42.1.2 Isotopes and Isobars

It was discovered early in the 20th century that not all atoms of the same element (same $Z$) have the same mass. There are a range of neutron numbers that happily form a nucleus with $Z$ protons, creating a group of nuclei having the same $Z$-value (i.e., they are all the same chemical element) but different $A$-values. The atoms of an element with different values of $A$ are called isotopes of that element.

Chemical behavior is determined by the orbiting electrons. All isotopes of one element have the same number of orbiting electrons (if the atoms are electrically neutral) and thus have the same chemical properties, but different isotopes of the same element can have quite different nuclear properties.

The notation used to label isotopes is ${}^{A}Z$, where the mass number $A$ is given as a leading superscript. The proton number $Z$ is not specified by an actual number but, equivalently, by the chemical symbol for that element. Hence ordinary carbon, which has six protons and six neutrons in the nucleus, is written ${}^{12}\text{C}$ and pronounced “carbon twelve.” The radioactive form of carbon used in carbon dating is ${}^{14}\text{C}$. It has six protons, making it carbon, and eight neutrons.

More than 3000 isotopes are known. The majority of these are radioactive, meaning that the nucleus is not stable but, after some period of time, will either fragment or emit some kind of subatomic particle in an effort to reach a more stable state. Most of these radioactive isotopes are created by nuclear reactions in the laboratory and have only a fleeting existence. Only 252 isotopes are stable (i.e., nonradioactive) and occur in nature. We’ll begin to look at the issue of nuclear stability in the next section.

The naturally occurring nuclei include the 252 stable isotopes and a handful of radioactive isotopes with such long half-lives, measured in billions of years, that they also occur naturally. The most well-known example of a naturally occurring radioactive isotope is the uranium isotope ${}^{238}\text{U}$. For each element, the fraction of naturally occurring nuclei represented by one particular isotope is called the natural abundance of that isotope.

Although there are many radioactive isotopes of the element iodine, iodine occurs naturally only as ${}^{127}\text{I}$. Consequently, we say that the natural abundance of ${}^{127}\text{I}$ is 100%. Most elements have multiple naturally occurring isotopes. The natural abundance of ${}^{14}\text{N}$ is 99.6%, meaning that 996 out of every 1000 naturally occurring nitrogen atoms are the isotope ${}^{14}\text{N}$. The remaining 0.4% of naturally occurring nitrogen is the isotope ${}^{15}\text{N}$, with one extra neutron.

Series of nuclei having the same $A$-value (the same mass number) but different values of $Z$ and $N$ are called isobars. For example, the three nuclei ${}^{14}\text{C}\text{,}{}^{14}\text{N}$, and ${}^{14}\text{O}$ are isobars with $A=14$. Only ${}^{14}\text{N}$ is stable; the other two are radioactive.