5.5 Isotopes of the Elements

What is an Isotope?

Isotopes are various forms of the same element that have the same number of protons but a different number of neutrons. As the number of neutrons of an atom changes, so does its relative isotopic mass. The relative isotopic mass (also called mass number) is the sum of the protons and neutrons present in that isotope.

Mass Number (A) = Number of Protons + Number of Neutrons

Isotope symbols for elements are used to represent specific isotopes of atoms and include mass number (A) in superscript, atomic number (Z) in subscript, followed by the element symbol (X) in normal case (Figure 5.5a).

The number of protons in the nucleus of an atom is its atomic number (Z). This is the defining trait of an element. Its atomic number (Z) determines the identity of the atom. For example, any atom that contains six protons is the element carbon and has the atomic number 6, regardless of how many neutrons or electrons it may have. If you change the atomic number to 7, you are no longer dealing with carbon atoms, but nitrogen atoms. A neutral atom must contain the same number of positive and negative charges, so the number of protons equals the number of electrons. Therefore, the atomic number also indicates the number of electrons in an atom. The total number of protons and neutrons in an atom is called its mass number (A). The number of neutrons in that atom is therefore the difference between the mass number (A) and the atomic number (Z).

In summary:

[latex]\begin{array}{r @ {{}={}} l} \text{atomic number (Z) = } & \text{number of protons} \\[1em] \text{mass number (A) = } & \text{number of protons + number of neutrons} \\[1em] \text{A - Z = } & \text{number of neutrons} \end{array}[/latex]

Examples of Isotopes and their Properties:

As mentioned above, the symbol for a specific isotope of any element is written by placing the mass number as a superscript to the left of the element symbol. The atomic number is sometimes written as a subscript preceding the symbol, but since this number defines the element’s identity (atomic number), as does its symbol, it is sometimes omitted, as shown in Figure 5.5b. The various isotopes for the element carbon and the number of each subatomic particle in that isotope are shown below:

• Carbon-12 (or ¹²C) has the atomic number 6 and mass number 12 (six protons and six neutrons). It contains six protons, six neutrons, and six electrons
• Carbon-13 (or ¹³C) has the atomic number 6 and mass number 13 (six protons and seven neutrons). It contains six protons, seven neutrons, and six electrons.
• Carbon-14 (or ¹⁴C) has the atomic number 6 and mass number 14 (six protons and eight neutrons). It contains six protons, eight neutrons, and six electrons

When reading a specific isotope symbol, it is read as “element, mass number.” For instance, in the case of magnesium, ²⁴Mg is read as “magnesium 24,” and can be written as “magnesium-24” or “Mg-24.” All magnesium atoms have the atomic number 12, which means they have 12 protons in their nucleus. These isotopes differ only because a ²⁴Mg atom has 12 neutrons in its nucleus, a ²⁵Mg atom has 13 neutrons, and a ²⁶Mg atom has 14 neutrons.

While the atomic mass (mass number) of individual isotopes of a given element is different, their physical and chemical properties remain mostly the same. However, isotopes do differ in their stability. Carbon-12 (¹²C) is the most abundant of the carbon isotopes, accounting for 98.93% of carbon on Earth. Carbon-13 (¹³C) is stable but only accounts for about 1.07% abundance in nature. Carbon-14 (¹⁴C) isotopes are found in trace amounts in nature and are unstable or radioactive. Radioactive isotopes may decay over time by emitting neutrons, protons, and/or electrons and energy to obtain a more stable form. Further information relating to radioactivity and radioactive decay would be covered in more detail in nuclear chemistry or nuclear physics.

Information about the naturally occurring isotopes of elements with atomic numbers 1 through 10 is given in Table 5.5a. Note that in addition to standard names and symbols, the isotopes of hydrogen are often referred to using common names and accompanying symbols. Hydrogen-2, symbolized ²H, is also called deuterium and sometimes symbolized D. Hydrogen-3, symbolized ³H, is also called tritium and sometimes symbolized T.

Scientists in Action: Mildred Cohn, PhD.

As you will learn, isotopes are important in nature and especially in human understanding of science and medicine. Let’s consider just one natural, stable isotope: Oxygen-18, which is noted in the Table 5.5a above and is referred to as one of the environmental isotopes. It is important in paleoclimatology, for example, because scientists can use the ratio between Oxygen-18 and Oxygen-16 in an ice core to determine the temperature of precipitation over time. Oxygen-18 (Figure 5.5c) was also critical to the discovery of metabolic pathways and the mechanisms of enzymes.

Mildred Cohn pioneered the usage of these isotopes to act as tracers, so that researchers could follow their path through reactions and gain a better understanding of what is happening. One of her first discoveries provided insight into the phosphorylation of glucose that takes place in mitochondria. And the methods of using isotopes for this research contributed to entire fields of study.

Learn more about Mildred Cohn’s work and her motivating story of overcoming gender and religious biases in the article Mildren Cohn (1913-2009) [New Tab].

WATCH Women in Chemistry: Mildred Cohn (18min 42sec).

Video Source: Science History Institute (2012, September 10). Women in Chemistry: Mildred Cohn [Video]. YouTube.

Attribution & References

Except where otherwise noted, this page is adapted by Jackie MacDonald from “2.3 – Atomic Structure and Symbolism ” In Chemistry 2e (Open Stax) by Paul Flowers, Klaus Theopold, Richard Langley & William R. Robinson is licensed under CC BY 4.0. Access for free at Chemistry 2e (Open Stax) / Reused section on isotopes, rewrote learning objectives.