The atomic mass of magnesium is the weighted average of the atomic masses of 1. all of the artificially produced isotopes of mg 2. all of the naturally occurring isotopes of mg 3. the two most abundant artificially produced isotopes of mg 4. the two most abundant naturally occurring isotopes of mg

2. All the naturally occurring isotopes of Mg.

Explanation:

You want to know the atomic mass of the magnesium you use in the lab. That’s “natural” magnesium. So, you must use the weighted average of all the naturally occurring isotopes in natural Mg.

1. and 3. are wrong. You won’t get the correct mass for natural Mg if you use only the artificial isotopes for your calculation.

4. is wrong. You must use all the naturally occurring isotopes. The two most abundant isotopes of Mg account for only 90 % of the atoms. If you ignore the other 10 %, your calculation will be wrong.

The atomic mass of Magnesium weighed is the average atomic mass of all the naturally produced isotopes of Mg.

The Mg used in the laboratory is the naturally occurring Mg. The natural isotopes of Mg have different atomic masses. So, the atomic mass of the weighed Mg is the average atomic mass of all naturally occurring isotopes.

Since there is no use of artificial isotopes in the reactions and in the laboratory, the atomic mass can't be the weighted average of artificially produced isotopes of Mg.

The average of two artificial isotopes will not all the isotopes and will not be correct.

The average of two natural isotopes does not consider all the other isotopes. And thus can not be the correct option.

For more information about atomic masses, refer to the link:

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