2.3 Isotopes

2026 Syllabus Objectives

Core: Define isotopes as different atoms of the same element that have the same number of protons but different numbers of neutrons
Core: Interpret and use symbols for atoms, e.g. ${}^{12}_{6}\text{C}$ , and ions, e.g. ${}^{35}_{17}\text{Cl}^-$
Supplement: State that isotopes of the same element have the same chemical properties because they have the same number of electrons and therefore the same electronic configuration
Supplement: Calculate the relative atomic mass of an element from the relative masses and abundances of its isotopes

Measuring the Mass of Atoms

The Need for a Standard ⚖️

A single atom cannot be weighed on a balance. However, scientists can compare the mass of one atom to another using an instrument called a mass spectrometer. Since we are comparing masses, the values obtained are relative values, and a standard is needed for comparison.

🔑 Key Term: A mass spectrometer is an instrument in which atoms or molecules are ionised and then accelerated; the ions are then separated according to their mass.

The Carbon-12 Standard

Carbon was chosen as the standard element because there are far more compounds containing carbon than any other element. The specific isotope carbon-12 is used as the reference standard.

One atom of carbon-12 is given the mass of exactly 12 units.

From this definition, we obtain the atomic mass unit (a.m.u.):

\text{1 a.m.u.} = \frac{1}{12} \times \text{mass of one atom of carbon-12}

Relative Atomic Mass

The masses of all other atoms are compared to carbon-12, giving us a series of values called the relative atomic mass ( $A_r$ ) for different elements.

🔑 Key Term: Relative atomic mass ( $A_r$ ) is the average mass of naturally occurring atoms of an element on a scale where the carbon-12 atom has a mass of exactly 12 units.

Examples of Relative Atomic Masses

Element	Atomic symbol	Relative atomic mass ( $A_r$ )
carbon	C	12
hydrogen	H	1
oxygen	O	16
calcium	Ca	40
copper	Cu	64
gold	Au	197

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