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What is the basic difference ?

  1. Mar 21, 2016 #1
    Atomic mass,molecular mass and formula mass: how are they different and how are they used differently ?
  2. jcsd
  3. Mar 21, 2016 #2
    An atomic mass unit (symbolized AMU or amu) is defined as precisely 1/12 the mass of an atom of carbon-12. The carbon-12 (C-12) atom has six protons and six neutrons in its nucleus. In imprecise terms, one AMU is the average of the proton rest mass and the neutron rest mass.
    In imprecise terms, one AMU is the average of the proton rest mass and the neutron rest mass. This is approximately 1.67377 x 10 -27 kilogram (kg), or 1.67377 x 10 -24 gram (g). The mass of an atom in AMU is roughly equal to the sum of the number of protons and neutrons in the nucleus.


    Both the terms relative atomic mass and atomic weight are sometimes loosely used to refer to a technically different standardized expectation value, called the standard atomic weight. This value is the mean value of atomic weights of a number of "normal samples" of the element in question. For this definition, "[a] normal sample is any reasonably possible source of the element or its compounds in commerce for industry and science and has not been subject to significant modification of isotopic composition within a geologically brief period."[3] These standard atomic weights are published at regular intervals by the Commission on Isotopic Abundances and Atomic Weights of theInternational Union of Pure and Applied Chemistry (IUPAC)[4][5] The "standard" values are intended as mean values that compensate for small variances in the isotopic composition of the chemical elements across a range of ordinary samples on Earth, and thus to be applicable to normal laboratory materials. However, they may not accurately reflect values from samples from unusual locations or extraterrestrial objects, which often have more widely variant isotopic compositions.

    The standard atomic weights are reprinted in a wide variety of textbooks, commercial catalogues, Periodic Table wall charts etc., and in the table below. They are what chemists loosely call "atomic weights."
    see https://en.wikipedia.org/wiki/Relative_atomic_mass
  4. Mar 21, 2016 #3
    The biggest problem is I don't even understand why atomic mass should be one twelfth the mass of one carbon 12 isotope.I know it's a relative way of finding the masses of other elements.I also know that it was chosen because it was a whole number (carbon 12 isotope has 12g of mass) but then practically how are we able to find the atomic mass of other elements ? Why do we do it that way ?
    I mean I amu is one twelfth the mass of one carbon 12 isotope but HOW ? And how do we find the atomic mass of other elements then ?
    I am sorry,I am confused !
  5. Mar 22, 2016 #4


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    Staff: Mentor

    See if this link helps: https://www.quora.com/Why-is-calculation-of-molecular-mass-based-on-carbon-12-and-not-hydrogen-1 [Broken]

    I believe mass spectrometry is commonly used.
    Last edited by a moderator: May 7, 2017
  6. Mar 22, 2016 #5


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    Gold Member

    Here is some history of atomic mass.

    Note that early on there was a proposal to have hydrogen as a basis ( I suppose H2 molecule ) with it listed as 2.

    Oxygen became a standard, with chemists using that which is found in nature , and physicists using the O-16 isotope as as their basis, leading to a discrepancy between the two groups for the masses of the elemental atoms.

    As a compromise, C-12 isotope was chosen as the basis, and we have what we have today.

    Years and years ago, a volume of gas would be "weighed", and with Avogadro's number, calculate the weight of an atom ( molecule ) of that gas.

    Nowadays, to get more accurate results, mass spectrometry is used.
  7. Mar 22, 2016 #6


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    Staff: Mentor

    Because we defined it to be so, by choosing - arbitrarily - 12.

    We could as well make it any other number, 5 or 42. These would be much less convenient to work with, but they are perfectly correct choices as well.
  8. Mar 22, 2016 #7
    If we use mass spectrometry,we can find weight of all elements in terms of grams.If that is true why do we use "amu" as it's unit? Why can't it be in grams ? And if 1 amu is one twelfth the mass of one carbon twelve isotope,then how do we find the masses of other elements after weighing them using mass spectrometry in terms of "amu," or "u"?
    My whole confusion is about how this unit amu and it's definition is helpful in weighing elements.
    Last edited by a moderator: May 7, 2017
  9. Mar 22, 2016 #8


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    Staff: Mentor

    Because we don't want to talk about atoms and molecules as being some multiple of 1.660539040(20)×10-24 g all the time. It's error-prone when doing basic calculations and an eyesore to constantly read.

    Here's how one type of mass spectrometer works: http://chemwiki.ucdavis.edu/Core/An..._Spectrometry/How_the_Mass_Spectrometer_Works

    So we just measure all the different variables when carbon-12 is used, and then when another atom or isotope is measured we can find the mass through various formulas and such.
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