What is the difference between mass number, atomic mass and average atomic mass? I know the mass number is the amount of protons + the amount of neutrons. The average mass is the weighed average of the isotopes that occur in nature. But then what is the (relative) atomic mass?

  • $\begingroup$ Related: Quick and simple explanation of molar mass, molecular mass and atomic mass and Units of mass on the atomic scale $\endgroup$ – user7951 Oct 12 '15 at 11:44
  • $\begingroup$ It is relative to the mass of another element. So, you may set the mass of carbon to be 12.000 and get all the masses of the elements relative to that; or, you may set the mass of oxygen to be 16.000 and get the masses of the elements relative to that. These two mass systems will be slightly different but internally consistent. Hope that helps. $\endgroup$ – user1945827 Oct 12 '15 at 11:48

It’s pretty simple:

  • The mass number is the integer you get if you count (and add up) the neutrons and the protons of a given element. Thus, a hydrogen atom of the $\ce{^1_1H}$ isotope has a mass number of $1$ (only proton), $\ce{^12_6C}$ has $12$ ($6$ protons, $6$ neutrons) and $\ce{^81_35Br}$ has $81$ (of which $35$ are protons, the remaining $46$ neutrons).

  • The atomic mass is what these atoms actually weigh in atomic mass units. For reasons that boil down to $E = mc^2$ (or so I believe) and the nonzero mass of an electron, this is not an integer except for one exception:

    • $\ce{^1H}$’s atomic mass is $1.007825032 1(4)$
    • $\ce{^12C}$’s atomic mass is exactly $12$. This is because $1~\mathrm{u}$ was defined as exactly $1/12$th of the mass of a carbon-12 atom.
    • $\ce{^81Br}$’s atomic mass is $80.916289(6)$
  • The average mass takes into account an elements different isotopes and their natural abundance and calculates an overall average. Thus, this is no longer defined on an isotopal basis but on an elemental one. The average masses of the elements discussed above are:

    • $\ce{H}$ has $1.00794 (7)$ (this is larger than the atomic mass of $\ce{^1H}$ due to the heavier isotopes.)

    • $\ce{C}$ has $12.0107 (8)$

    • $\ce{Br}$ has $79.904 (1)$ (this is lower than the atomic mass of $\ce{^81Br}$, because about half of the naturally occuring bromine atoms are $\ce{^79Br}$.)


well lets start it off with the example of a pizza, first imagine a pizza that is divided into 12 slices where 6 slices are meat free and the other 6 with meat and you were told that different pizzas have different amount of slices and different proportions of meat:non-meat ones the total number of pizza slices here is what we could refer to as "pizza number" now the seller tells you that this pizza(12 slices) weighs exactly 1200 grams and the other one (mini-pizza)weighs 600 grams now there weighs here is what we will refer to as pizza mass, now someone asks how much does this pizza weighs relative to the other we do bunch of calculations and we find out that the relative pizza mass of the mini should be on the scale of 1,and so the other pizza should be two{its double the mass} now what we have is relative pizza mass

now to sum it up lets consider the atom the atom nucleus does have proton and neutron and different atoms have different number of protons and neutrons in different proportions the number of both proton and neutron number is mass number it is practically helpful if we have different isotopes because each isotope have different mass number and so could be identified by their numbers,the atomic mass which is measured by mass spectrometer is the real mass of the atom which also varies between atoms if they were to sum up all those isotopic masses and take the average the result will be average atomic mass considering how small an atom is it will be awkward to use units like kg because the total mass will be small that you will need to write the number in standard form of such 1*10^-26 so the use the relative atomic mass to make it easier by using the scale of 12 amu(atomic mass unit) for a carbon 12 atom, and by that they can make other relative mass such as hydrogen(1.008),oxygen(15.999) and so on summary using an example

  • nitrogen atom with 7 neutron and 7 protons have mass number of 14 and thus isotope name is nitrogen-14
  • the atomic mass of nitrogen-14 is 2.32525265e-26
  • average atomic mass of all isotopes is 2.325x10^-26
  • relative atomic mass of all isotopes of nitrogen is 14.007 amu

*note:this is based on my personal experience if someone with any suggestions and arguments please consider a comment

*that is all what I have, thanks


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