By Ron B. Davis Jr., Georgetown University
It might at first seem that oxygen and nitrogen appear next to one another on the table but with nitrogen at element seven, and oxygen at element eight, these two substances differ at the atomic level only by one proton/electron pair. The periodic table helps us understand why these two elements are so common in our atmosphere, and yet so different in their chemical properties.
Why Are Oxygen and Nitrogen So Abundant?
In some ways, it isn’t hard to see why oxygen and nitrogen together make up about 99% of the air we breathe. Nitrogen is element number seven. Like hydrogen and helium, nitrogen is one of the few elements whose atomic number also tells you how abundant it is, since nitrogen is the seventh most common element in the universe. Oxygen is even more abundant, coming in as the third most abundant element in the universe. That’s right: it’s hydrogen, helium, then… element number 8 oxygen is third.
There are two reasons for this extra abundance of oxygen. First, oxygen has an even number of protons. So the Oddo-Harkins rule predicts that it should be more abundant than nitrogen. Second, while two is the magic number that makes helium so stable, the next magic number for nuclear stability after two is eight.
This means that element number eight, oxygen contains a magic number of protons. What’s more, its most common isotope, oxygen-16, has a magic number of neutrons as well, making it one of those rare nuclei that is doubly-magic and therefore doubly stable.
Diatomic Oxygen, Diatomic Fluorine, and Diatomic Nitrogen
We find oxygen and nitrogen diatomic molecules in high abundance in our atmosphere. Oxygen comes from group 16 of the table, its atoms are two electrons shy of an octet. This means that when two atoms of oxygen join to make a molecule, they join by forming a double bond, sharing two pairs of electrons—in other words, a total of four electrons—to achieve an octet. Fluorine atoms, by contrast, only share two electrons—a single pair—so they form only a single bond that is much weaker.
Scientists can approximate a molecule’s stability by envisioning the energy required to pull apart two bonded atoms into their constituent, neutral atoms. This energy to pull two bonded atoms apart is called the average molar bond enthalpy. It only takes 155 kJ of energy to pull apart one mole of fluorine molecules. By contrast, pulling apart the double bond shared by two oxygen atoms takes 495 kJ per mole—more than three times as much.
Nitrogen comes from group 15, meaning that its atoms are three electrons shy of an octet. So diatomic nitrogen is expected to contain six shared electrons, bonding in three pairs, for a triple bond. Nitrogen’s triple bond is vastly more stable than oxygen’s double bond, requiring 941 kJ on average to separate a mole of diatomic nitrogen into atoms. That’s more than six times fluorine’s single bond.
This extreme stability of molecular nitrogen makes nitrogen unable to support many of the kinds of chemical reactions that oxygen can—reactions like combustion and respiration that help to power our homes, or vehicles and even ourselves.
This article comes directly from content in the video series Understanding the Periodic Table. Watch it now, on Wondrium.
Shared Physical Characteristics
Fluorine, oxygen and nitrogen do share many physical properties. Their shared region on the right side of the second period makes for small diatomic molecules that strongly prefer the gas phase. All three become gases at very low boiling points that are strikingly similar. And all three liquify only at temperatures close to those found on the dark side of the Moon.
Gases liquified at very low temperatures are called cryogens, and liquid nitrogen is one of the most widely used cryogens in the world today. Liquid oxygen achieves roughly the same temperature, but oxygen’s reactivity gives its liquid phase a much different use.
Engineers at NASA have relied heavily on the reactive power of liquid oxygen, using huge tanks of liquid oxygen to enhance the combustion of common fuels like hydrogen or natural gas in booster rockets that provide enough thrust to take vehicles to the edge of space. As the liquid oxygen is released from its tank, it instantly vaporizes and creates highly concentrated oxygen gas that accelerates the combustion of fuels.
The Reason for the Balanced Amount of Reactivity in the Atmosphere
But the reactivity of oxygen can be a double-edged sword. NASA engineers learned this to their horror, during a ground rehearsal test with the Apollo 1 crew. To make breathing easier, as well as reduce outward pressure on the hull in the vacuum of space, their command module was filled with pure oxygen. Unfortunately, there was a momentary surge of voltage in one of the module’s components, and that was enough to cause a fire. All three astronauts died within seconds.
This tragedy illustrates the importance of having nitrogen and oxygen working together in our atmosphere. Later Apollo missions launched using a mixture of oxygen with nitrogen to reduce the risk of fire. Similarly, the oxygen in earth’s atmosphere is essentially diluted with unreactive nitrogen, making our resulting atmosphere far less capable of supporting fires.
So in that sense it’s fortunate that these two elements have such differing reactivity. The two of them work in concert to offer a balanced amount of reactivity in our atmosphere, allowing air to support life, without putting us one lightning strike away from spontaneously setting the entire planet on fire!
Common Questions about the Abundance of Oxygen and Nitrogen
After hydrogen and helium, oxygen takes the third place among the other most abundant elements, and element number seven which is nitrogen is the seventh in this regard.
These elements both appear in the same region of the periodic table, making them prefer the gas phase. They liquify at a very low temperature similar to the one at the dark side of the moon and they turn to gas at a strikingly similar boiling point.
Oxygen and nitrogen work together in the atmosphere to keep the level of reactivity completely balanced. Pure oxygen in the Earth’s atmosphere, by itself, is very reactive and can set our planet on fire. To keep this from happening, unreactive nitrogen is mixed with the oxygen, making it less capable of supporting fire.
