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Isotope Notation

 

Subscripts and superscripts can be added to an element’s symbol to specify a particular isotope of the element and provide other important information. The atomic number is written as a subscript on the left of the element symbol, the mass number is written as a superscript on the left of the element symbol, and the ionic charge, if any, appears as a superscript on the right side of the element symbol. If the charge is zero, nothing is written in the charge position. If the charge is +1 or −1, the convention is to write + or − (without the 1) as a superscript on the right. If the charge is +2, +3, −2, or −3, we write 2+, 3+, 2−, or 3− as the superscripts. 

 Image showing the general form of isotope notation with the element symbol with the atomic number as a subscript on the left, the mass number as a superscript on the left, and the charge (if any) as a superscript on the right

Examples are below.

Most abundant hydrogen isotope 

Image of the isotope symbol for hydrogen-1. The H has the atomic number 1 as a subscript on the left and the mass number 1 as a superscript on the left 

Most abundant isotope of uranium

 Image of the isotope symbol for uranium-238. The U has the atomic number 92 as a subscript on the left and the mass number 238 as a superscript on the left 

A sodium cation, Na+ 

 Image of the isotope symbol for the +1 cation of sodium-23. The Na has the atomic number 11 as a subscript on the left, the mass number 23 as a superscript on the left, and the plus superscript on the right

An aluminum cation, Al3+  

Image of the isotope symbol for the +3 cation of aluminum-27. The Al has the atomic number 13 as a subscript on the left, the mass number 27 as a superscript on the left, and the three plus superscript on the right  

An iodine anion, I−

Image of the isotope symbol for the minus one anion of iodine-127. The I has the atomic number 53 as a subscript on the left, the mass number 127 as a superscript on the left, and the minus superscript on the right

An oxygen anion, O2−

Image of the isotope symbol for the minus two anion of oxygen-16. The O has the atomic number 8 as a subscript on the left, the mass number 16 as a superscript on the left, and the two minus superscript on the right

Because all of the isotopes of an element have the same atomic number, the atomic number is often left off the isotope notation. Another way of naming isotopes uses the name of the element followed by the isotope’s mass number. For example, carbon-14 can be described in two ways:

Image of the isotope symbol for carbon-14. The C has the atomic number 6 as a subscript on the left and the mass number 14 as a superscript on the left          Image of the isotope symbol for carbon-14. The C the mass number 14 as a superscript on the left. The atomic number is left off.

All isotopes of an element have essentially the same chemical characteristics, and there is usually no need for the chemist to distinguish between them, but sometimes the differences between isotopes are very important. For example, although the iodine atoms found in nature are almost 100% iodine-127, iodine-131 can be formed in nuclear reactions. A major difference between these isotopes is that iodine-127 atoms are stable, and atoms of iodine-131 are unstable and undergo radioactive decay. Because these isotopes have virtually the same chemical properties, -1 ions of each are absorbed by our thyroid glands in the same way (thyroid tissue specifically absorbs and stores iodine, whereas other body tissues do not). A physician who suspects that a patient has a malfunctioning thyroid gland may perform a diagnostic test in which a very small amount of sodium iodide made with iodine-131 is administered. Instruments for detecting the levels and locations of the resulting radioactive emissions can then be used to study the thyroid gland’s activity.