According to the periodic table , a single carbon has a mass of 12.011 amu. what is the mass of a single carbon atom in grams?
Carbon has two naturally occurring isotopes, carbon-12 and carbon-13. The more common isotope of carbon is carbon-12 with an abundance of 98.93% while carbon-13 has an atomic mass of 13.00335 amu. Carbon-12 is what was used to define the amu/Dalton. It is defined as 1/12 the mass of an atom of carbon 12. So an unbound ground state carbon-12 atom weighs 12 amu. The chemists resisted making the amu one-sixteenth the mass of an oxygen-16 atom; it would change their atomic weights by about 275 parts per million. Making the amu one-twelfth the mass of a carbon-12 nucleus, however, would lead to only a 42 parts per million change, which seemed within reason.
2 Answers
There is no single carbon atom with that mass of 12.011 units.
Explanation:
Natural carbon is formed by 98.93% of isotope C-12 with mass 12.000000 units, and 1.07% of isotope C-13 with mass 13.003355 units. If you take the average of ten thousand atoms, you get the following weighed average mass of carbon atoms:
Now, if we are interested to know the masses of actual carbon atoms in grams, we should take the most abundant isotope, C-12.
Its mass is exactly 12 u because of the definition of the 'unified atomic mass unit' or u as one-twelfth of that isotope carbon.
That mass corresponds to #1.660539040*10^-24 g#, thus the mass in grams of a single atom of isotope 12 is #12.000000*1.660539040*10^-24 g = 19,926468*10^-24 grams#.
I hope this would be useful
Well, no carbon atom has exactly a mass of #12.011 'amu'#..
Explanation:
The periodic table gives the average mass of all the elements, that is, taking the average mass of all the elements' isotopes, and then taking their abundance on Earth into account as well.
For instance, carbon does not only exist as #'^12C#, but can also exist as #'^13C# and #'^14C#. However, the latter isotopes are not as common as #'^12C#, and so the mass of carbon on average would be a little bit over #12#, i.e. #12.011#.
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If you still want to know what will the average mass of a carbon atom in grams, we can find that out as follows.
Stoichiometry - What Is The Relation Between The Amu And ..
So, the average mass of a carbon atom will be approximately #2*10^-23# grams.
Related questions
Taken from http://www.sizes.com/units/atomic_mass_unit.htm
History of the atomic mass unit
Stanislao Cannizzaro (1826–1910), the pioneer in this field, adopted the hydrogen atom as a standard of mass and set its atomic weight at 2. Others accepted the idea of using a specific atom as a standard of mass, but preferred a more massive standard in order to reduce experimental error.
As early as 1850, chemists used a unit of atomic weight based on saying the atomic weight of oxygen was 16. Oxygen was chosen because it forms chemical compounds with many other elements, simplifying determination of their atomic weights. Sixteen was chosen because it was the lowest whole number that could be assigned to oxygen and still have an atomic weight for hydrogen that was not less than 1.
12 Amu Carbon To Grams
The 0=16 scale was formalized when a committee appointed by the Deutsche Chemische Gesellschaft called for the formation of an international commission on atomic weights in March 1899. A commission of 57 members was formed. Since the commission carried on its business by correspondence, the size proved unwieldy, and the Gesellschaft suggested a smaller committee be elected. A 3-member International Committee of Atomic Weights was duly elected, and in 1903 issued its first report, using the 0=16 scale.5
Taking isotopes into account
The discovery of isotopes complicated the picture. In nature, pure oxygen is composed of a mixture of isotopes: some oxygen atoms are more massive than others.
Chemical Elements.com - Carbon (C)
This was no problem for the chemists’ calculations as long as the relative abundance of the isotopes in their reagents remained constant, though it confirmed that oxygen’s atomic weight was the only one that in principle would be a whole number (hydrogen’s, for example, was 1.000 8).
Physicists, however, dealing with atoms and not reagents, required a unit that distinguished between isotopes. At least as early as 19276 physicists were using an atomic mass unit defined as equal to one-sixteenth of the mass of the oxygen-16 atom (the isotope of oxygen containing a total of 16 protons and neutrons).
Carbon Amu Mass
In 1919, isotopes of oxygen with mass 17 and 18 were discovered.7 Thus the two amu’s clearly diverged: one based on one-sixteenth of the average mass of the oxygen atoms in the chemist’s laboratory, and the other based on one-sixteenth of the mass of an atom of a particular isotope of oxygen.
Cached
In 1956, Alfred Nier (at the bar in the Hotel Krasnapolski in Amsterdam) and independently A. Ölander8, both members of the Commission on Atomic Masses of the IUPAP, suggested to Josef Mattauch that the atomic weight scale be based on carbon-12. That would be okay with physicists, since carbon-12 was already used as a standard in mass spectroscopy. The chemists resisted making the amu one-sixteenth the mass of an oxygen-16 atom; it would change their atomic weights by about 275 parts per million. Making the amu one-twelfth the mass of a carbon-12 nucleus, however, would lead to only a 42 parts per million change, which seemed within reason.
Mattauch set to work enthusiastically proselytizing the physicists, while E. Wichers lobbied the chemists.9 In the years 1959–1961 the chemists and physicists resolved to use the isotope carbon-12 as the standard, setting its atomic mass at 12.
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