What is an Isotope: Explanation, Review, and Examples | Albert Resources (2024)

Understanding isotopes is essential to moving forward with any study of chemistry. In the following blog post, we will explore isotopes and their significance. We’ll cover what an isotope is, how to write its symbols, and methods to find its abundance. This post will also review important foundational concepts such as atomic structure, number, and mass.

Introduction to Isotopes

What is an Isotope?

Isotopes are variations of an element that have the same number of protons but different numbers of neutrons. Having a different number of neutrons results in the same elements having different mass numbers. Subsequently, isotopes are crucial in understanding how elements behave and vary.

How to Write Isotopes

Scientists use a specific notation to represent isotopes, which includes the element’s symbol and mass number. For example, carbon-14 is written as ^{14}{C}, where 14 represents the mass number (the sum of protons and neutrons) and {C} is the symbol for carbon.

Examples of Isotopes

So, let’s explore some examples. Hydrogen has three isotopes: protium (^1{H}), deuterium (^2{H}), and tritium (^3{H}). While they all have one proton, their neutron counts vary, resulting in different mass numbers.

In another case, carbon-14 (^{14}{C}) is an isotope used in carbon dating to determine the age of ancient artifacts. This isotope has 6 protons and 8 neutrons, giving it a mass number of ^{14}{C}.

Isotopes and Atomic Structure

Atomic Number and Mass Number

In general, the atomic number of an element represents the number of protons in its nucleus. Specifically, it defines the element’s identity and determines its place on the periodic table. For example, carbon ({C}) has an atomic number of 6, indicating that it contains 6 protons.

The mass number, on the other hand, is the sum of protons and neutrons in an atom’s nucleus. It represents the total number of particles (protons and neutrons) in the nucleus. In particular, let’s take the example of hydrogen. Hydrogen usually has one proton, but its isotopes, deuterium and tritium, have different mass numbers due to varying numbers of neutrons. Deuterium (^{2}{H}) has a mass number of 2 (1 proton + 1 neutron), while tritium (^{3}{H}) has a mass number of 3 (1 proton + 2 neutrons).

Because neutrons are neutral particles, the number of neutrons can vary without affecting the identity of the element. If an element has more protons, then it would no longer be the same element. The number of neutrons can vary and this is what causes isotopes to occur. The atomic number does not change. However, its mass does change, which results in the isotope. Let’s look more deeply into the structure of the atom.

Protons, Neutrons, and Electrons

Recall that protons are positively charged particles found in the nucleus of an atom. They carry a charge of +1 and contribute to the atom’s atomic number. Neutrons, on the other hand, are neutral particles also located in the nucleus. They have no charge and contribute to the atom’s mass number.

Lastly, electrons orbit around the nucleus in specific energy levels or shells. These negatively charged particles balance the positive charge of the protons and contribute to the atom’s overall neutrality. The number of electrons in an atom is equal to the number of protons, maintaining electrical balance. Review more about the parts of the atom in the following blog post.

Let’s consider chlorine ({Cl}). Chlorine has an atomic number of 17, meaning it contains 17 protons in its nucleus. The most common isotope of chlorine, chlorine-35 (^{35}{Cl}), has a mass number of 35 (17 protons + 18 neutrons). This isotope also has 17 electrons to balance the positive charge of the protons.

In summary, understanding the atomic structure and the interplay between protons, neutrons, and electrons is crucial to comprehending the behavior of elements and the variations observed in isotopes. This video from the American Chemical Society shows the structure of isotopes and some more examples.

Isotopes, Abundance, and Average Atomic Mass

Isotopic Abundance

Isotopic abundance refers to the relative amount or percentage of each isotope of an element found in a naturally occurring sample. Different isotopes of an element may have different abundances in nature. This variation in abundance is due to factors such as the element’s origin and the environment. The graph that follows shows the natural abundance of sulfur isotopes as an example.

How to Find the Abundance of an Isotope

To determine the abundance of an isotope, scientists use techniques such as mass spectrometry and spectroscopy. These methods allow for the measurement and analysis of isotopic ratios within a sample. By comparing the intensity of specific isotopic peaks, the abundance of each isotope can be determined. You can read more about mass spectrometry and how it’s used to study isotopes in this related article.

How to Calculate Atomic Mass of Isotopes

The atomic mass of an element accounts for the mass of each isotope and its respective abundance. To calculate the average atomic mass, we multiply the mass of each isotope by its abundance (expressed as a decimal or percentage), and then sum the products.

Calculation of Average Atomic Mass using Isotopic Abundance

Let’s consider the element chlorine ({Cl}) as an example. Chlorine has two naturally occurring stable isotopes: chlorine-35 (^{35}{Cl}) with an abundance of approximately 75\%, and chlorine-37 (^{37}{Cl}) with an abundance of around 25\%. In order to calculate the average atomic mass of chlorine, we multiply the mass of each isotope by its abundance and sum the results:

The average atomic mass of chlorine can therefore be calculated:

Cl = (^{35}{Cl}\text{ mass} \times 0.75) + (^{37}{Cl}\text{ mass} \times 0.25)

By substituting the respective masses and abundances, we can calculate the average atomic mass of chlorine.

Calculating the average atomic mass enables scientists to establish a weighted average of isotopic masses and better understand the behavior and properties of elements.

Isotopes in Everyday Life

Medical Applications of Isotopes

Chemistry and isotopes play a crucial role in various medical applications, such as aiding in diagnostics, treatments, and research. One example is the use of technetium-99m (^{99\text{m}}{Tc}), a radioactive isotope, in nuclear medicine. It is used for diagnostic imaging techniques such as single-photon emission computed tomography (SPECT). Technetium-99m emits gamma rays, which allow medical professionals to visualize internal organs and identify potential abnormalities.

Isotopes in Environmental Sciences

Isotopes also contribute to our understanding of the environment and natural processes. For instance, the stable isotope oxygen-18 (^{18}{O}) is utilized in environmental studies. By analyzing the isotopic composition of water samples, scientists can determine the sources and pathways of water, study hydrological cycles, and gain insights into climate change. Isotope analysis of ice cores, for example, provides valuable information about past climate conditions and long-term environmental changes.

Understanding the diverse applications of isotopes in everyday life helps us appreciate their significance beyond the confines of the laboratory. Isotopes are valuable tools in the fields of medicine and environmental sciences, enabling advancements in diagnostics, treatments, and our understanding of the world around us.

Conclusion: What is an Isotope?

In conclusion, by delving into isotopes, we gained a deeper understanding of the atomic structure. We’ve also seen the interplay between protons, neutrons, and electrons, and their impact on an element’s properties. Once you comprehend what an isotope Is, you can appreciate its wide-ranging implications, from advancing medical diagnostics and treatments to providing insights into environmental processes.