Back To CourseBasics of Astronomy
28 chapters | 325 lessons
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What does a staircase have to do with atoms and electrons? Well, the obvious answer is it is made of atoms, which contain electrons. But enough of that, smarty-pants. It can also be used as a good metaphor for this lesson's concepts involving atoms, electrons, and transitions. You'll see in a bit what I mean by that.
The transition, or the movement, of an electron between energy levels, in an atom can occur in more than one way. For an electron to transition to a higher energy level, it must absorb energy, just like it takes energy to lift a rocket upwards into the sky or to lift a heavy weight above your head.
If an atom collides with another atom, ion, or electron, the atom can become excited. An excited atom is an atom where an electron has moved from a lower to a higher energy level. You know how when two football players forcefully collide it looks like the helmet jumps up off of their head? Well, a collision can also provide enough energy to get an electron to jump up off of a lower energy level and into a higher energy level.
Another way an atom can become excited is by absorbing a photon, a small bundle of electromagnetic radiation. Electromagnetic radiation, or light, is a form of energy that has wave-like properties. Therefore, the energy of a photon depends on its wavelength; the shorter the wavelength, the higher the energy.
Only photons of specific wavelengths can be absorbed by an atom. Basically, the photon is like an energy drink for a little electron, giving it more buzz to jump up higher into a higher energy level.
But just like people may have a preference for one energy drink to another, atoms have preferences for the kinds of photons they can absorb. Only photons with energies equal to the energy difference between two energy levels in a specific atom can be absorbed. This means electrons can occupy only certain permitted energy levels; there are no energy levels in between.
Consequently, when just the right kind of photon hits an atom, it's absorbed, and the electron jumps to a higher energy level. Since atoms have many energy levels, depending on how energetic this photon is, the electron can jump up one or more energy levels. All of this also means that more than one wavelength of light can be absorbed by a single atom to get an electron to jump into a higher energy level.
If a continuous spectrum of photons (a complete arrangement of colors) shines on a group of identical atoms, these atoms, like sponges, will understandably absorb only certain kinds of photons from the continuous spectrum. When this happens, an absorption-line spectrum will be produced. Because an absorbed wavelength of light removes a color from the original continuous spectrum, the resulting absorption spectrum is also called a dark-line spectrum.
Excited atoms cannot stay excited for long, however, and so the electron must eventually jump down to a lower energy level. As it does so, the electron emits a photon with energy (and thus wavelength) equal to the difference in energy levels between the two levels the electron jumps in between. The photons that are emitted in such a fashion make bright colorful lines against a dark background. These are known as bright-line or emission-line spectra.
Let me try and put all of the confusing core concepts of this lesson into a more simple metaphor. Let's pretend you're an electron. A photon of a specific energy (or wavelength) can be like a specific energy drink. The energy levels can be like steps in a staircase in your home.
You are now standing at the bottom step, the lowest possible energy level in the atom. You have a few energy drinks with different strengths next to you. You know that to jump from the bottom step up, you need energy. A little bit of energy to jump to the second step but a lot more energy to jump from the bottom all the way up to the third step in one fell swoop.
Okay, now you take a sip of the first energy drink. You don't move. Why not? It's because that drink didn't provide just the right amount of energy for you to transition between two steps. You can only jump onto a fully-fledged step. You can't jump to a fourth or a half of a step; such a thing doesn't exist on the staircase.
So, if the drink doesn't give you exactly the right amount of energy to jump onto a solid step, you're not going to jump onto anything, are you? Of course not! You'll crash and burn if you do that.
Okay, now you take a sip of a second energy drink. Boom! You jump to the third step. Nicely done. That drink gave you just the right amount of energy to do so. That means a photon of a specific wavelength was absorbed.
But no one really likes to just stand on a step on a staircase for long. So, you jump down a step, to a lower energy level. As you land from your jump, you emit a photon.
Transitions like this that occur in the hydrogen atom, the most abundant atom in the universe, can be grouped into well-known series, including the Lyman series, Balmer series, and Paschen series.
The arrows pointing up on the image on your screen represent absorption of energy.
The arrows pointing down represent emission of energy. Longer arrows represent larger amounts of energy than shorter arrows. This means that longer arrows represent shorter-wavelength photons.
Knowing this, let's take a look at the Lyman series, which starts at the atom's ground state (or energy level 1). They are long arrows and represent the shorter-wavelength ultraviolet part of the spectrum, one that is invisible to the human eye.
The Balmer series are the shorter arrows starting at the second energy level, representing lower energies and longer wavelengths. The three longest wavelength Balmer lines are visible to the human eye.
Even shorter than the Balmer series are the Paschen series, starting at level 3. They have lower energies than the Balmer series and thus even longer wavelengths. Hence, they fall into the longer wavelength, and invisible to the human eye, infrared part of the spectrum.
The transition, or the movement, of an electron between energy levels can occur thanks to the absorption or emission of a photon, a small bundle of electromagnetic radiation. Electromagnetic radiation, or light, is a form of energy that has wave-like properties.
The energy of a photon depends on its wavelength; the shorter the wavelength, the higher the energy. Remember, only photons of specific wavelengths can be absorbed by an atom. Thus, if a continuous spectrum of photons shines upon a group of the same atoms, they will only absorb specific kinds of photons. This will remove these photons from the continuous spectrum, resulting in an absorption spectrum, which has dark lines representing the wavelengths that were absorbed.
Since excited atoms can't be excited for long, the electrons jump down to a lower energy level, and as this occurs, the electron emits a photon with energy (and thus wavelength) equal to the difference in energy levels between the two levels the electron jumps in between. This means an emission-line spectrum is produced, one that has bright lines representing the wavelength emitted, juxtaposed against a dark background.
Transitions in the hydrogen atom can be grouped into many series, including the Lyman series, which represents invisible ultraviolet wavelengths; the Balmer series, which includes some wavelengths of visible light; and the Paschen series, representing infrared wavelengths.
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Back To CourseBasics of Astronomy
28 chapters | 325 lessons