Why is atomic mass based on Carbon-12?



I’ve been trying to read and comprehend the reasons why, but I still don’t get it.

In: Chemistry

Well originally it was decided to use hydrogen. Made sense, but technically challenging to actually weigh. Physicists/chemists moved on to oxygen which was easier. And hey. Oxygen is pretty common, so if we’re picking an arbitrary element Oxygen makes enough sense.

Then we discovered isotopes which screwed everything up. Some guys decided to use the average of oxygen isotopes found in nature, and others decided to use oxygen-16. Unfortunately the two different groups never reconciled and both systems remained in common usage.

Eventually a big international conferenced decided on C-12 which put the unit of atomic weight between the values of the two Oxygen based units.

Basically cause its really abundant and almost always you can find the same amount of Neutrons and protons and electrons inside it. In other words, this element is really abundant and pretty much the most consistent, in reality the mass is not 12, it is 12.something because there are some atoms that do variate, but 99% of the time they are the same.

Early on in learning about atoms it became clear that most of the mass is concentrated in the nucleus and that mass comes from protons and neutrons. It became clear that protons and neutrons are both about the same mass, to the point where you can pretty accurately describe the mass of an atom with just a whole number: the mass of Carbon 12 is 12, the mass of Oxygen 16 is 16, etc.

However, as more investigation was done it became clear that things aren’t quite that simple. Protons don’t have the same mass as neutrons and Carbon 12 isn’t exactly 3/4 the mass of Oxygen 16. This led to a need for a specific atomic mass unit.

When there’s a unit there needs to be a definition. The initial definition used was 1/16 the mass of oxygen, but this took two slightly different forms. One form specified that you should take 1/16 the *average* mass of a mixture of oxygen isotopes that matches what you find in nature. That definition was preferred in chemistry. The other form specified that you should take 1/16 the mass of Oxygen-16.

The first definition had more momentum behind it, but it’s also the less precise definition. How can you be certain that you have the right mixture of oxygen? It turned out that there’s no correct answer here: some isotopes are more common in oxygen taken from air, while others are more common when you take oxygen from water (the sun’s rays can cause atoms to form different isotopes, which skews the ratios in air).

The goal was to find a definition that is close to the value that was already in use while having the predictability of the physicists’ definition. Carbon-12 was up to the task, and is nice for being an atom that’s abundant and stable, while also having equal protons and neutrons.

That decision was made for practical convenience, not based on any fundamental physical truths. It would have been just as valid to define 1 amu (or dalton, if you prefer) as the mass of a hydrogen-1 atom, 1/4 the mass of a helium-4 atom, or 1/8 the mass of Beryllium-8 (which decays in about 10 attoseconds–that’s below femptoseconds, picoseconds, and nanoseconds, each a factor of 1000 shorter than the next). Obviously Beryllium-8 is a poor choice here for practical reasons, but it would have been valid as a definition nonetheless.

Well, it had to be something. And the unit ‘[yocto](https://en.wikipedia.org/wiki/Yocto-)gram’ hadn’t been invented. The choice of C-12 was a compromise between competing standards.

> The existence of two distinct units with the same name was confusing, and the difference (about 1.000282 in relative terms) was large enough to affect high-precision measurements. … For these and other reasons, in 1961 the International Union of Pure and Applied Chemistry (IUPAC), which had absorbed the ICAW, adopted a new definition of the atomic mass unit for use in both physics and chemistry; namely, 1/12 of the mass of a carbon-12 atom. This new value was intermediate between the two earlier definitions, but closer to the one used by chemists (who would be affected the most by the change).[11][12]