And how do scientists measure the average amount of molecular mass for each atom, they cant check every atom of the particular element in the universe to determine its molecular mass , and if scientists take samples from areas rich in a particular type of isotopes, wouldnt it affect the total average molecular mass of an element.
In: Chemistry
The isotope distribution actually affects the mass of molecules build with them. For example a water molecule made with deterium (1 proton 1 neutron) has a higher mass instead of protium hydrogen (1 proton 0 neutron. Water with an high fraction of deterium water molecules has a higher density than normal occuring water (which just have a very small amount of deuterium water). That’s why it’s called heavy water.
The element weight you can find in the periodic table is a an average value of the natural occuring isotope distribution (on earth). That’s the reason why these are not integers, as you would expect for integer amount of protons and neutrons, but fractional values. The atoms in hydrogen gas has a slightly higher mass than 1 u, as these are not only protium with 1 u, but also a few deuterium atoms with 2u.
The isotope distribution of elements on earth does not vary that much, so these average values are totally fine for most applications. Depending on the source you might find slightly different numbers for the elements however (as it depends on what materials were exactly measured).
If you need really exact values for a specific sample, you can do mass spectroscopy to determine the isotope distribution of a sample, and therefore it’s average atom weights.
Would it not be more odd if every atom created of a given element had the same amount of neutrons?
>they cant check every atom of the particular element in the universe
You can draw conclusions of a population without polling every member of that population. They’re just the average atomic mass you’d expect to find in a generic sample. There can definitely be reasons that a sample would *not* follow the average as well. When processing uranium from Oklo, it was found that there was less U-235 than you would normally expect. It turns out nearly two billion years ago, the conditions were just right (concentration of uranium, water as a neutron moderator) that the uranium deposit went critical and was a naturally occurring fission reactor.
On average, however, samples are average, and have a predictable distribution of isotopes.
Elements are produced by various means in nature, some of which are stable, some of which are not, and the mixture of not-stable isotopes with stable isotopes is very predictable. This is the reason why atomic weights on the periodic table are an average value, not a specific integer number as you might expect.
Let’s take a very well known example for this. The molecular weight of the most stable isotope of carbon is 12. But you’ve no doubt heard of Carbon-14. C14 is constantly being produced in the atmosphere by the action of cosmic rays on nitrogen – when nitrogen (7 protons) interacts with a high energy particle, it can produce carbon-14 (6 protons), which is unstable, with a half-life of 5700 or so years. The production of C14 from N14 in the atmosphere is predictable and constant, leading to a state of equilibrium between these two isotopes, and that proportion exists within every living carbon-based life form. When a living thing dies, however, it ceases to interact with its environment, and the unstable C14 gradually decays into stable C12.
The average mass is just for earth, baring any anomalies like naturally occurring reactors, the isotope distribution does not vary much, as most of the stuff on earth has been here for the same time.
So you just take a couple of samples from all over the planet, and determine the isotopes of an element, and in virtually all cases the distribution is doing to be near identical.
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