A 2.000 g sample of magnesium was burned in air to form an oxide.
After the product was purified it was found to have a mass of 3.317 g. What is the empirical formula for the product?
This is how I tried to answer...
1 mol of Mg= 24.3g
1 mol of O= 16.0 g
2.0 g Mg = 0.0823 mol Mag
1.317 g O (3.317-2.0) = 0.0823 mol O
empirical formula is MgO
Looks OK to me. Just to be picky watch what you are doing with the units and balancing equations, for example
2.0 g Mg = 0.0823 mol Mg
does not balance in terms of number or units.
Similarly
1.317 g O (3.317-2.0) = 0.0823 mol O
does not balance in terms of number or units.
Better as
Number of moles of magnesium is
2.0 g / 24.3 g mol^-1= 0.0823 mol
and number of moles of oxygen is
1.317 g /16 g mol^-1 = 0.0823 mol
To determine the empirical formula of the product, we need to find the ratio of the elements present in the compound. In this case, we're given the mass of magnesium before it reacted (2.000 g) and the mass of the product after purification (3.317 g).
To find the mass of oxygen that combined with magnesium, we subtract the mass of magnesium from the total mass of the product:
Mass of oxygen = Mass of product - Mass of magnesium
Mass of oxygen = 3.317 g - 2.000 g = 1.317 g
Now, we can calculate the moles of magnesium and oxygen using their respective molar masses. The molar mass of magnesium (Mg) is 24.31 g/mol, and the molar mass of oxygen (O) is 16.00 g/mol.
Moles of magnesium = Mass of magnesium / Molar mass of magnesium
Moles of magnesium = 2.000 g / 24.31 g/mol ≈ 0.082 mol
Moles of oxygen = Mass of oxygen / Molar mass of oxygen
Moles of oxygen = 1.317 g / 16.00 g/mol ≈ 0.082 mol
After obtaining the moles of both elements, we divide each by the smallest number of moles to get the simplest whole-number ratio. In this case, both magnesium and oxygen have approximately 0.082 mol. Therefore, the ratio is 1:1.
Thus, the empirical formula of the product is MgO, which represents one atom of magnesium and one atom of oxygen.