It's difficult to think about atoms without picturing them.
Of course, as a physicist I know the image in my head isn't really
what atoms look like.
Still, these pictures can be useful as long as you don't take them too literally.
To really understand atoms, we do experiments to discover how they behave.
One experiment excites a tube filled with hydrogen
gas using electricity so that it emits a coloured light.
If we look at the light through a diffraction grating,
we see that the light is made up of narrow bands of colour.
This is its spectrum.
But notice how the bands of light are very narrow
and separated by dark gaps of no light at all.
Different gases will have different spectra.
Here is argon, and here's neon.
Unlike a rainbow, the spectrum of a pure element
is made up of a very few specific wavelengths.
Niels Bohr proposed a model of the atom that took these results very seriously.
In the Bohr model,
the atom's electrons are distributed in energy shells surrounding the nucleus.
Electrons only change shells if they absorb or release
a specific amount of energy.
The specific wavelengths of light we observe are the result of electrons
emitting energy as they change from a higher shell to a lower one.
Because every element has a different number of protons
and electrons, every element will produce its own unique spectra.
The big step forward with Bohr's model was that it
described the way the energy of electrons is quantized.
But here's the thing.
There were many reasons that physicists, including Bohr,
knew that this model wasn't the full picture.
In fact, a variation of this same spectra experiment revealed
that electrons could behave in ways the Bohr model couldn't explain.
Let's repeat the same experiment with neon gas.
Only this time we'll add an electromagnet so we can surround the gas tube
with a magnetic field.
With the electromagnet off, we see the distinct spectral lines for neon.
When we turn the electromagnet on, we see individual lines
split into three or more bands.
In the presence of a magnetic field, the electrons
emit slightly different wavelengths of light
This result is called the Zeeman effect.
It's subtle, but very significant.
More importantly, the Bohr model can't explain it.
The experiment showed that electrons
have a magnetic behaviour that we need to understand better.
Physicists realized that because electrons have a charge
they would create a magnetic field as they move.
And if they were rotating in an atom, that is, if they have angular momentum,
then the magnetic field would be similar to that of a bar magnet.
In physics, angular momentum is always conserved.
For example, if I sit here holding the axle of a spinning wheel
and I tilt the axle, my stool turns in response.
That response is angular momentum being conserved.
The Einstein-de Haas experiment tests to see if electrons have angular momentum.
In the experiment, we take an iron cylinder and suspend it from a thread.
Then we surround the bar with a coil of wire.
Sending an electric current through the coil creates
a magnetic field along the length of the iron bar.
The moment the magnetic field is turned on, we observe a rotation in the iron bar.
This behaviour is angular momentum being conserved.
The magnetic field causes a change in the angular momentum of the electrons
inside the cylinder, and the cylinder rotates in response.
The experiment shows that electrons have some kind
of angular momentum, which physicists call spin.
We call the behaviour spin.
But for all kinds of reasons, electrons cannot actually spin
like a wheel or a top.
Even though the picture of a spinning particle is intuitive and we often picture
electrons this way, electron spin is something else entirely.
We know this because of another experiment.
In the Stern-Gerlach experiment, a piece of silver is heated in an oven
to create a beam of silver atoms that is directed at a screen.
Between the oven and the screen, two different-shaped
magnets, create a non-uniform magnetic field along the beam path.
The experiment tests how electrons are affected
by this magnetic field by looking at where the silver atoms arrive at the screen.
If electron spin were like a spinning top with each axis oriented randomly,
the silver atoms would be deflected by different amounts
and arrive at the screen in a continuous distribution.
But that's not what we observe.
What actually happens is that the silver atoms hit the screen
in only two zones, one up and one down.
This surprising result shows that electron
spin doesn't have many possible values, but only two.
Unlike spinning objects like wheels or tops, electron spin is quantized,
meaning that it is limited to only two possible values.
It took physicists a while to sort all this out, but spin
along with principal number, magnetic number and orbital number
constitute a full description of an electron in an atom.
It wasn't until we had sorted out the complete description of electrons
that experiments like the Zeeman effect could be explained.
Of course, we're still left with the problem of trying to picture atoms.
Imagination can lead us to important insights,
but if we cling to these pictures too closely, they can mislead us too.
Ultimately, scientists model atoms with mathematical equations,
and these describe what we observe
about atoms, but not necessarily what they look like.
When it comes to atoms, pictures are never the whole picture.