alright, we're going to talk about the
formation of spectral lines... both emission and absorption lines. I want
to remind you of this image 'here'... this very simplified model of
(Bohr's model of) the atom. For excitation of an electron who is
exciting it? That's not really obvious from the diagram, is it? It's
a 'packet' of light that comes by... of just the right energy to 'boost' the
electron to a higher energy state. The correspondence is... that when the
electron comes cascading downward again... back to its lower energy state... (that's
just the tendency of all material in this
universe that we live in... to seek out its lowest energy state) when it de-excites...
it must do so by liberating part of its energy. It does so by emitting
a photon of the exact same wavelength and energy.
Let's sum all that up... in text, that we can write down...
put on flash cards and memorize! Absorption.... that means the atom is
absorbing energy, yes? It can boost the electron from
the second or to the second or even higher
energy state, possible. You're probably wondering how many
energy states are there? ...that are possible in an atom?
That's the part that Bohr didn't have 'quite right' in his simplified model.
The energy states were 'one', 'two', 'three', 'four' and they're all about
equally spaced. Today, we know that there's an infinite
number of possible energy states! But here's the 'rub'.... they get
closer and closer to each other as you get higher and higher and higher.
So, even though they go up to infinity they all get 'shoved' closer and closer to
each other. The last orbit is not out toward
infinity.... it's actually quite close to the nucleus of the atom!
If you get above that 'threshold energy', you'll eject the electron
altogether... something that we call ionization.
Okay, but let me not get ahead of myself!
The electron comes 'back down' again... to a lower energy state. We say
the atom is 'decaying' or de-exciting. It can either jump from...
let's say it's the third or fourth or fifth energy state.... down to the ground
state. That's a possibility! Or, it can cascade back down to the
ground state from 'three' to 'two'... 'two' to 'one' (I guess one would be the
ground state), but it can do it in 'steps' as opposed to all in 'one fell swoop'. This little image 'here' (which is from
your current textbook) does a great job explaining the two
different possibilities here. There's even more, but they're the
general possibilities! In part (a) here, it's 'direct' decay so what
happens? An ultraviolet photon (UV photon) comes in.
That's what it 'takes' to boost Hydrogen's electron from the ground state
to any of the higher energy states. It's a big jump to the first
excited state! It takes the energy of an ultraviolet photon.
It shows that this atom is 'boosting' its electron to this
first excited state, which we're going to call energy level number 'two' (the first
excited state). Then, very quickly thereafter, what
happens? Follow the red arrow to the right... then almost instantly, the electron goes
cascading back down again to the ground state.
If it took a very specific UV photon to boost it from the ground state to the second state... then the exact same frequency (exact same energy) photon is given off by
the reverse... cascading from level two back down to level one!
Okay, that's 'direct decay.' In frame (b), down here, the 'cascading'
process, this shows an ultraviolet photon coming in
and enough energy to 'not' excite into the first excited state BUT to the second
excited state! So, there is a certain corresponding energy to go from
ground state to the second state... and 'it' found 'that' exact frequency and it
'noticed' it and 'grabbed' it... out of thin air! Yeah, it was passing through.
Then, what happens over on the right is one of two possibilities: it's possible for that electron to go from level
'three' down to level 'one' (the second excited state) down to the 'ground state'.
We say in 'one fell swoop' and give off the same
corresponding ultraviolet photon that it just absorbed!
Or, it can do it in a two-step process. It can cascade from level 'three' to level
'two' giving off... (which would be the red
visible light photon) that Hydrogen is so famous for giving off!
Then, it could cascade from level two down to level one,
as well! That's a bigger jump, even though the cartoon doesn't really show
it! That's going to give a more energetic photon
(an ultraviolet photon) but not as energetic as 'this' ultraviolet
photon, over here. So, suffice it to say, that it's possible to
do it in 'multiple steps', as well... if you are more than one
excited level above the ground state! Okay, let's write some of this down, in
more formal language! Excitation: the act of moving an electron to
a higher energy level... because you just took energy
'out of thin air'.... light that was passing through you.
you have stolen somebody's energy (the light energy) to
allow yourself (the atom) to be at a higher energy state, So, that's
gonna yield an 'absorption line'... that color that
you just 'stole' is no longer evident in the spectrum that you might be
looking at. The two methods of excitation that we're
going to talk about are... It's possible (in a very dense... well it
doesn't have to be 'super' dense, but in a dense setting) like a liquid or a gas that's at room temperature on Earth...
that collisions of atoms happen all the time. It's possible for them to bump
into each other and transfer energy, that way and get their electrons to excite. But in interstellar space, where the density is
very low.... there's only one atom and a few photons passing by. Then, it's
going to be when it absorbs a specific photon with
sufficient (or really, the exact) amount of energy needed
to boost it to a higher energy state. The reverse of that (de-excitation)
is when an electron is descending to lower energy levels.... either
the one 'right' below it, or it could be multiple jumps downward.
Because the atom has just lost energy, by the conservation of energy you
know what just happened? The atom gave off that corresponding
energy in the form of a photon.... in terms of a quanta of light.
It yields an emission-line. So, within the atom itself, we know it
gives off emission and absorption lines. The electron loses the energy, and
this exact energy is given off by the exact frequency of that photon. So, that's
why it corresponds to such an 'exacting' line
or exacting color on the spectrum. That's an exact amount of energy (the difference
between energy level 3 and energy level 5, for
instance). Now, you're saying to yourself 'well,
then why do all the atoms have different spectra?' That's because (believe it, or not) every
single element (every single item on the periodic table...
Hydrogen, Helium, Nitrogen, Oxygen, Carbon.... all have a unique
structure to their energy levels for the atom!
Which is hugely helpful! That means that all of them are going to look different
in the end, doesn't it? Now, ionization.... is the process of giving an atom so much
energy that you've given it enough to do just that!...
Boost it out of the atom all-together! To escape all-together. To get it beyond the
'nth' excited state. As you might imagine.... is a lot of energy!
Well, it is! But what happens to the energy states (they get closer and closer
to each other), so.... getting past the 'infinite' one,
if you will.... is actually not that hard, in an
environment where there's lots of energetic photons around! So, if your
electron is no longer bound to the atom after the transition takes place...
that's another way of saying it has been 'ionized'...
or ionization. The amount of energy that corresponds
to getting your atom ionized, yeah alleviating it of one of its electrons,
or more... is called the ionization energy. Okay, those are important terms up
there.... excitation... de-excitation... ionization and ionization energy.
If you're a 'flashcard kind of a person', these are all 'flashcard-worthy'
terms! Continuing.... clearly an absorption
spectrum, by Kirchoff's laws happen when you have a
cooler (intermediate) gas between a hot source
and the observer. Absorption spectra are created when
atoms, that's the intervening cooler gas, absorb
photons from the hot opaque source behind it....
of just 'the right' energy to get their electrons to
excite to higher energy levels. If you have a multi-electron atom
(which Hydrogen is not), right? Hydrogen was the simplest one!
But for Helium, Lithium, Boron, Beryllium, Carbon, etc...
You get much more complicated! spectra, meaning.... many more lines can show up because there's many more possibilities in terms of
transitioning those electrons within the atom, itself.
Down below, 'here'... we have ... on the left, in frame (a), you've got Helium,
right? With just two protons and two neutrons in the nucleus...
and two electrons that orbit it. It shows the
mean spacing between, let's say... the ground state of the electrons and the
nucleus as being about 0.05 nanometers... or about five Angstroms.
That's because Helium is bigger than Hydrogen... it's about five times its
size! On the right... here in frame (b), you've got
six electrons so it must be Carbon, right? Or, with six protons and six
neutrons in the nucleus. Clearly, as you get more and more 'complex' with
more and more electrons to deal with.... then you have many more permutations and possibilities of... electronic change. So, yeah...
we kind of like Hydrogen because of its simplicity! But they (atoms) get complex very, very quickly! In the Hydrogen atom's case... this is
one of the 'In Focus' in your current textbook....
It's a pretty complicated 'little' page! But, it's just
so 'helpful' for us astronomers, because ninety percent of the universe is
Hydrogen! Maybe not so much on Earth, but the
universe at large is! This shows the
corresponding energy levels of the Hydrogen atom, itself....
just with a single proton in the nucleus and one electron having these possible
differences in in orbital energy. This
equation 'over here' on the left.... please do not memorize that! I will never ask you
about it again, i promise! The 'nth' level (the first level
energy level is the ground state level) and the first excited state is the
'n' equals two-level. It says 'If you would like to know the energy that
the electron has in the second energy level.... you could plug
'two' into the equation here. One over four. (one over two squared)
and it would tell you how many electron volts that's what 'eV'
is 'for'... that the electron actually has. So, this is the only 'almost' perfectly
understood atom - Hydrogen, because of its simplicity! But over here... on the right
(it kind of gets lost) it shows the Balmer series. In the Balmer series (the visible light
spectral lines from Hydrogen)... remember? there was a red line, a teal, and a violet line. We call it Hydrogen-
alpha, Hydrogen-beta, and Hydrogen- gamma... The Hydrogen-alpha line occurs when the
electron cascades from level three down to level two.
Or, from the second excited state, rather to the first excited state.
Because of that exact amount of energy difference between those two
levels... we know that the energy of the photon
that's emitted is of an 'exacting' color... an exacting frequency.
Likewise, when the electron cascades from the third-excited state down to the
first-excited state... it gives us the Beta line (or, that's the
teal line). So, the Balmer series is when electrons cascade from higher energy
levels down to the first-excited state. That's just
the 'sweet spot' for giving us 'visible' light photons. Bigger jumps
would give you even more energetic light. It would show up in the ultraviolet, for
instance! Less energetic light.... smaller jumps
between electronic energy states... would give you less
energetic light, like infrared. Now, we understand (we didn't,
originally). Balmer didn't understand what's going on with the atom either, but
he noticed these very 'orderly' pattern of lines in the Hydrogen spectrum.
We're going to talk a lot about the Balmer series!
This says that the emission OR the absorption lines
arising from transitions (changes in) electronic energies between the second
energy level (that's the first-excited state)
and 'all others' except the one below it (not the ground state)
So from upper levels down to level 'two' or from two
up to 'upper levels' - either of those!... give rise to what we call the Balmer
series of lines, which are the visible wavelength lines. Red, is for the simplest
transition from level three to level two, in this diagram here. It
shows that a photon 'came in' of just this 'red' color.
That excited the electron from energy level two to energy level three.
But, if you leave it (atom) alone for just a little bit, then it comes cascading back
down again... trying to seek out its lowest energy state. It gives off the exact same color of red... 'here'.
This diagram has the energy levels correct! Notice that between two and
three, it's a pretty big gap! Between three and four... it's much less.
Four and five.... smaller, still. Five and six, smaller still, yeah? They get
'bunched up' as you get to higher and higher
energy states.... but they're quantized! meaning...
They can only be at very exacting energy levels! The Hydrogen-specific series.... the Lyman
series of lines... are the lines that occur when
transitions occur (this is showing the energy state. Energy level one is the
ground state) and to go from the ground state up to a
higher energy state requires a big boost in energy!
It's going to be ultraviolet photons that are possible for
causing these transitions. Very specific wavelengths of ultraviolet light!
Those transistors are too big to give off visible light.
In the middle case, the Balmer series (which is so famous
because it was in the visible part of the spectrum) which we were all capable of
seeing long before any other parts of the spectrum... are the
transitions that occur from higher energy states down to level two (the
first excited state) or from level two up to higher energy
states. This is the one we'll keep talking
about because it's visible light and
that's what we're able to observe clearly on Earth, so easily! Just for trivia purposes...
the last one here is called the Paschen series.... with the 'sch' in it. That's for
smaller transitional levels that are caused by less energetic photons.