Q.The need of the term average atomic mass.

Step 1. Most elements as they occur in nature exist as a mixture of two or more isotopes, each with a different individual atomic mass because they carry different numbers of neutrons.

Step 2. Since any real, weighable sample of the element automatically contains all of its naturally-occurring isotopes together, in their natural fixed proportions, using just ONE isotope's mass would give incorrect results for real calculations (like moles or molar mass) done on real samples.

Step 3. The average atomic mass — each isotope's mass multiplied by its percentage abundance, summed and divided by 100 — correctly reflects the mass behaviour of a real, natural sample of the element, which is why it is the value used everywhere in practical chemistry (and the value printed in the periodic table).

Explain that natural elemental samples are isotope mixtures, so only a weighted average correctly represents a real sample's mass behaviour.

• Explaining only WHAT average atomic mass is, without addressing WHY it is needed (i.e., without mentioning that natural samples are isotope mixtures).