Final answer:
In a high school chemistry context, a student is asked to calculate the mass of impure iron produced from a reaction involving hematite and aluminum, taking into account the purity of the reactants and products. Stoichiometric relationships and molar masses are used to determine the theoretical yield, which is then adjusted for the desired purity of the produced iron.
Step-by-step explanation:
The question relates to a typical high school chemistry problem involving the concept of percent yield. The given scenario describes a reaction between hematite with a known purity and aluminum to produce impure iron. When calculating the quantity of the produced iron, it's necessary to account for the purity of the reactants and the desired purity of the product.
The hematite has a purity of 60%, meaning that only 60% of the 20 grams is actual iron oxide (Fe2O3). Therefore, the amount of pure Fe2O3 is 12 grams. The stoichiometry of the reaction yields iron, whose molar mass is 55.85 grams/mol. Assuming 100% efficiency, the mass of iron from 1 gram of Fe2O3 would yield 0.70 grams of iron.
However, we are looking for impure iron with a purity of 70%. This means the actual yield of pure iron must be adjusted for this final purity. We use stoichiometric calculations to derive the mass of pure iron and then adjust for the desired purity level to determine the mass of impure iron obtained.
Remember that depending on the actual conditions of the reaction, the actual yield may be less than the calculated theoretical yield, and the percent yield can vary accordingly. Proper laboratory technique and accurate measurements are vital for approximating theoretical yields in practice.