Isotopes (Grade 10 NSC Matric Physical Sciences): Revision Notes
Isotopes
What are isotopes?
Isotopes are atoms of the same element that have identical numbers of protons but different numbers of neutrons in their nuclei. This means they share the same atomic number (Z) but have different mass numbers (A).
The key point to understand is that the chemical behaviour of an element is determined by the number of protons and electrons it contains. When a neutron is added or removed from an atom's nucleus, the chemical properties remain unchanged because the electron arrangement stays the same.
This means isotopes of the same element will always occupy the same position in the periodic table and behave identically in chemical reactions.
Key characteristics of isotopes
Isotopes share several important features:
- Same chemical properties - They react in identical ways because they have the same number of electrons
- Same atomic number - They contain the same number of protons
- Different mass numbers - They contain different numbers of neutrons
- Different nuclear stability - Some isotopes may be more stable than others
- Same element identity - They are all versions of the same element
Whether a carbon atom contains 6, 7, or 8 neutrons, it will always be carbon with the element symbol C, because it contains 6 protons. The number of protons defines the element's identity.
Isotope notation
Isotopes can be written using the notation E-A, where:
- E represents the element symbol
- A represents the atomic mass (mass number)
Worked Example: Chlorine Isotopes
- Cl-35 has an atomic mass of 35 u (17 protons + 18 neutrons)
- Cl-37 has an atomic mass of 37 u (17 protons + 20 neutrons)
Both isotopes contain 17 protons (which makes them chlorine), but they differ in their neutron count.
Natural occurrence of isotopes
In nature, different isotopes of the same element occur in specific percentages called natural abundance. These percentages are not usually equal - one isotope is typically more common than others.
Using chlorine as our example:
- Cl-35 makes up 75% of all chlorine atoms on Earth
- Cl-37 makes up the remaining 25% of chlorine atoms
This natural variation in abundance affects the average atomic mass we see listed in the periodic table. The periodic table values are weighted averages, not the mass of any single isotope.
Calculating average atomic mass
The average atomic mass of an element accounts for all its naturally occurring isotopes and their abundances. Here's how to calculate it:
Formula:
Worked Example: Average Atomic Mass of Chlorine
Given information:
- Chlorine-35: 75% abundance, mass = 35 u
- Chlorine-37: 25% abundance, mass = 37 u
Step 1: Calculate the mass contribution of Cl-35
- Mass contribution = u
Step 2: Calculate the mass contribution of Cl-37
- Mass contribution = u
Step 3: Add the contributions together
- Average atomic mass = u
This calculated value of 35.5 u matches the atomic mass for chlorine shown in the periodic table, confirming our calculation is correct.
Key Points to Remember:
- Isotopes are atoms of the same element with different numbers of neutrons
- Chemical properties remain identical because protons and electrons are unchanged
- Isotopes are written using E-A notation (element symbol-mass number)
- Natural abundance determines how common each isotope is in nature
- Average atomic mass is calculated using weighted averages of all isotopes and their abundances