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 ammonium chloride, which creates conditions more conducive to rapid rusting. Second, we will heat the steel wool to dryness after several days of soaking and the heat will carry the reaction to completion.
This process will still take 4 to 7 days to complete so that this lab will extend over more than one lab period. By the end of the lab 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.
How many moles of iron did you react?
How many grams of oxygen reacted with the iron and how many moles is this?
What is the stoichiometric ratio of iron to oxygen in the compound of iron oxide?
What is the balanced chemical equation for the reaction of iron with oxygen gas to make rust?
extra fine steel wool (000 or 0000)
1 M ammonium chloride (NH4Cl)
One class of chemical reactions are called Oxidation/Reduction 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 an oxidation 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. A simple salt 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 ammonium chloride (the chloride salt of household ammonia, NH4Cl) in this lab. The salt in the solution facilitates 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 salt solution we can expect all of the iron in the steel wool to rust by the time a week has passed. The advantage of NH4Cl is that it turns directly from a solid to a gas (it sublimes) when heated. This will allow us to remove it from the iron oxide in the final procedure of the lab.
The following list does not cover all possible hazards, just the ones that can be anticipated. Move slowly and carefully in the lab: haste and impatience have caused more than one accident.
Ammonium chloride solution is mildy corrosive. Keep it off your skin and clothing and away from your eyes. Wash away any spills with plenty of water.
When you heat something it looks the same as when it was cool. Do not touch anything you have heated over the burner until you are sure it has cooled off. Avoid a serious burn by being careful!
Use a wet paper towel to test hot objects to see whether they are cool enough to touch.
Stay away from the fumes that rise from the crucible when you heat it. This 'smoke' is ammonium chloride and it is a strong irritant. Heat in a fume hood or in a well-ventilated room.
There are three parts to the procedure for this lab. First, you must find the mass of iron and prepare it for a week of rusting. Second, you must monitor the iron during that week and ensure that it stays damp with NH4Cl solution. Third, and finally, you must drive off the remaining NH4Cl solution and complete the rusting of the iron by heating and find the mass of iron oxide.
Find the mass of your crucible.
Measure out a mass of about 3 g of steel wool. Record the actual mass of steel wool that you use.
Compress the steel wool into a tight ball and put it in your crucible. Find and record the mass of the combination of the two.
When you are ready, take the crucible and steel wool to the beaker of 1 MNH4Cl that your instructor has prepared. Dunk the steel wool in the solution using your tongs until thoroughly wet (about 5 - 10 s).
Allow the solution to drip off into the beaker and replace the steel wool in the crucible.
Make a labelled notecard on which to set your crucible and put it in the designated storage area.
Over the next week of classes look at your steel wool once a day and add more NH4Cl solution as needed to keep it damp.
Record any changes as time progresses.
On the final day of this lab set up the apparatus as displayed by your teacher and as depicted at right. Set up a ring stand to support a ring with clay triangle and prepare your burner to heat the crucible.
Wait until your teacher improves air circulation by setting up one or more fans in the windows to blow air out of the room. Heating the crucibles sublimes the ammonium chloride and sends it into the air as a fine smoke.
Heat your crucible with its now very rusty contents using the burner at high heat. The crucible will glow red hot during the heating: continue heating for at least ten minutes. Stand back and do not look down into the crucible during this process as the fumes can easily get in your eyes or lungs. It is easier to see the rising white smoke if you move the burner out from under the crucible.
Turn off the burner and allow the crucible to cool until you can touch it. This may take some time. Test the temperature of the apparatus using a wet paper towel.
Find the mass of the crucible plus the dried out rust and record it.
Heat again to red hot and allow to cool. Weigh again. If the two masses match then you are finished. Continue this process until you get the same mass twice. This is called heating to constant mass.
The rust can be disposed of in the trash or kept for further experiments. Be sure to crush it up and admire its rich red-brown color.
Clean up your lab station and scrub out the crucible. Be sure to wash your hands thoroughly before taking off your goggles and leaving the lab.
Now you will use your data to determine the empirical formula of rust. You should collect your data into a neat table showing the mass of the crucible, the mass of steel wool (iron), the mass of the crucible plus iron oxide, and the mass of iron oxide. Additionally, you should add the calculated mass of oxygen which was added to the mass in the crucible and the number of moles of each reagent: iron and oxygen.
Answer these questions using complete sentences. Show work for all calculations.
Calculate the number of moles of iron that reacted.
Calculate the mass of oxygen that reacted by subtracting the mass of iron you started with from the final mass of the rust.
Calculate number of moles of oxygen atoms that reacted. (In this case it is appropriate to use the chemical formula O and not O2.)
Calculate the ratio of oxygen atoms to iron atoms in the formula by dividing the moles of oxygen by the moles of iron. This will result in a ratio of this form: O/Fe = x/1. The number will not be a whole number (more than likely) and the ratio will express the number of moles of oxygen which combines with one mole of iron atoms.
Find the smallest whole-number ratio of the compound by multiplying the ratio by integers until you come up with a whole number. Your teacher can demonstrate this technique for you.
Use the following as a demonstration:
take a reaction of element X with element Y
moles of X 2.42
moles of Y 6.46
Y 2.67 5.33 8.01
- = ---- = ---- = -----
X 1 2 3
so the molar ratio is 8:3 and the formula is X3Y8
Determine the empirical formula of the compound. Using a reference book or the internet find out whether your formula is correct for iron(III) oxide.
Using the correct formulas for elemental iron and oxygen write the balanced chemical equation for the reaction that produced your product.
It is possible in this procedure to obtain data which give a ratio of O/Felower than the true ratio for iron(III) oxide. What physical casuse would produce data that give such a result? Do not just mention ‘human error’ and instead give specific reasons why the ratio was either too low between oxygen and iron.
It is also possible that the O/Fe ratio is too large. What actual events in the lab could lead to such a result?
In your experimental set-up which chemical was the limiting reactant? Explain.
What is the importance of using ammonium chloride as the electrolyte which allowed the iron to rust quickly? Why would sodium chloride be a poor choice given the rest of the experimental procedure?
In a professional-quality typed document answer the questions above. By typing them show work for all calculations necessary to answer the questions.