The atomic mass of carbon is 12.01 g/mol, represented by the symbol 'C'. This isn't the mass of a single carbon atom, but rather a weighted average reflecting the different isotopes of carbon found in nature.
Understanding Atomic Mass
The atomic mass isn't simply the sum of protons and neutrons in the nucleus (the mass number). It's a weighted average considering the abundance of each carbon isotope. Carbon has two primary isotopes:
- Carbon-12 (¹²C): Makes up about 98.9% of naturally occurring carbon. Its mass number is 12 (6 protons + 6 neutrons).
- Carbon-13 (¹³C): Makes up about 1.1% of naturally occurring carbon. Its mass number is 13 (6 protons + 7 neutrons).
There are trace amounts of other isotopes, but these two dominate the calculation.
Calculating the Average Atomic Mass
The average atomic mass is calculated by considering the relative abundance of each isotope:
- Multiply: The mass of each isotope by its percentage abundance (expressed as a decimal).
- Add: The results from step 1 together.
Example:
- ¹²C: (0.989)(12 amu) = 11.868 amu
- ¹³C: (0.011)(13.003354 amu) = 0.143037 amu
Adding these together gives approximately 12.01 amu (atomic mass units), which is consistent with the reported atomic mass of carbon (12.01 g/mol). The difference in units (amu vs. g/mol) arises because amu refers to a single atom while g/mol refers to a mole of atoms (Avogadro's number).
Resources and Further Reading
For a deeper dive into atomic mass calculations and isotope abundances, refer to these resources:
