In this lab, you will perform the replacement reaction as shown below. By weighing the iron that is added to the copper(II) solution, and weighing the copper produced in the reaction, you will be able to determine whether iron is oxidized to iron(II) or iron(III).
Cu2+(aq) + Fe(s) -----> Fe+?(aq) + Cu(s)
In your introduction, you should state the reaction and identify what is oxidized and what is reduced in the reaction. You should then write the two balanced equations corresponding to the two possibilities (iron(II) and iron(III)). Then, assuming you start with 2.00 g iron, calculate the mass of copper produced in each case.
Obtain about 50 mL copper(II) sulfate solution in a 250 mL beaker. Weigh out the iron and place in the copper sulfate solution. Stir with a glass rod, observing any changes in the appearance of the contents of the beaker and recording them in your lab book. Weigh a piece of filter paper and record the weight in your lab book. When all the iron has reacted, filter the copper and rinse thoroughly with distilled water. Let dry overnight and mass the filter paper and beaker.
¥determine the # of moles of copper expected for both possibilities (ferrous and ferric) as you did for the prelab
¥determine the # of moles of copper actually produced (don't forget to subtract the mass of the filter paper!)
Which iron ion was produced? Justify. Compare your results with another groups. Do their results agree with yours? Does this reaction agree with what you would expect from the electromotive series?
What were some sources of error in this lab? How could you improve this experiment?