Question

Question #e2f35

Answer 1

I fear there may be an error in the numbers you have given, or if these are experimental results, then you need to understand the sources of error in your experiment and how that limits the accuracy of your answer.

To work out an empirical formula for your data, you start withe masses of each type of atom, then divide by the relative atomic mass to get the number of moles of each, and turn this into a mole ratio. That's what an empirical formula is - the simplest whole-number ratio of the numbers of different types of atoms present in a compound.

              Mg                         O

Mass 0.17 0.07
`A_r` 24.3 16.0
Mol. 0.007 0.0044
Ratio. 1.6 : 1

So with your numbers we have 1.6 magnesiums in the formula for every oxygen. When we get fractions like this, we have to decide wether to round up, in which case we'd get `Mg_2O` as the formula or go to the nearest fraction I.e. 1.5 : 1 = 3 : 2 so we'd get `Mg_3O_2` as the empirical formula.

Neither of these are the actual empirical formula of magnesium I ice which is MgO. But your answer is typical of what happens if you weigh a piece of magnesium then burn it in a crucible and weigh the magnesium oxide formed, and calculate the mass of oxygen it combined with by subtraction.

The problem in the experiment is that some of the magnesium oxide escapes as smoke, and sometimes not all of the magnesium is burnt, so the mass of magnesium oxide measured is lower than it should be. This means the mass of oxygen comes out too low, and thus we see less of it in the empirical formula than there should be.