The History Behind the Atomic Mass Unit
Created | Updated Jan 25, 2006
The 'atomic mass' (more often also called 'atomic weight', which is slightly inaccurate but officially tolerated) describes, as the name cleverly says, the mass of an atom, the tiniest piece of elementary matter. It can be given in any unit of mass, like the kilogramme. As one can imagine the mass of an atom is very small, around 10-26 kg. So, instead of using this unhandy small number with the '10-26' all the time, it is a lot easier to use a different unit for the mass of atoms - the 'atomic mass unit'.
The 'atomic mass unit' is abbreviated 'amu' or more commonly just 'u' and corresponds to 1.6605 10-27 kg - the genesis of this unit is explained in more detail later. The unit corresponds - very roughly - to an effective mass of an average proton and/or neutron.
One of the most important scientific findings which finally led to the the development of the periodic table of the elements was that atoms of the same kind have the same mass. Say, all iron atoms have the same mass.1 For this reason the 'atomic mass' has been used as a specific property that defines what is an element, (in other words, the 'atomic mass' is what atoms of the same type have in common). This element specific mass is tabulated in the periodic table of the elements. It gives the average atomic mass of an element under consideration of the natural occurrence of isotopes (this is explained in more detail later).
As an examplary exercise, try to figure out the mass of 500 gold atoms (one gold atom has a mass of 196.7 u) in kilogrammes. Or try to figure out what element you have, provided you know you have exactly 500 atoms with a total mass of 1.6654 10-22 kg. The answers are in this footnote2)
Details: The Mess with the Mass
From Democritus to Dalton
The 'atom' is a very old concept. According to the Greek philosopher Democritus of Abdera (460 BCE) an atom would be the smallest piece of matter, which cannot be divided into smaller pieces. Throughout history this philosophical concept was forgotten and rediscovered many times. Around the 18th Century this concept infiltrated the scientific world. In the scientific definition of that time an atom would be the smallest piece of elementary matter, an element being 'matter homogenous in its properties and not be divisible into even more fundamental elements'. The common picture was that an atom would behave exactly like the element, with the only difference that it is a lot smaller. Antoine Lavoisier, Joseph Proust and John Dalton, to name a few, refined this picture from a chemical perspective. In the end, it was found that the elements that form a compound always come in a very specific ratio of mass. In that way it was possible to assign a relative mass to the individual elements. Dalton started synthesizing all kinds of compounds and calculating relative masses for elements. He figured out that the lightest one is hydrogen and then suggested using it as a reference with a mass of 1 unit. Elements, and consequently the atoms, now had a specific mass that could be described in units of hydrogen. This mass was originally termed the 'atomic weight' (which only recently mutated to the more accurate term 'atomic mass'). The 'atomic weight' is a fundamental property of an element.
Dalton's choice for hydrogen as the reference for the atomic weight is comprehensible. However, hydrogen is an unhandy element from a chemist's perspective. It is difficult to synthesize compounds with hydrogen. Oxygen for that matter was a lot easier to deal with: All a chemist had to do is 'burn' the element, or oxydize it. The weight of the element before and after the oxydation would lead to the relative weight relative to oxygen. Many scientists preferred to use a table based on oxygen.
Boyle, Avogadro, Loschmidt et al.
A big problem of the atomic unit was that all the data was gained by weighing chemical products. The source for errors is very big and difficult to circumvent. As a consequence the values for the relative atomic masses elements did not always fit the picture and varied substantially from author to author. The idea that elements had a well-defined and uniform atomic weight was nice, but didn't seem to be realistic.
Another fundamental question that needed to be addressed was: How much does one atom weigh in kilogrammes? Or, put differently, how many atoms of oxygen are in a litre of oxygen? The answer came, quite naturally, becasue oxygen is a gas, from the gas-theory thermodynamics people (Joseph Gay-Lussac, Robert Boyle, Amadeo Avogadro and others). And the, somewhat cryptic, answer was: A certain number of atoms, or molecules3 of a gas takes up the same volume, no matter what kind of gas it is.4. In other words, take fifty molecules of any gas and you'll end up with the same volume if the pressure is constant.
The next thing the scientists figured out was, that the number of 'atomic weight units' in units was the same number of the weight in grammes when the volume of the gas (at atmospheric pressure) was exactly 22,4 litres. And so Amadeo Avogadro (1776-1856) came up with a number, which posthumously was named after him: The Avogadro number. This number, of which the unit is the 'mole' (abbreviated 'mol'), tells us the number of molecules in 22,4 litres of a gas, at normal atmospheric pressure. Avogadro didn't know how much that was in numbers, however.
Had he survived until 1870 he would come across other people's works that found out that this reasoning also worked for solutions of solids, however this was a lot more difficult to find out, because of the different densities and other parameters. An Austrian guy named Johann Loschmidt was the first person to estimate this number5 by considering the volume and the density of perfect crystals and by estimating the volume of an atom. This number, even today, is not known with ultimate accuracy because it is not possible to count atoms in a straightforward manner. It's value is 6.022 x 1023 units.
Getting it all together
We saw that the concept of atoms is older than the possibility to characterize atoms. Dalton's concept of relative weights in units of hydrogen were a step in the direction of characterizing atomic properties. By the end of the 19th century scientists figured out that a certain number of particles, say one mole, of any gas take up the same volume (at constant temperature and pressure), say 22.4 litres. The next step is to react 22.4 litres of a gas (oxygen) with anything else, iron for example. If matter consisted of atoms (this is a theory), then each oxygen atom will react with one iron atom, forming one molecule of ironoxyde. Weighing the formed ironoxyde, and dividing it by one mole gives the mass of one ironoxyde molecule. If the initial mass of the unreacted iron is known, it is easy to find the difference - which corresponds to the atomic mass of iron. Phew! And there are more problems ahead...
More Problems
There has been a lot of confusion between chemists and physicists in the end of the 19th Century when the atomic properties were being described in more detail. Chemistry had not figured out all the laws of stoichiometry, and used the terms 'molecule' and 'atom' intergangeably. The procedure described above (reacting 22.4 litres of oxygen with an iron block) was not reliable. Physics in turn was not ready yet as it was not known what 'atoms' and 'matter' looked like at all.
The idea behind the term 'atom' and 'element' also changed abruptly when Ernest Rutherford, James Chadwick, Niels Bohr and other folks started to figure out how an atom really looks like.
The atomic unit was chosen to be roughly equivalent to the average of the masses of the neutron and the proton in a nucleus6. It was chosen, deliberately, to be 1/12 of the mass of a 12C carbon isotope (which consists of 6 neutrons, 6 protons and 6 electrons) - that's also the reason why it is not exactly 1, but 1.661 10-27 kg. For example: The Oxygen-16 isotope716O is made of 8 protons and 8 neutrons, hence it has a mass of 16 u.
To complicate things a bit further, the elements have a variable number of neutrons. Carbon (C, 6 protons) for example exists with 6 (12C) and 7 (13C) neutrons8. These slightly different 'versions' of carbon, and generally of other elements, are called isotopes. Now, there is a natural occurrence or distribution of isotopes: Boron (B, 5 protons), for example, occurs naturally as a mix of 20% 10B (with 5 protons and 5 neutrons) and 80% 11B (with 5 protons and 6 neutrons). Now, anything containing boron will contain the different isotopes in a 20:80 proportion. The atomic mass given in the periodic table considers this, so, for this reason one will find that the value for the atomic mass of boron in units is 10.81 u istead of 10 or 11. It is an average value considering the natural occurrence.
Now summing all this up, the atomic unit is a hodgepodge of masses, and the molecular mass an average value considering the natural occurrence.
The atomic masses are tabulated in the periodic table of the elements. The most practical unit, however, is the 'atomic unit' which is abbreviated 'u', it corresponds to 1/12 the mass of a 12C atom, or 1.66 10-27 kg. By using this unit one can avoid using the unhandy
b) The element has an atomic weight of: 1.6654 10-22 kg / 500 * 1.6605 10-27 kg = 200.59 u. This value is found for Mercury, Hg, in the periodic table.3Historically the distinction between atoms and molecules came a lot later. So, in old literature molecule and atom are often used interchangeably.4This is slightly incorrect as we know today, but only in very precise measurements the difference becomes detectable.5In older Literature one will often find Avogadro's number named after Loschmidt.6The atomic mass unit is not exactly equivalent to the average between the mass of a proton (1.673 10-27 kg) and a neutrons(1.675 10-27 kg). The atomic mass unit is actually smaller (by 1% or so), since some mass is 'used up' in the binding energy of the nucleus, but that's just a detail.799.8% of all oxygen is Oxygen-168There are more variations, like with 5 or 8 neutrons, but these are very, very rare.
