DETERMINATION OF THE EMPIRICAL FORMULA OF IRON OXIDE
Compounds are substances that consist of more than one element chemically bound together. The elements in a compound cannot be physically separated but can only be distinguished by the use of chemical reactions. In this lab you will create a chemical compound from two elements: iron and oxygen. The iron will come from fine steel wool and the oxygen is provided by the metabolism of plants: our atmosphere is about 20% oxygen.
The compound you will make is a common one and you probably know it by the name rust.
Rusting is a slow process that, under normal conditions, will take years. We will be speeding it up by doing two things. First, we will soak the steel wool in a solution of vinegar (6% acetic acid), which creates conditions more conducive to rapid rusting by removing the protective outer coating on the steel wool, exposing pure iron. Second, we will expose the steel wool to air for several hours, which will carry the reaction to completion.
This process will take a few hours to complete, so that this lab will extend over more than one lab period. By the end of the second day, you will have the mass of iron and the mass of iron oxide formed from it. This data will be all you need in order to determine the empirical formula of iron oxide.
1. How many moles of iron did you react?
2. How many grams of oxygen reacted with the iron and how many moles is this?
3. What is the stoichiometric ratio of iron to oxygen in the compound of iron oxide (the
chemical name for rust)?
One class of chemical reactions is called Oxidation/Reduction (redox) reactions. The prototypical oxidation reaction is one in which an element combines with oxygen; hence the name. More generally, oxidation reactions occur whenever an element or compounds loses electrons. Reduction reactions are required to separate metals from their ores and are found in such everyday products as rust removers such as Iron Out. When a reaction is classified as a reduction it means that the chemical involved has gained electrons. Oxidation reactions are always accompanied by reduction reactions. The formation of rust is classified as a redox reaction: iron loses electrons and oxygen gains electrons.
The exchange of electrons between reagents in this type of reaction is facilitated by a medium which can carry electrons. For example, batteries store potential chemical energy and the reaction can be started by completing a circuit with a wire or device. Batteries are just oxidation/reduction reactions that can be harnessed to provide an electric current. Another way to carry electrons is to carry them indirectly on ions in solution. An ionic solution is a good carrier of electricity and is therefore a perfect medium for oxidation/reduction reactions. This is the reason that you will use vinegar (a weak acid) in this lab. The hydrogen ions in the solution facilitate the exchange of electrons between the iron in the steel wool and the oxygen in the air. Without this electrolyte (a chemist’s name for a solution that carries electricity) the reaction
would take far too long. But with the acid solution we can expect all of the iron in the steel wool to rust after only a few hours.
Safety goggles, steel wool, vinegar, balance, foreceps or crucible tongs
Procedure: Day 1
1. Obtain a sample of steel wool with a mass between 0.150 g and 0.400 g. Record the
mass in the data table.
2. Place the piece of steel wool into a beaker containing 6% acetic acid. Soak the steel wool
for about 2 minutes. On a dry piece of paper towel, record your initials so that you may
identify your sample later.
3. After about two minutes have passed, remove the steel wool from the acetic acid using
forceps and blot dry with another paper towel.
4. Using the forceps, carefully loosen the steel wool so that air can interact with it. Try not
to separate any pieces completely from each other.
5. Place the paper towel with your initials on it and your steel wool sample in the hood to
oxidize and dry.
6. The vinegar may be placed down the drain and the beakers rinsed and dried. Wipe the
lab table and wash your hands.
Procedure: Day 2
1. Using forceps to transport your sample, measure and record the mass of your iron oxide.
2. When you are done with your measurement, the paper towel and iron oxide may be
placed in the trash.
Mass of iron (g)
Mass of iron oxide (g)
Mass of oxygen (g)
Analysis (Answer the following on a separate sheet of paper and attach)
1. Calculate the number of moles of iron that reacted.
2. Calculate the mass and number of moles of oxygen atoms that reacted. (In this case it is
appropriate to use the chemical formula O and not O. Why?) 2
3. Calculate the ratio of iron atoms to oxygen atoms in the formula by dividing the moles of
each element by the smaller number of moles. This will result in a formula with
subscripts greater or equal to 1. If the numbers are very close to whole number integers,
then this is also the empirical formula. Sometimes, however, half- or quarter-numbers
are obtained and in order to obtain the smallest ratio of WHOLE number subscripts (i.e.
the empirical formula), you must multiply all the subscripts by the same number to get
the empirical formula.
4. Find the smallest whole-number ratio of the compound by multiplying the ratio by
integers until all the subscripts are WHOLE numbers.
Use the following as an example:
take a reaction of element X with element Y
moles of X = 2.42
moles of Y = 6.46
XY2.426.46 2.42 2.42
(XY)x 2 ; XY[not quite there yet…] 12.6625.36
(XY)x 3 ; XY ? XY12.6637.98 38
so the molar ratio is 8:3 and the formula is XY 38
5. Determine the empirical formula of the compound. Using a reference book find out
whether your formula is correct.
6. Using the correct formulas for elemental iron and oxygen write the balanced chemical
equation for the reaction that produced your product.
7. In your experimental set-up, which chemical was the limiting reagent? How did you use
this fact to your advantage when calculating the empirical formula of the product? Discussion & Conclusion
Explain any theories, laws, or major concepts that this lab illustrated for you. Also discuss your accuracy in determining the empirical formula of iron (III) oxide (rust). Be sure to include
possible sources of error and how they would have affected your experimental empirical formula. For example, if some of the steel wool was left on the paper towel when blotting it dry after soaking in vinegar, how would this have affected your empirical formula determination?