Do you ever wonder where light comes from or how it is produced? In this lesson, we are going to use our knowledge of the electron configurations and quantum numbers to see what goes on during the creation of light.
The Bohr Model and Atomic Spectra
Imagine it is a holiday, and you are outside at night enjoying a beautiful display of fireworks. They are exploding in all kinds of bright colors: red, green, blue, yellow and white. Later on, you're walking home and pass an advertising sign. One of the bulbs is emitting a blue light and the other has a bright red glow. What produces all of these different colors of lights? What is responsible for this? The answer is electrons. Did you know that it is the electronic structure of the atoms that causes these different colors to be produced?
The rainbow colors are listed in order of increasing energy.
The Hydrogen Atom
We're going to start off this lesson by focusing on just the hydrogen atom because it's a simple atom with a very simple electronic structure. It only has one electron which is located in the 1s orbital. Recall from a previous lesson that 1s means it has a principal quantum number of 1. This means it's in the first and lowest energy level, and because it is in an s orbital, it will be found in a region that is shaped like a sphere surrounding the nucleus.
All we are going to focus on in this lesson is the energy level, or the 1 (sometimes written as n=1). This little electron is located in the lowest energy level, called the ground state, meaning that it has the lowest energy possible. When you write electron configurations for atoms, you are writing them in their ground state.
So, if this electron is now found in the ground state, can it be found in another state? Absolutely. If this electron gets excited, it can move up to the second, third or even a higher energy level. But what causes this electron to get excited? Adding energy to an electron will cause it to get excited and move out to a higher energy level. The more energy that is added to the atom, the farther out the electron will go. In a later lesson, we'll discuss what happens to the electron if too much energy is added.
A couple of ways that energy can be added to an electron is in the form of heat, in the case of fireworks, or electricity, in the case of neon lights. When these forms of energy are added to atoms, their electrons take that energy and use it to move out to outer energy levels farther away from the nucleus.
The Rutherford model shows electrons as being situated randomly around the nucleus.
Now, those electrons can't stay away from the nucleus in those high energy levels forever. Eventually, the electrons will fall back down to lower energy levels. If the electrons are going from a high-energy state to a low-energy state, where is all this extra energy going? Energy doesn't just disappear.
As electrons transition from a high-energy orbital to a low-energy orbital, the difference in energy is released from the atom in the form of a photon. A photon is a weightless particle of electromagnetic radiation. Electromagnetic radiation comes in many forms: heat, light, ultraviolet light and x-rays are just a few. The most important feature of this photon is that the larger the transition the electron makes to produce it, the higher the energy the photon will have.
This is where things get interesting.
High-energy photons are going to look like higher-energy colors: purple, blue and green, whereas lower-energy photons are going to be seen as lower-energy colors like red, orange and yellow. Remember those colors of the rainbow - red, orange, yellow, green, blue and violet? Those are listed in the order of increasing energy. Essentially, each transition that this hydrogen electron makes will correspond to a different amount of energy and a different color that is being released. This is called its atomic spectrum.
The Bohr Model
The Bohr model shows electrons situated in specific energy levels.
So, who discovered this? In the early 1900s, a guy named Niels Bohr was doing research on the atom and was picturing the Rutherford model of the atom, which - you may recall - depicts the atom as having a small, positively-charged nucleus in the center surrounded by a kind of randomly-situated group of electrons. While Bohr was doing research on the structure of the atom, he discovered that as the hydrogen atoms were getting excited and then releasing energy, only three different colors of visible light were being emitted: red, bluish-green and violet. If the electrons were randomly situated, as he initially believed based upon the experiments of Rutherford, then they would be able to absorb and release energy of random colors of light. His conclusion was that electrons are not randomly situated. Instead, they are located in very specific locations that we now call energy levels.
This led to the Bohr model of the atom, in which a small, positive nucleus is surrounded by electrons located in very specific energy levels. For example, whenever a hydrogen electron drops from the fifth energy level to the second energy level, it always gives off a violet light with a wavelength of 434.1 nanometers. A wavelength is just a numerical way of measuring the color of light. Also, whenever a hydrogen electron dropped only from the third energy level to the second energy level, it gave off a very low-energy red light with a wavelength of 656.3 nanometers.
These findings were so significant that the idea of the atom changed completely. What was once thought of as an almost random distribution of electrons became the idea that electrons only have specific locations where they can be found. This is where the idea of electron configurations and quantum numbers began.
Many Electron Atoms
We now know that when the hydrogen electrons get excited, they're going to emit very specific colors depending on the amount of energy that is lost by each. This also happens in elements with atoms that have multiple electrons. However, because each element has a different electron configuration and a slightly different structure, the colors that are given off by each element are going to be different. Each element is going to have its own distinct color when its electrons are excited - or its own atomic spectrum. That's what causes different colors of fireworks! For example, when copper is burned, it produces a bluish-greenish flame. When sodium is burned, it produces a yellowish-golden flame. When magnesium is burned, it releases photons that are so high in energy that it goes higher than violet and emits an ultraviolet flame. You wouldn't want to look directly at that one!
Each element has its own atomic spectrum.
When neon lights are energized with electricity, each element will also produce a different color of light. In fact, the term 'neon' light is just referring to the red lights. Blue lights are produced by electrified argon, and orange lights are really produced by electrified helium. These atomic spectra are almost like elements' fingerprints. The color a substance emits when its electrons get excited can be used to help identify which elements are present in a given sample. Scientists use these atomic spectra to determine which elements are burning on stars in the distant outer space.
I hope this lesson shed some light on what those little electrons are responsible for! To me, it is one of the most interesting aspects of the atom, and when it comes down to the source of light, it's really just a simple process. First, energy is absorbed by the atom in the form of heat, light, electricity, etc. Second, electrons move out to higher energy levels. They get excited. Third, electrons fall back down to lower energy levels. They can't stay excited forever! Finally, energy is released from the atom in the form of a photon.
After watching this lesson, you should be able to:
- Define ground state, photon, electromagnetic radiation and atomic spectrum
- Summarize the Bohr model and differentiate it from the Rutherford model
- Explain how electrons emit light and how they can emit different colors of light