Isotopes are different versions of a chemical element. They have the same number of protons and electrons, but their neutron counts vary.

This means they can have different atomic masses, which can change their properties.

What are Isotopes?

• Isotope: Atoms of the same element with the same number of protons but different numbers of neutrons.

• Example: Let's look at carbon. It has three important isotopes:

  • Carbon-12: This has 6 protons and 6 neutrons. Its atomic mass is about 12.
  • Carbon-13: This has 6 protons and 7 neutrons. Its atomic mass is about 13.
  • Carbon-14: This has 6 protons and 8 neutrons. Its atomic mass is about 14.

How Atomic Mass is Different

The atomic mass of an isotope is figured out by adding up the number of protons and neutrons in its nucleus.

• Calculation:
  • Atomic mass = Number of protons + Number of neutrons
  • For Carbon-12: 6 + 6 = 12
  • For Carbon-13: 6 + 7 = 13
  • For Carbon-14: 6 + 8 = 14

Natural Abundance and Stability

Different isotopes can also be found in different amounts and have different stability:

• Carbon-12: Makes up about 98.89% of the carbon found in nature.
• Carbon-13: Makes up about 1.11% of natural carbon.
• Carbon-14: This is present in tiny amounts (about 0.0000000001%). It is radioactive, which means it changes over time, and is used for dating old materials.

Importance and Uses

The different atomic masses of isotopes help in many areas of science and industry:

• Carbon-14 Dating: This is used to find out how old ancient organic materials are.
• Medical Imaging: Some isotopes, like Carbon-11, are used in scans to see inside the body.
• Nuclear Energy: Isotopes like Uranium-235 and Uranium-238 have different properties and are important for nuclear reactions.

In summary, isotopes mainly differ in their atomic mass because of the different number of neutrons they have. This difference, along with their natural abundance and stability, allows them to be used in various scientific fields and applications.