38.4: Isotopes

Isotopes are atoms of a given element that vary in their weight. Atoms are made of protons, neutrons, and electrons. The number of protons in an atom determine its identity as an element. For instance, all oxygen atoms have eight protons. All carbon atoms have six protons. Usually, the number of protons (which carry a positive charge) is matched by the number of electrons (which carry a negative charge). This balance makes the atom electrically neutral. If there are more protons than neutrons, the atom is a cation. If there are more electrons than protons, it makes it an anion. We’ll ignore ions for the present discussion.

It’s the number of neutrons that matters for determining isotopic identity. In addition to its six protons, a given atom of carbon can contain six neutrons, seven neutrons, or eight neutrons. When you add the number or protons to the number of neutrons, you get the atomic mass of that atom.

6 + 6 = 12

6 + 7 = 13

6 + 8 = 14

So that means carbon atoms come in three isotopic varieties: \(\ce{^{12}C}\), \(\ce{^{13}C}\), and \(\ce{^{14}C}\). You’ve probably heard of \(\ce{^{14}C}\) (carbon-14) as the basis of dating recent organic remains, such as wood, bone, or tissue. It is radioactive, meaning the atom is unstable. Eight is one neutron too many, and the forces that hold the atom together are incapable of maintaining its structure over the geological long-term. As a result, it’s not useful to us for the current discussion. It’s also a tiny amount of all carbon, about one part in a thousand. So let’s set \(\ce{^{14}C}\) aside and focus now on \(\ce{^{12}C}\) and \(\ce{^{13}C}\) instead.

\(\ce{^{12}C}\) and \(\ce{^{13}C}\) are not radioactive; the forces that hold their atoms together are durable over geological spans of time. Hence, we call them “stable” isotopes. \(\ce{^{12}C}\) is “light carbon.” \(\ce{^{13}C}\) is “heavy carbon.” Most carbon atoms are the lightweight kind, \(\ce{^{12}C}\). It makes up about 99% of all carbon atoms. But there is also a measurable amount of heavy \(\ce{^{13}C}\), comprising about 1% of all carbon on Earth.

It turns out that various natural processes separate light carbon from heavy carbon. These acts of separation, called isotope fractionation, enrich some parts of the Earth system in \(\ce{^{12}C}\), while depleting other parts. For instance, photosynthesizing plants prefer \(\ce{^{12}C}\)-based \(\ce{CO2}\) molecules, meaning that plants (and the animals that eat the plants) are enriched in \(\ce{^{12}C}\), but that act of removing \(\ce{^{12}C}\) from the atmosphere makes it relatively depleted in \(\ce{^{12}C}\). Or if you consider it from the perspective of the less-common \(\ce{^{13}C}\) atoms, then it’s the other way around: life is the reservoir that’s depleted in heavy carbon, and the atmosphere is enriched in heavy carbon. The ratio between these two isotopes is something we use as an indicator of past conditions on Earth, such as its climate. The ratio of \(\ce{^{12}C}\) and \(\ce{^{13}C}\) gets skewed between the major reservoirs in the Earth system as climate shifts. The same is true of heavy oxygen and light oxygen (\(\ce{^{16}O}\) vs. \(\ce{^{18}O}\)), or light hydrogen and heavy hydrogen (\(\ce{^1H}\) vs \(\ce{^2H}\), the latter of which is also called D for “deuterium”). Critically, with all these isotopes, we can measure that skew.

This video discusses three different elements (hydrogen, oxygen, and carbon) in terms of natural fractionation of their isotopes, and relates that fractionation to Earth’s prevailing climate:

Because plants prefer \(\ce{^{12}C}\) over \(\ce{^{13}C}\), during times when there are high rates of carbon burial happening on Earth, much of the \(\ce{^{12}C}\) ends up underground, being converted to coal, oil, or natural gas. That means the atmosphere will be relatively enriched in \(\ce{^{13}C}\) instead. Because the ocean exchanges gases with the atmosphere, the waters of the ocean will also be enriched in \(\ce{^{13}C}\) instead of \(\ce{^{12}C}\).

When climate is cool, there is less energy available to drive evaporation, so ocean water tends to be enriched in heavier isotopes: water molecules containing deuterium or \(\ce{^{18}O}\). Evaporated water molecules are more likely to be isotopically lighter, and so will the precipitation that they drop. During ice ages, the snow that falls to make glacial ice will be a reservoir of isotopically lightweight water, both in terms of hydrogen and oxygen.