Atoms are the "building blocks of matter." Anything that has mass and occupies space (by having volume) is made up of these teeny tiny little units. That goes for the air you breathe, the water you drink and your body itself.
Isotopes are a vital concept in the study of atoms. Chemists, physicists and geologists use them to make sense of our world. But before we can explain what isotopes are — or why they're so important — we'll need to take a step back and look at atoms as a whole.
Our Atomic World
As you probably know, atoms have three main components — two of which reside in the nucleus. Located at the center of the atom, the nucleus is a tightly packed cluster of particles. Some of those particles are protons, which have positive electrical charges.
It's well-documented that opposite charges attract. Meanwhile, similarly charged bodies tend to repel one another. So here's a question: How can two or more protons — with their positive charges — coexist in the same nucleus? Shouldn't they be pushing each other away?
That's where the neutrons come in. Neutrons are subatomic particles that share nuclei with protons. But neutrons don't possess an electrical charge. True to their name, neutrons are neutral, being neither positively nor negatively charged. It's an important attribute. By virtue of their neutrality, neutrons can stop protons from driving one another clear out of the nucleus.
"Elementary, My Dear Watson"
Orbiting the nucleus are the electrons, ultra-light particles with negative charges. Electrons facilitate chemical bonding — and their movements can produce a little thing called electricity. Protons are no less important. For one thing, they help scientists tell the elements apart.
You might have noticed that in most versions of the periodic table, each square has a little number printed in its upper righthand corner above the element symbol. That figure is known as the atomic number. It tells the reader how many protons are in the atomic nucleus of a particular element. For example, oxygen's atomic number is eight. Every oxygen atom in the universe has a nucleus with exactly eight protons; no more, no less.
Without this very specific arrangement of particles, oxygen wouldn't be oxygen. Each element's atomic number — including oxygen's — is totally unique. No two elements can have the same atomic number. No other element has eight protons per nucleus. By counting the number of protons, you can identify an atom. Just as oxygen atoms will always have eight protons, nitrogen atoms invariably come with seven. It's that simple.
Neutrons do not follow suit. The nucleus in an oxygen atom is guaranteed to harbor eight protons (as we've established). However, it might also contain anywhere from four to 20 neutrons. Isotopes are variants of the same element that have different numbers of neutrons (and thus potentially different physical properties). They do, however, tend to have the same chemical properties.
Now, each isotope is named on the basis of its mass number, which is the total combined number of neutrons and protons in an atom. For example, one of the better-known oxygen isotopes is called oxygen-18 (O-18). It's got the standard eight protons plus 10 neutrons.
Ergo, the mass number of O-18 is — you guessed it — 18. A related isotope, oxygen-17 (O-17), has one fewer neutron in the nucleus. O-16, then, has the same number of protons and neutrons: eight. Among this trio, O-16 and O-17 are the lighter isotopes, and O-16 is also the most abundant isotope of the three.
Some combinations are stronger than others. Scientists classify O-16, O-17 and O-18 as stable isotopes. In a stable isotope, the forces exerted by the protons and neutrons hold each other together, permanently keeping the nucleus intact.
On the flip side, the nucleus in a radioactive isotope, also called a "radioisotope," is unstable and will decay over time. A radioactive isotope has a proton-to-neutron ratio that's fundamentally unsustainable in the long run. Nobody wants to stay in that predicament. Hence, radioactive isotopes will shed certain subatomic particles (and release energy) until they've converted themselves into nice, stable isotopes.
O-18 is stable, but oxygen-19 (O-19) is not. The latter will inevitably break down — fast! Within 26.88 seconds of its creation, a sample of O-19 is guaranteed to lose half of its atoms to the ravages of radioactive decay.
That means O-19 has a half-life of 26.88 seconds. A half-life is the amount of time it takes 50 percent of an isotope sample to decay. Remember this concept; we're going to connect it to paleontology in the next section.
But before we talk fossil science, there's an important point that needs to be made. Unlike oxygen, some elements do not have any stable isotopes whatsoever. Consider uranium, one of the most well-known radioactive elements. In the natural world there are three isotopes of this heavy metal, and they're all radioactive, with the atomic nuclei in a constant state of decay. Eventually, a chunk of uranium will turn into an altogether different element on the periodic table.
Don't bother trying to watch the transition in real time. The process unfolds very, very slowly.
Getting Dates (And Staying Healthy)
Uranium-238 (U-238), the element's most common isotope, has a half-life of about 4.5 billion years! Gradually, this will become lead-206 (Pb-206), which is stable. Likewise, uranium-235 (U-235) — with its 704 million-year half-life — transitions into lead-207 (Pb-207), another stable isotope. (Both U-238 and U-235 are examples of naturally occurring isotopes.)
To geologists, this is really useful information. Let's say somebody finds a slab of rock whose zircon crystals contain a mixture of U-235 and Pb-207. The ratio of these two atoms can help scientists determine the rock's age.
Here's how: Let's say the lead atoms vastly outnumber their uranium counterparts. In that case, you know you're looking at a pretty old rock. After all, the uranium's had plenty of time to start transforming itself into lead. On the other hand, if the opposite is true — and the uranium atoms are more common — then the rock must be on the younger side.
The technique we've just described is called radiometric dating. That's the act of using the well-documented decay rates of unstable isotopes to estimate the age of rock samples and geologic formations. Paleontologists harness this strategy to determine how much time has elapsed since a particular fossil was deposited. (Although it's not always possible to date the specimen directly.)
You don't need to be a prehistory buff to appreciate isotopes. Medical practitioners use some of the radioactive varieties to monitor blood flow, study bone growth and even fight cancer. Radioisotopes have also been used to give farmers insights into soil quality.
So there you have it. Something as seemingly abstract as the variability of neutrons affects everything from cancer treatment to the mysteries of deep time. Science is awesome.
Originally Published: Jun 17, 2019