The Project: Introduction to
Atoms and Isotopes

modified from the University of Georgia website

The nucleus of atoms consists of protons and neutrons. Those two basic particles are often called nucleons. Since nucleons (protons (positively charged) and neutrons (no charge)) weigh approximately 1 unit each on the scale used to measure such things, the atomic weight of an atom can be treated as the same as the total number of its nucleons. Atomic weights of many of the elements listed in the Periodic Table are not neat whole numbers. This shows that these weights are averages of the differing atomic weights of two or more forms of that element.

A chemical element's atomic number is the number of positive charges (the number of protons) in the nucleus of each of its atoms (see this website for a nice periodic table!). This number (the number of protons) is the main characteristic of a given element. It never changes for all atoms of that element. So, if some atoms of an element have a different atomic weight from others, the difference must lie in the number of neutrons. Atoms of the same atomic number but different atomic weights are called isotopes.

Isotopes of elements can exist in both stable and unstable (radioactive) forms. Most elements that are important to biology (including carbon, hydrogen, oxygen, nitrogen, and sulphur) have two or more stable isotopes. The lightest isotope of each of these elements (lowest atomic weight) is found (in the earth, air and water) in much greater abundance than the other isotopes of the same element. Among stable isotopes, the most useful as biological “tracers” are the heavy isotopes of carbon and nitrogen. Carbon has an atomic number of 6, and nitrogen has an atomic number of 7. In the “common” form of the element, the number of protons and neutrons are equal in the nucleus. Therefore, the “common” isotope of carbon has an atomic weight of 12, while the “common” isotope of nitrogen has an atomic weight of 14. The atomic weights of isotopes are usually indicated by the number preceding the element abbreviation, so that C12 stands for the carbon isotope of atomic weight 12, while N14 stands for the nitrogen isotope that has atomic weight 14.

Nitrogen and carbon are extremely important. These elements are found in the earth, water, the atmosphere, and in all living things. Each of these elements has a heavy isotope (C13 and N15) with a natural abundance of ~1% or less of its total on Earth, and a light isotope (C12 and N14) that makes up almost all of the remainder.

Stable isotopes can be used in scientific experiments to learn things about Earth and its systems.

The following table shows the relative abundance of different isotopes of carbon, hydrogen, oxygen, nitrogen, and sulphur:

Table 1.1. Average Terrestrial (on Earth’s land) Abundances of the Stable

modified from the University of Georgia website

Isotopes of Major Elements (AN = atomic number)

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This special report was made possible by NSF Marine Geology and Geophysics Award Nos. OCE-0326573 to Fredrick Prahl (Oregon State University)and OCE-0324299 to Brian N. Popp (University of Hawaii) for study of "Alkenone Production and Productivity in Contrasting Surface Water Environments in the North Pacific Ocean."