Radioactive Isotopes

Non-Radioactive Isotopes

Let's start by talking about isotopes. All elements can have isotopes, which are forms of that element that have the same number of protons but differ in the number of neutrons in its nucleus.

The atomic notation of an isotope looks like this: \( ^{A}_Z\text {X} \), where A is the mass number of the isotope, Z is the atomic number and X is the chemical symbol of the element.

For example, hydrogen has three naturally occurring isotopes: \( ^{1}_1\text {H} \), \( ^{2}_1\text {H} \), and \( ^{3}_1\text {H} \), with \( ^{1}_1\text {H} \) being the most abundant isotope of hydrogen (99.985%) . In terms of mass notation, we would write them as hydrogen-1, hydrogen-2, and hydrogen-3, where the number next to the element's name is simply the isotope's atomic mass.

Hydrogen-1, also known as protium, containing one proton (and consequently one electron) and zero neutrons in its nucleus. Hydrogen-2, also called deuterium, has 1 proton, 1 electron and 1 neutron. Hydrogen-3, also known as tritium, has one proton, one electron and 2 neutrons.

Because these isotopes have different masses, they behave differently in chemical reactions. For instance, a water molecule possessing a \( ^{2}_1\text {H} \) atom ( \( ^{2}\text {H}_2 \text {O} \) ) will undergo evaporation slower than a water molecule with a \( ^{1}_1\text {H} \) atom because it is heavier than \( ^{1}\text {H}_2 \text {O} \). Factors affected by the isotope composition of a molecule include, freezing point, boiling point, and Vapor Pressure.

• The melting point of \( ^{1}\text {H}_2 \text {O} \) is 0.00 °C at 760 torr, whereas the melting point of \( ^{2}\text {H}_2 \text {O} \) is equal to 3.81 °C at 760 torr.

Isotopes that possess a stable nucleus are known as non-radioactive isotopes. When an isotope is stable, it means that its nucleus will not spontaneously undergo what is called radioactive decay.

The elements that tend to have two or more stable isotopes are those elements that have a low atomic weight, such as hydrogen, Nitrogen, oxygen, sulfur, and Carbon! The image below shows the stable isotopes of oxygen: oxygen-16, oxygen-17, and oxygen-18.

But, how do we know whether an isotope is considered nonradioactive (stable) or radioactive (unstable)? Well, this depends on whether that isotope is found relative to the belt of stability. If the isotope is on the belt of stability, then it will be nonradioactive!

The belt of stability tells us that lighter isotopes (isotopes with up to 20 protons) are most stable when the neutron number/proton number ratio is close to 1. We won't go into more details at this time.

Radioactive Isotopes Examples

Now that know what nonradioactive isotopes are, let's dive into radioactive isotopes. Radioactive isotopes are unstable isotopes.

During nuclear decay, radiation is emitted by the emission of nuclear particles. There are different types of nuclear decay that an isotope might experience.

• If a radioactive isotope has a high neutron-to-proton ratio (high mass number compared to the mass number on the Periodic Table), it would be found above the belt of stability and undergo beta (β) decay. In beta decay, the nucleus loses one neutron and gains one proton.
• A nucleus containing a low neutron-to-proton ratio (low mass number than that on the Periodic Table) would be found below the belt of stability and would undergo positron emission or electron capture. In this case, the nucleus gains one neutron and loses on proton.
• Nuclei with atomic number > 83 tend to undergo alpha (α) decay, losing two proton and two neutrons.

Let's look at some examples of radioactive isotopes, also known as radioisotopes. Tritium (\( ^{3}_1\text {H} \) ) is a radioactive isotope of hydrogen. Tritium is used in the manufacture of glow in the dark painting, and also in H-bombs.

Another common radioactive isotope is nitrogen-16 (\( ^{16}_7\text {N} \) ). Nitrogen-16 has a mass higher than the mass of Nitrogen in the periodic table (14), so its nucleus will most likely undergo beta decay and have a nucleus like that of oxygen.

$$ ^{16}_{7}\text{N }\longrightarrow \text{ }^{0}_{-1}\text{e + }^{16}_{8}\text{O} $$

Radioactive Isotopes of Carbon

Carbon-12 (\( ^{12}_6\text {C} \) ) and carbon-13 (\( ^{13}_6\text {C} \) ) are both considered stable isotopes of carbon. However, there are some isotopes of carbon that are considered unstable, and and therefore radioactive.

Carbon-14 is a radioactive isotope of carbon (\( ^{14}_6\text {C} \) ). It has 6 protons and 8 neutrons and will most likely undergo beta decay to decay into a stable (nonradioactive) isotope: nitrogen-14.

$$ ^{14}_{6}\text{C }\longrightarrow \text{ }^{0}_{-1}\text{e + }^{14}_{7}\text{N} $$

Radioactive Isotopes Half-life

The half life of a radioactive isotope is referred to as the amount of time taken for \( \frac{1}{2} \) of a radioisotope sample to decay.

For example, iodine-131 (\( ^{131}_{53}\text {I} \)) is a radioactive isotope of iodine that undergoes beta decay, producing the nonradioactive isotope \( ^{131}_{54}\text {Xe} \) and releasing a beta (electron) particle. The half-life of this radioactive isotope is 8.0 days.

$$ ^{131}_{53}\text{I }\longrightarrow \text{ }^{0}_{-1}\text{e + }^{131}_{54}\text{Xe} $$

So, if you started with 10.0 mg of iodine-131, after 8.0 days, you would have 5.00 mg of iodine-131. After 2 half-lives (16 days), you would have 2.5 mg of iodine-131!

Let's look at an example of a problem.

Radioactive Isotopes Uses

To finish off, let's talk about the uses of radioactive isotopes. As we said at the start, carbon-14 is used in Carbon Dating.

For example, to figure out how old an animal or human bone is, the half-life of carbon-14 (5730 years) is used. For example, a bone from a prehistoric animal with 25% of the activity of carbon-14 would mean 2 half-lives. So, the amount of years passes since that animal died would be 11,000 years.

$$ 2.0 \text{ half-lives}\times \frac{5730\text{ years}}{1 \text{ half-life}} = 11,000 \text{ years} $$

Radioactive isotopes also have uses in medicine. For example, a radioactive isotope of gallium, \( ^{68}\text {Ga} \), is used in the detection of pancreatic cancer, whereas \( ^{67}\text {Ga} \) is used in abdominal imaging and also for tumor detection.

Phosphorus-32 is used in the treatment of leukemia, excess red blood cells and also pancreatic cancer. Another radioactive isotope, Yttrium-90 is used in the treatment of liver cancer. Currently, for the detection of bone lesions, and also in brain scans, Strontium-85 is commonly used.

Hopefully, you are not more confident in your understanding of radioactive isotopes!