thermal radiation.
At room temperature this radiation is emitted
in the infrared range, which is invisible
to the human eye, but can be seen with night-vision
devices.
Let's plot the dependence of radiation intensity
on its frequency.
We obtained a function with a definite maximum.
The precise form of the function depends on
the chosen temperature.
If a metal plate is heated to high temperatures,
then the radiation will be mostly in the visible red range ("red-hot").
So increasing the temperature will shift the
maximum to higher frequencies.
By further heating, it will move to the ultraviolet
region of the spectrum.
However, the classical physics predicts that
the intensity should keep up increasing with frequency forever.
Such a prediction contradicts everyday observations.
It shows the failure of classical physics,
which is reflected in the name of the paradox
- "the ultraviolet catastrophe".
Historically the birth of quantum mechanics
is associated with solving this problem.
It was Max Planck who succeeded in deriving
the formula correctly describing thermal radiation.
He assumed that energy is emitted only by
discrete portions - quanta.
The quantum of the electromagnetic field was
called a photon.
The energy of a photon is proportional to
its frequency.
The Planck constant serves as the coefficient
of proportionality.
Now it belongs to the list of fundamental
physical constants, just like the speed of light.
