This lecture is going to be about absorption and emission spectra. I'm assuming in this lecture
that you've seen the previous lecture on electron energy levels, and if you haven't I've left a
link to that in the video description. You'll remember from that lecture that electrons exist
at specific energy levels in atoms, and every single time an electron goes down an energy level
it emits a photon of that exact same amount of energy. So just observing this electron here,
I can see that every time it drops a certain amount of energy in electron-volts,
it emits a photon of the same energy. And if we want to increase an electron's energy
in an atom we have to add a photon to it where the photon is absorbed by the atom,
and the photon has to have that exact amount of energy required to get the electron to the correct
energy level because electrons cannot exist between energy levels.
In another lecture on photons which I’ve also linked in the description, I talked about how
if we know the energy of a photon we can also automatically know its wavelength using this
equation. Before I use that equation I just need to convert the energy to Joules. So when I do that
this is what I get. So I can find the wavelengths of each of these 3 emitted and absorbed photons,
and these are what I find them to be. So these are the wavelengths of photons which can be
absorbed or emitted by this atom, and if a photon that doesn't have one of those exact wavelengths
interacts with this atom, what's going to happen is the photon is just going to move
straight through the atom without affecting it and without being absorbed by it. That's going
to be a really important concept here. This is all review, we covered this in the previous lecture,
but this is going to be really central to this lecture about spectra.
I’m going to say that we have a source of white light and some kind of detector that can detect
all the different wavelengths of light that are impacting it, and white light actually contains
all possible wavelengths of photons, so if we shine this white light on this detector that
can detect which wavelengths are present, this is what it would look like. The detector would
record every possible wavelength of light as being present, because white light contains all possible
wavelengths, and in this detector we can imagine we order the wavelengths from smallest to largest
going from left to right. I'm going to shine that light on this detector again, but this time I'm
going to pass it through the element that we were looking at in the opening of this video. So that
element is now between the white light and the detector. So what would the detector detect now?
Well, every possible wavelength of light would come out of the white light source and begin to
fill up our detector with all the possible wavelengths of light, but I know there are
a few wavelengths that this atom can absorb. So what happens when one of those wavelengths
interacts with the atom? Because again, white light contains every possible wavelength,
so it also contains those 3 wavelengths. If that wavelength of 5.4x10^-7 meters appeared, the
atom could absorb this photon and it would become electron energy. So when that happens the photon
does not reach the detector, so there's going to be one very small part of our chart of wavelengths
that's going to be just a black line, because no light was detected at that wavelength, because the
light was kind of eaten up by the atom before it got to the detector. And if we're ordering
this light by the wavelength, that black line would appear right at 5.4x10^-7 meters, because
that was the wavelength absorbed by the atom. We can now observe the other wavelengths. Other
wavelengths would continue to pass through the atom, but when we hit the next wavelength that the
atom could absorb, that wavelength would also be absorbed into the energy of one of the electrons,
and so again a black vertical line would appear at that exact wavelength of 8.9x10^-7 meters.
More light would pass through and finally there would be one more black line,
because this wavelength of 1.4x10^-6 meters would also be absorbed by the atom. Finally
photons with a very long wavelength would pass through and complete the chart.
So we call this an “absorption spectrum.” An absorption spectrum shows the spectrum
of electromagnetic wavelengths passed through a substance, a collection of atoms, with dark lines
due to the absorption of specific wavelengths. So if I had a large amount of atoms of the substance
and I passed electromagnetic waves of every possible wavelength through the substance
and measured which wavelengths came out the other side, almost all wavelengths would make
it through the substance, but there would be a few very specific wavelengths that would not
make it through, that would cause black lines to appear on my diagram. I would call this diagram
an absorption spectrum for this specific atom. We can now imagine that we have atoms of the
same element, and this time we can place them in total darkness and see if we can measure any
of the photons being emitted by the element. So I know that if electrons in this element
were to lose energy, they could emit a photon of 5.4x10^-7 m and a photon of 8.9x10^-7 meters
and a photon of 1.4x10^-6 meters. So those would be the 3 places on my chart where I would detect
photons. Everything else would be black. There wouldn't be any other EM waves hitting this chart,
so we would call this an emission spectrum because this shows the spectrum of electromagnetic
wavelengths emitted by a substance with bright lines due to the emission of specific wavelengths.
You can see that the places where the lines show up on an absorption or an emission spectrum
are the wavelengths of photons which the substance can absorb or emit. In the last lecture we learned
how to find the energy, frequency, and wavelength of any photon that an atom can absorb or emit,
and here you can see that these spectra diagrams show us those wavelengths that the atom can absorb
or emit physically showing up on a chart we can make of all possible electromagnetic waves passing
through or emitted from the substance. Because all atoms of the same element have the same energy
levels, they absorb and emit the same photons, and so they produce the same absorption and emission
spectra. So as an example, all hydrogen atoms have these energy levels that I've listed here.
I can run through and find all the possible energies of the photons that can be absorbed
or emitted by a hydrogen atom, and I'll use the equation to find the wavelength. This is what I
get when I do that. So if I fill in an absorption or emission spectrum using these wavelengths, it's
going to look something like this. This is not drawn exactly to scale, but because all hydrogen
atoms have the exact same energy levels they all have the same possible energy changes, which means
they all emit or absorb the exact same photons with the same energies and the same wavelengths.
So every single time you make an absorption or emission spectrum of a hydrogen atom,
it's always going to look exactly the same, something like this, with the wavelengths
in the exact same spots. The spectrum of an element is like a fingerprint that we can use
to identify which elements light has passed through or been emitted from. As an example,
these are the emission spectra of different common elements. I'm just going to roll through,
and you can see how each one is kind of like a fingerprint for that element. We
can always identify which element we're seeing in an emission spectrum because we already know what
that element’s spectrum looks like, so if we're given an emission spectrum we can just match it
up with the correct element. This is actually how we know which elements exist in stars. We look at
the light from stars and see which wavelengths are missing and then match the missing wavelengths to
the spectra of elements. So this is how we know the chemical composition of the entire
universe. You can see in the picture that I’ve posted that this is the sun's absorption spectrum,
and you have all these different fingerprints of elements kind of superimposed on top of each
other, and we can just pick up the individual patterns and say “okay,
this set of wavelengths is produced by hydrogen, this is produced by iron,
by calcium, by magnesium, etc.” so this is how we know the exact chemical composition of the sun
and all stars that exist in the universe. This next idea is really central to the topic,
this gets asked about a lot on tests: absorption and emission spectra are evidence for the fact
that electron energy levels are quantized. We can understand why by asking what the spectra would
look like if electron energy were not quantized. This would mean that instead of existing at very
specific energy levels, electrons could exist at any level of energy around the nucleus.
It's difficult to say exactly what would happen here, because we would live in a very different
universe if this were true, but one thing that would happen is because the electrons
are able to exist at any energy level they can now absorb any wavelength of photon.
So any and all wavelengths could be absorbed by these electrons,
so it could be that the spectra would just be black all the time. It could also be that
maybe after reaching a certain limit of energy intake, and maybe when all the electrons have
escaped the nucleus because they've gained so much energy, whatever photons just happen
to randomly be left could pass through the nucleus and pop up as random lines on our chart.
I would imagine that if the same thing happened again because those wavelengths are random,
this would result in different wavelength readings on the chart,
so instead of getting the same fingerprint every time for each element, we would get just
random lines in random positions. This is probably not what would actually happen,
and it's probably just a matter of time before a professional physicist
hops into the comments and tells me that I’m wrong and something different would happen,
but we don't live in that universe so that doesn't really matter so much. What does matter
is that the one thing that would not happen if energy levels weren't quantized is the exact
same absorption and emission lines showing up for every single atom of the same element.
The fact that all atoms of the same element always show the exact same spectra pattern
means that the photon energies the electrons can emit or absorb are always exactly the same,
because if it's always the same wavelength that's absorbed, it's always the same photon energy
that's absorbed, and this means that the electrons can only experience very specific energy changes,
which means their energy must only exist in specific quantities and cannot exist between
those quantities. So again if we have the same spectra for all atoms of the same element, that
implies that the same wavelengths of photons are absorbed by every single atom of the same element,
which means the same energy of photons is absorbed, which means the electrons in each
atom of the same element can only change energy by the same specific amounts, which means that
the electron energy is quantized that can only exist in these very specific amounts, and cannot
exist between them. If it could exist between them then we would get additional wavelengths
that were absorbed or emitted, and that's not what we see in the real world. We always see the exact
same spectra patterns for all atoms of the same element. So this is how we know that electron
energy is quantized. So absorption and emission spectra are evidence that electron energy is
quantized. Just a few important notes to finish up the lecture: the way absorption and emission
spectra are drawn mistakenly implies that they only detect wavelengths on the visible spectrum,
from visible violet to red light. They're very often drawn as rainbows like what I've drawn here,
which implies that they only deal with visible light, but absorption and emission spectra
actually measure the full electromagnetic spectrum, so they can measure any wavelength
or frequency of photon, and remember visible optical light is only one very small part of
the much larger electromagnetic spectrum, because all photons with the same wavelength also have the
same frequency. We can also draw absorption and emission spectra with frequency instead
of wavelength. So because the speed of light is equal to wavelength times frequency, I can use
this equation to solve for the frequency along with the fact that this is the speed of light.
So if I wanted to convert this into a frequency diagram, this is what it would look like. I just
replaced all the wavelengths with frequencies instead, so we're also totally allowed to draw
a spectra with frequency instead of wavelength. This final note is hyper-specific, but each energy
change only occurs between two specific levels in an atom, and doesn't occur twice in the same atom.
this means that each line on an absorption or emission spectrum is associated with two
specific energy levels of the electron, and only those two. So as an example, these are 6 of the
energy changes for hydrogen, and each of these energy changes only appears once. As an example,
there's never going to be another change between two energy levels in hydrogen
that's equal to 10.2 electron-volts. That's just never going to happen, so that specific wavelength
on the spectrum is associated specifically with a change from the n = 1 to n = 2, or vice versa,
level of electron energy. So each individual line on the spectrum only represents a change in energy
between two very specific levels. You can see that the purple line is from n = 1 to n = 4.
The orange line is from n = 2, to n = 3 and there will never ever ever be a situation on a spectrum
where one line goes with 2 different pairs of energy levels like the orange line could never
possibly be associated with like n = 5 to n = 8. As an example, it can only be associated with one
pair of levels because each individual energy change only occurs once in an atom's energy
levels. And that's everything that you need to know about absorption and emission spectra.