by Bob Hamers
Millions of people across the US recently got their first view of the Aurora Borealis, or “Northern Lights”, as the night sky turned a rainbow of unusual colors. When I mentioned the aurora in my last blog post about astrophotography for Valentine’s Day, little did I know that this summer would be such an amazing opportunity to see this colorful phenomenon in action. In recent months, those of us in Madison, WI have been treated to several beautiful displays of green, red, and purple auroras; and I was able to capture some of them on camera:
(Photos by Bob Hamers, using an iPhone 11)
But where do these beautiful colors come from?
Auroras are caused by charged particles ejected from the sun that travel to earth and interact with our atmosphere. This summer we were at the point in the sun’s 11-year cycle when its magnetic field is most unstable; forming sunspots, flares, and coronal mass ejections when enormous amounts of highly energetic charged particles are hurled into space. When this occurs and the particles head toward the earth, they get caught in our magnetic field and spiral around, concentrating near the magnetic north pole (which is in northern Canada) and magnetic south pole. The auroras form as these charged particles crash into gas molecules in the upper fringes of the earth’s atmosphere – about 100 km (62 miles) above the earth’s surface. And this is where chemistry comes in!
Our atmosphere has a rather constant composition of 78% nitrogen and 21% molecular oxygen (O2) up until 100 km (62 miles) above mean sea level, an altitude known as the Karman Line. At altitudes higher than the Karman Line, the composition suddenly changes: N2 and O2 decrease, and atomic oxygen (O) becomes the main component:
(Figure adapted from Wikipedia image by Amaurea)
To understand why Auroras form at this height, we need to talk about electron orbitals and spin. Electrons in an atom can have different states or “levels” of different energies. For example, the three most important configurations of the four key electrons in an oxygen atom are shown here as up and down arrows arranged in three different energy levels:
Spectroscopists label the three levels as 1S,1D, and 3P. The meaning of these designations isn’t important here, but each arrangement of electrons has its own unique energy. When electrons change from one configuration to another, they have to absorb or emit energy so that the total energy is conserved. (You may remember this phenomenon from Cathy Murphy’s post about “Electrocuting a Pickle”, as well as my Valentine’s Day post.)
Under normal conditions, most O atoms are at the lowest-energy (most-stable) 3P configuration. But when those charged particles from the sun ionize nitrogen molecules, which then collide with the oxygen in the atmosphere, some of these O atoms are excited up to the high-energy 1S configuration level by various collisions. After these O electrons get excited up to the highly energetic 1S configuration, you’d think they would release their energy (as light!) and jump down to the 1D or 3P configuration. Instead they get stuck. For complicated reasons having to do with the laws of quantum mechanics, the electrons in each of these configurations can’t easily jump between these different configurations. Spectroscopists call these “forbidden” transitions. But “forbidden” isn’t truly forbidden: these transitions are really just kinetically unfavorable. If the atmospheric pressure is low enough that there aren’t many other molecules around to interact with, then after some time (a few tenths of a second), the electrons in these 1S atoms will jump down in energy anyway, even though it’s supposedly forbidden, and give off light. And where do we have low pressure with molecules not bumping into each other very much? High up in the atmosphere! When these electrons jump from the 1S to 1D configuration they emit light at 557 nm, which is green. So, we have a green aurora.
(Photo by Bob Hamers)
Once in the 1D configuration, the electrons still have a lot of energy, but they get stuck again—and stuck even worse than before, because now in order to drop down to the 3P configuration they also have to change the spin of one of the electrons (indicated by the arrows on the diagram). Once again, even though it’s unfavorable, if there is no other way to get rid of its remaining energy, the electron will eventually drop from 1D to the lowest-energy level (3P) configuration. The energy is emitted as light at 630 nm, which is red.
But if both red and green are created by oxygen atoms, why do we see red and green coming from different locations in the sky? At altitudes lower than ~100 km above mean sea level (where the pressures are higher), the O atoms get excited, but collisions with other gas molecules jostle the electrons and let them drop to lower configurations without emitting light. When this happens, we say the emission is quenched.
Ultimately the aurora depends on a subtle balance of having enough collisions with N2+ and other species to excite oxygen electrons up to the 1S configuration, but few enough collisions so that the O atoms emit light instead of being quenched.
This balance is reached at altitudes near the Karman Line, 100 km (62 miles) straight up. Green light can be emitted from altitudes of ~ 100 km above mean sea level, but red light is produced at even higher altitudes because the red 1D-3P emission is more easily quenched by collisions with other gas molecules at lower altitude. So, green light is emitted from lower down in the atmosphere, while red light is only emitted from higher up.
The auroras sometimes show other colors too, due to emission from nitrogen and other gas molecules. Green and red are the most common, but nitrogen (N2+) in particular can give rise to a beautiful purple color from very high altitudes. Auroras can give off many different colors, and are surely one of Nature’s great spectroscopic spectacles!
