Today, we want to examine the physics of why this is in terms of the structure of atoms.
Here is an overview of this subject that you can refer to while we are going alone.
The entire key to understanding all of this is the recognition that each atom has a unique set of atomic energy levels . This will become clearer as we run the simulations and make some measurements. We will also have a transmitter quiz at the end of all this to see, collectively, how well we are groking the concept.
The basic ideas of atomic energy levels and transitions will become clear when we run the simulation. For now, the details are summarized in the diagrams below.
We begin with the simple concept of Energy levels in an atom as shown below
The key point is that atomic energy levels or discrete (or quantised). These levels are occuppied by electrons. Therefore, in any given atom, an electron can only be in a certain, discrete energy state. Each element has a unique set of energy levels.
The concept of discrete is important. A photon is either all or nothing. You can not take a photon, steal part of its energy, and still have that photon left. So, in general, there is a one to one correspondence between photon generation or absorption and the movement of electrons to different energy levels.
Photon Emission:
If the electron in an upper energy level (say E3) falls to
a lower energy level (say E2) one photon is emitted
whose energy is exactly equal to the energy difference between levels
E3 and E2.
For photons, their energy is inversely proportional to their wavelength. Short wavelength photons (e.g. UV photons) have more energy than long wavelength photons (e.g. IR photons). In the optical, blue photons have shorter wavelength (around 4000 angstroms) than red photons (around 7000 angstroms) and therefore have more energy.
All photons travel at the speed of light, independent of their energy or wavelength. |
What the observer therefore sees is photon emission at a unique wavelength and that is what the spectra of elements are.
Its corresponding emission spectrum
Photon Absorption:
Now, suppose there is an electron that's in the lowest energy state ( E1 in the diagram above). This lowest energy state is called the ground state.
This electron can be moved to an excited state by absorbing a photon in the following manner.
Suppose the energy difference between levels E1 and E2 is 4 units of energy. Therefore, if an incoming photon has precisely 4 units of energy, that photon can be "absorbed" by the atom, causing the electron to move from energy level E1 to energy level E2.
However, electrons do not like to stay excited and therefore this electron
would quickly transition back down to the ground state and therefore
emit a photon of energy = 4 units. To
the external observer, no net absorption would occur in this case.
Now let's say the energy difference between levels E1 and E3 is 6 units and further, that there is some rule for this atom (in real life called a transition probability), that tells the atom if an electron is in the excited state E3 it can not fall back directly to the ground state (thereby emitting a photon of energy 6 units) but must first fall back to energy state E2 (and therefore emit a photon of energy 2 units) and then go from E2 to E1 (and emit a photon of energy 4 units).
For this case then, an external observer would see no photons coming from the source with energy = 6 units. They have all been absorbed by the atom. This produces a lack of photons at a precise wavelength which is then the absorption line seen in the stellar spectrum.
Below is a spectrum of a star like our Sun. Next week we will be measuring features in real stellar spectra. The "dips" in the spectra are due to absorption lines, places in the spectra where photon energy has been removed due to the excitation of an electron from a lower atomic level to an upper one, in the manner described above.
So if we consider the following simplified 3 level atom we can see the mixture of emission and abosrption, each of which is uniquely associated with either the generation of a photon or the removal (absorption) of a photon of a specific energy/wavelength.
Finally note that if we give "too much" energy to the atom, the electron will be stripped away and there will be no electrons in a bound atomic energy state. This process is known as "ionization" and we will discuss this further later.
In the example above, suppose that E4 is 7 units of energy above the ground state and its the highest energy level in this atom. This means that if this atom receives more than 7 units of energy, it will become ionized.
Since photon energy is inversely proportional to wavelength and since hot stars radiate more of their energy at short wavelengths, then we expect ionization to occur (easily) in hot stars as most of the emitted photons have energies greater than the ionization energy of an atom.
Illustrative Summary of different types of spectra:
Now let's do some stuff related to discrete spectra and photon energies.
Note as a useful reference: optical photons are in the range of 3500 - 7500 angstroms (this is basically what your eye can see).
So, each group with a laptop should do the following:
For each chosen atom, run the simulation and then report your measurements using this tool
The exercise for you is to move the electron to an excited state in the simulation and then observe its behavior as it falls down to the ground state. Record the wavelength of the various photons that are emitted along with the energy level difference between the upper and lower energy states that correspond to the photon emission.
Each Chemical Element has a Unique set of atomic energy levels and therefore has a unique spectrum |