Please provide values below to convert Atomic mass unit [u] to gram [g], or vice versa.

Atomic Mass Unit to Gram Conversion Table

Experiments have shown that 1 amu = 1.66 × 10 −24 g. Mass spectrometric experiments give a value of 0.167842 for the ratio of the mass of 2 H to the mass of 12 C, so the absolute mass of 2 H is mass of 2H mass of 12C × mass of 12C = 0.167842 × 12amu = 2.104104 amu The masses of the other elements are determined in a similar way. Numerical value: 1.660 539 066 60 x 10-27 kg: Standard uncertainty: 0.000 000 000 50 x 10-27 kg: Relative standard uncertainty: 3.0 x 10-10: Concise form 1.660 539 066 60(50) x 10-27 kg: Click here for correlation coefficient of this constant with other constants. The original standard of atomic weight, established in the 19th century, was hydrogen, with a value of 1. From about 1900 until 1961, oxygen was used as the reference standard, with an assigned value of 16. The unit of atomic mass was thereby defined as 1 / 16 the mass of an oxygen atom. In 1929 it was discovered that natural oxygen contains. An atomic mass unit is the same thing as grams per mole (1 amu = 1 g/mol). It is also the same thing as a dalton (1 amu = 1 Da). So if you don't know the amu for one of your elements, you can search for this particular isotope online to find the amu and natural abundance specific to that particular isotope.

How to Convert Atomic Mass Unit to Gram

1 u = 1.6605402E-24 g
1 g = 6.0221366516752E+23 u

Example: convert 15 u to g:
15 u = 15 × 1.6605402E-24 g = 2.4908103E-23 g

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Why is the atomic mass unit (amu), rather than the gram, used to express atomic mass?

1 Answer

Because atoms are ridiculously small.

And the #'amu'# is numerically equivalent to the #'g/mol'#. For instance, if I were to be so lucky as to isolate #'1 atom'# of #'N'#, it would have a mass of

#14.007 cancel'amu' xx (1.6605 xx 10^(-24) 'g')/(cancel'1 amu')#

#= ul(2.326 xx 10^(-23) 'g')#

which is immeasurably small. We don't care for masses that small because we physically can't see or measure it. Instead, we care for masses we can touch, like #'1.000 g'# or #'12.50 g'#.

And that involves:

#1.000 cancel'g N' xx cancel'1 mol N'/(14.007 cancel'g N') xx (6.022 xx 10^23)/(cancel'1 mol')#

#= ul(4.299 xx 10^22 'N atoms')#

Amu Chemistry

#12.50 cancel'g N' xx cancel'1 mol N'/(14.007 cancel'g N') xx (6.022 xx 10^23)/(cancel'1 mol')#

#= ul(5.374 xx 10^23 'N atoms')#

You can clearly tell that this number of atoms is impossible to count. And so Avogadro's number, #6.022 xx 10^23 'mol'^(-1)#, was invented to describe this many particles...

#4.299 xx 10^22##'N atoms' xx ('1 mol')/(6.022 xx 10^23)#

#=##ul'0.0714 mols N'#

#5.374 xx 10^23##'N atoms' xx ('1 mol')/(6.022 xx 10^23)#

Value Of 1 Amu In Kg

#=##ul'0.8924 mols N'#

Value Of 1 Amu In Kg

And as you can see, these numbers look much nicer and more physically useful.

How To Find Amu

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