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What is the difference between:

atomic mass

relative atomic weight

relative isotopic weight

standard atomic weight

atomic weight

relative atomic mass

I am told by Wikipedia that: relative atomic weight = atomic weight

and a the start of the article it says: For a single given sample, the relative atomic mass of a given element is the weighted arithmetic mean of the masses of the individual atoms (including their isotopes) so standard atomic weight = relative atomic weight

But at https://en.wikipedia.org/wiki/Relative_atomic_mass#Standard_atomic_weight#Definition it says:

It is a synonym for atomic weight, though it is not to be confused with relative isotopic mass. Relative atomic mass is also frequently used as a synonym for standard atomic weight and these quantities may have overlapping values if the relative atomic mass used is that for an element from Earth under defined conditions. However, relative atomic mass (atomic weight) is still technically distinct from standard atomic weight because of its application only to the atoms obtained from a single sample; \textbf{it is also not restricted to terrestrial samples, whereas standard atomic weight averages multiple samples but only from terrestrial sources}

I am told here that:

The standard atomic weight (Ar, standard) of a chemical element is the weighted arithmetic mean of the relative atomic masses (Ar) of all isotopes of that element weighted by each isotope's abundance on Earth.

But in this video I am told this is the definition of relative atomic mass.

Here I am told:

Atomic mass (ma) is the mass of an atom. A single atom has a set number of protons and neutrons, so the mass is unequivocal (won't change) and is the sum of the number of protons and neutrons in the atom.

Atomic weight is a weighted average of the mass of all the atoms of an element, based on the abundance of isotopes. The atomic weight can change because it depends on our understanding of how much of each isotope of an element exists.

In other words: Atomic weight = standard atomic weight

Also, which of these definitions is used to convert grams of a substance into moles

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    $\begingroup$ Related: Quick and simple explanation of molar mass, molecular mass and atomic mass $\endgroup$
    – user7951
    Jul 23, 2020 at 18:26
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    $\begingroup$ I have a bad feeling that the terminology is not properly standardised and that chemists are terminologically confused. However, this may not matter much in practice as the specific meaning is almost always obvious from context. I'm happy to be proved wrong, though. $\endgroup$
    – matt_black
    Jul 23, 2020 at 18:34
  • $\begingroup$ chemistry.stackexchange.com/questions/38082/… $\endgroup$
    – Mithoron
    Jul 23, 2020 at 19:08
  • $\begingroup$ @matt_black, I agree that the way atomic masses or atomic weights are taught in general chemistry is very confusing and superficial. $\endgroup$
    – AChem
    Jul 24, 2020 at 1:03
  • $\begingroup$ Thanks these websites clear some of the ambiguity $\endgroup$
    – user96130
    Jul 24, 2020 at 8:42

2 Answers 2

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Relative Atomic Mass vs Atomic Weight

Relative Atomic mass ($m_a$) is the measurement of the mass of an atom relative to carbon-12. It's dimension is M.[1] It can be measured in Da/u, kg or g.* [3] $$m_a \approx \mathrm{Mass\ Number} \times {1\ \mathrm{u}}$$ except for carbon-12 where they are the same.

Atomic weight is:

Relative atomic mass vs Standard atomic mass

Relative atomic mass or Relative atomic weight is the weighted average of the isotopes of an element according to their abundance in a given sample compared to the atomic mass of carbon-12. It's dimension is M/M or 1. [2] [4]

Standard atomic mass is the same except that the samples are restricted to terrestrial samples normally considering/averaging results from many different samples or results given by the CIAAW.

Relative Isotopic Mass vs Relative Atomic Mass

Relative Isotopic mass considers the mass of an isotope compared to the atomic mass of carbon-12. This is also dimensionless.

*Universal mass constant vs kg , g ,etc

The universal mass constant ($m_u$) is defined as $\frac{1}{12}$ of the mass of carbon-12 which defines 1 Da or 1 u. All measurements in Da are comparisons to the mass of a carbon-12 atom. Since the discovery of Avogadro's constant, we know: [3] $$ 1\ {\rm{Da}} = 1.660\ 539\ 066\ 60(50)\times 10^{-27}\ \mathrm{kg}$$

Molar Mass vs Molecular Mass (or Formula Mass)

Molar Mass is the grams of a substance in 1 mole of it measured in gram per mole. It's dimension is $\mathbf{M N^{-1}}$.

Molecular mass or molecular weight is the sum of the atomic masses of all the component atoms of each element measured in u. It's dimension is M. Formula mass or formula weight is the sum of the atomic masses of component ions (not sure of the difference between this and molecular mass - in terms of quantity).[2]

[1] see the answer to this Q by M Farooq

[2] What is the difference between "molecular mass", "average atomic mass" and "molar mass"?

[3] Mass number, (relative) atomic mass, average mass

[4] Quick and simple explanation of molar mass, molecular mass and atomic mass

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If you are confused, it is not a surprise because chemists remained confused for 200 hundred years. Welcome to the world of real science.

Atomic weight is the older term for atomic mass. Modern day hair splitters would say that mass and weight are different things but historically in the old chemical literature, they meant one and the same thing.

Nobody can repeat two centuries of work here, but briefly, atomic mass of an element meant that it is the mass of any atom that combines with the mass of one oxygen atom.

Now the key question is which leads to opening the Pandora box. What is the mass of one oxygen atom? People had to choose a number, because nobody knew the mass of oxygen atom. Some said, let us fix it to 200, some fixed it to some other number. This arbitrary "mass" of oxygen atom kept on changing, finally, someone said, let us fix it exactly to 16.000000000... (ad infinitum) units.

Notice that I did not need any crutches of moles or grams per mole. This is all modern day stories fed to fresh minds. We are still in early 1920s.

Now once the mass of oxygen is fixed, you can relatively easily assign an atomic weight of other elements with respect to oxygen

For example, atomic mass of hydrogen was determined to be 1.008 which is the mass which would combine with exactly 16 units of oxygen.

Now you can understand that chemists never talk about absolute atomic masses, they are all relative to some element, so it is relative atomic mass.

Today, the reference is carbon-12 (after 1960s) and there are many more nuances which may not be of interest at this moment.

Think of relative atomic weights like this. Suppose I want to assign an "atomic mass" to you, but I do not know your mass. However, I know your friend's mass which is arbitrarily assigned as 100 lb. Using your friend as a reference we can assign an atomic mass to you indicating whether you are heavier or lighter than your friend.

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