Atoms are the building blocks of matter. The screen you’re reading this on, the brain you’re reading with, they’re all very organized groups of atoms. They interact in specific ways, obeying specific rules, to maintain the shape and function of objects.
None of it works, however, unless the right atoms are involved. If you try to put the wrong ones into a protein or water molecule, it breaks apart. It’s like trying to cobble together a picture using pixels of the wrong colors.
Given how rigorous chemistry is on this, it’s surprising to see how much variety these ‘right’ atoms can get away with. Each element on the periodic table encompasses whole families of atoms who behave the same despite some important differences — isotopes.
What are isotopes?
Isotopes are atom families that have the same number of protons, but different numbers of neutrons. The term is drawn from ancient Greek words isos and topos, meaning ‘equal place’, to signify that they belong to the same elements on the periodic table.
Atoms are made of a dense core (nucleus) orbited by a swarm of electrons. The protons and neutrons that form the core represent virtually all of an atom’s mass and are largely identical except for their electrical charges — protons carry a positive charge, while neutrons don’t have any charge. The (negatively charged) electron envelope around the core dictates how atoms behave chemically.
The kicker here is that since neutrons carry no charge, they don’t need an electron nearby to balance them out. This renders their presence meaningless in most chemical processes.
To get a bit more technical, the number of protons within an atom’s nucleus is its ‘atomic number’ (aka the ‘proton number‘, usually notated ‘Z‘). Since protons are positively charged, each atom worth its salt will try to keep the same number of electrons in orbit to balance out its overall electric charge. If not, they’ll try to find other charge-impaired atoms and form ionic compounds, like literal salt, or covalent bonds — but that’s another story for another time.
What’s important right now is to keep in mind that these atomic numbers identify individual elements. The atomic number is roughly equivalent to an element’s numeric place in the periodic table, and in broad lines dictates how an element tends to behave. All isotopes of an element have the same atomic number. What they differ in is their ‘mass number‘ (usually abbreviated ‘A‘), which denotes the total number of protons and neutrons in an atom’s core.
In other words, isotopes are atoms of the same element — but some just weigh more.
For example, two isotopes of Uranium, U-235 and U-238, have the same atomic number (92), but mass numbers of 235 and 238, respectively. You can have two isotopes of the same mass, like C-14 and N-14, that aren’t the same element at all, with atomic numbers 6 and 7, respectively. To find out how many neutrons an isotope harbors, subtract its atomic number from its mass number.
Do isotopes actually do anything?
For the most part, no. Generally speaking, there’s little to no difference in how various isotopes of the same element behave. This is partly a function of how we decide what each element ‘is’: roughly three-quarters of naturally-occurring elements are a mixture of isotopes. The average mass of a bunch of these isotopes put together is how we determine those elements’ standard atomic weights.
But, chiefly, it comes down to the point we’ve made previously: without differences in their electron shell, isotopes simply lack the means to change their chemical behavior. Which is just peachy for us. Taken together, the 81 stable elements known to us can boast some 275 stable isotopes. There are over 800 more radioactive (unstable) isotopes out there — some natural, and some we’ve created in the lab. Imagine the headache it would cause if they all behaved in a different way. Carbon itself has 3 stable isotopes — would we even exist today if each had its own quirks?
One element whose isotopes do differ meaningfully, however, is the runt of the periodic table: hydrogen. This exception is based on the atom’s particular nature. Hydrogen is the simplest chemical element, one proton orbited by one electron. Therefore, one extra neutron in the core can significantly alter the atom’s properties.
For example, two of hydrogen’s natural isotopes, H-2 and H-3, have 1 and 2 neutrons respectively. Carbon (Z=6) has 2 stable isotopes: C-12 and C-13, with 6 and 7 neutrons respectively. In relative terms, there isn’t a huge difference in the neutrons’ share in their cores: they represent 50%, and 66.6% of the atoms’ weight in H-2, H-3, and 50% and 54-ish% of the total mass in C-12 and C-13. In absolute terms, though, the difference is immense: one neutron will double the mass of a hydrogen atom — two neutrons will triple it. For comparison, a single neutron is just 16.6% of a carbon atom’s mass.
While isotopes are highly similar chemically, they do differ physically. All that weight can alter how isotopes of light elements, hydrogen especially, behave. One example of such differences is the kinetic isotope effect — basically, heavier isotopes of the same element tend to be more sluggish during chemical reactions than lighter isotopes. For heavier elements, this effect is negligible.
Another quirky property of isotopes is that they tend to behave differently when exposed to infrared range than the ‘default’ elemental atoms. So, molecules that contain isotopes will look different to the same molecule sans isotopes when seen through an infrared camera. This, agian, is caused by their extra mass — the shape and masses of atoms in a molecule change how it vibrates, which in turn, changes how they interact with photons in the infrared range.
Where do isotopes come from?
Long story short, isotopes are simply atoms with more neutrons — they were either formed that way, enriched with neutrons sometime during their life, or are originated from nuclear processes that alter atomic nuclei. So, they form like all other atoms.
Lighter isotopes likely came together a bit after the Big Bang, while heavier ones were synthesized in the cores of stars. Isotopes can also form following the interaction between cosmic rays and energetic nuclei in the top layers of the atmosphere.
Isotopes can also be formed from other atoms or isotopes that have undergone changes over time. One example of such a process is radioactive decay: basically, unstable isotopes tend to shift towards a stable configuration over time. This can cause one unstable isotope to change into a stable one of the same element, or into isotopes of other elements with similar nucleic structures. U-238, for example, decays into Th-234.
This process, known as beta decay, occurs when there are too many protons compared to neutrons in a nucleus (or vice-versa), so one of them transforms into the other. In the example above, the uranium atom is the parent isotope, while the thorium atom is the daughter isotope. During this process, the nucleus emits radiation in the form of an electron and an antineutrino.
What are isotopes good for?
One of the prime uses for isotopes is dating (like carbon dating). One particular trait of unstable isotopes is that they decay into stable ones — but they always do so with the exact same speed. For example, C-14’s half-life (the amount of time needed for half of all isotopes in a sample to decay) is 5,730 years.
C-14 is formed in the atmosphere, and while an organism is alive, it ingests about one C-14 atom for every trillion stable C-12 isotopes through the food it eats. This keeps the C-12 to C-14 ratio roughly stable while it is alive. Once it dies, intake of C-14 stops — so by looking at how many C-14 atoms a sample has, we can calculate how far down C-14’s half-life it’s gone, meaning we can calculate its age.
At least, in theory. All our use of fossil fuels is pumping more C-14 isotopes into the atmosphere than normal, and it’s starting to mess up the accuracy of carbon dating.
To see how many C-14 atoms something has, we use accelerator mass spectrometry — a method that separates isotopes via mass.
PET (Positron-emission tomography) scans use the decay of so-called ‘medical isotopes‘ to peer inside the body. These isotopes are produced in nuclear reactors or accelerators called cyclotrons.
Finally, we sometimes create ‘enriched’ materials, such as enriched Uranium, to be used in nuclear reactors. This process basically involves us weeding through naturally-occurring uranium atoms via various methods for heavier isotopes, then separating those. The metal we’ve already weeded for heavier isotopes (which are more unstable and thus more radioactive than ‘regular’ uranium) is known as ‘depleted uranium’.