Theory

Elements are defined by the number of protons (Z) in the nucleus of a nuclide or isotope-specific atom. The number of neutrons (N) defines the isotope of that element. The atomic weight (A) of a nuclide is equal to the sum of both protons and neutrons in the nucleus, or
A=N+Z
¹²Carbon: protons=6, neutrons=6, atomic weight=12
¹³Carbon: protons=6, neutrons=7, atomic weight=13

A single element can have two or more mass numbers due to differences in the number of neutrons that can occur in the nucleus. These different forms of a single element are called isotopes. While protons have a positive charge neutrons have no charge, so the number of neutrons does not affect the charge of a molecule.

Fig. 1
Plot of Z vs. N for nuclides up to tn (Z=50) showing the "stable" valley of the nuclides. Increases or decreases in N for given element produces increasingly unstable isotopes (decreasing T¹/₂). (Clark, I.D., and Fritz, P. 1997)

Elements usually have a common isotope that is the form most often found in nature. Because carbon, oxygen, and hydrogen are the elements that make up all organic matter, biologists are often interested in the isotopes of these elements. Each has common and rare forms. For instance, 98.8% of carbon atoms have a mass number of 12; the notation for this form is ¹²C. 1.1% of carbon has a mass number of 13 notes as ¹³C (See Table 1).

Table 1 Isotopes and Percent Abundances of Important Isotopes.

Element	Isotope	% Abundance
Carbon	12C	98.89
	13C	1.11
Nitrogen	14N	99.63
	15N	0.37
Oxygen	16O	99.759
	17O	0.037
	18O	0.204
Hydrogen	1H	99.985
	2H	0.015