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Lab Preparation of Ferrous Sulfate

WARNING: This procedure involves working with sulfuric acid. If you choose to attempt it, you do so entirely at your own risk. 


Ferrous sulfate (FeSO4.7H2O), traditionally known as "green vitriol" or "copperas", forms beautiful blue-green crystals of the monoclinic system.  Ferrous sulfate is useful in chemistry as a reducing agent and a source of ferrous ions.  It can also act as a catalyst;  an example is Fenton's Reagent, which is used to destroy organic chemical wastes.
Ferrous sulfate's systematic name is iron (II) sulfate, with the "II" referring to the +2 valence state of iron.  Iron (III) would instead refer to the +3 or "ferric" valence state.  Despite the descriptiveness of the systematic naming convention, the terms "ferrous" and "ferric" are here to stay.  Chemists are a stubborn lot.  Actually, "ferrous" and "ferric" just sound better, in this writer's opinion.
It's important to decide how much sulfuric acid impurity is tolerable in the crystals you're going to grow;  this is most important during the crystal growing step.  The lowest pH values translate to more acid in the crystals.  Higher pH values, on the other hand, mean there will be greater oxidative loss of your ferrous sulfate.  This article assumes that H2SO4 impurities won't be a problem in the experiments you perform later, but this may not be true for every experiment. 
If for some reason you need ferrous sulfate with very low acid content, you'll have to take your chances with huge losses to oxidation by trying to grow the crystals at a higher pH (near neutral).  Such a solution can be made by adding slightly more steel wool to the acid than it can theoretically dissolve.  It may then become necessary to do the whole experiment under an inert atmosphere (nitrogen or argon).  There will probably still be traces of H2SO4 in the crystals despite this extra effort.

Crystallizing out of H2SO4
Above:  crystals of FeSO4 7H2O growing from a sulfuric acid - ferrous sulfate solution (click for larger image). 

Safety goggles MUST be worn at all times when doing the procedure or handling the compounds, even when just opening a container of something or checking on the reaction.  Wear rubber or vinyl gloves, as well;  sulfuric acid solutions are corrosive to skin, eyes, and clothing.
De-greasing the steel wool involves acetone.  You must keep this away from ignition sources, including live electrical appliances.  Acetone is similar to gasoline in that it has a very low flash point;  vapors can creep along the bench or floor and find ignition sources, flashing back into the container and causing a disaster.  Be careful not to generate static electricity.  Also, keep acetone away from strong oxidizing agents.  The work area must have adequate ventilation so fumes don't build up, but don't pull vapors through a fan whose motor could generate sparks.  This also goes for anyone working with gasoline, spray paints, lacquer thinner, etc.
Ferrous sulfate is not nearly as toxic as mercury, chromium, lead, or nickel compounds.  However, it must not be ingested or treated carelessly.  Ingestion can be fatal.  Please obtain the MSDS for ferrous sulfate;  J.T. Baker has one at this link.  (While you're at it, make sure you get and read the MSDS for sulfuric acid and the rest of your reagents).  Store ferrous sulfate in a place where children cannot get to it.  Your chemical storage cabinet(s) should be locked whenever you're not getting or putting away reagents. 
The crystals you will grow in this experiment will probably contain some sulfuric acid trapped in their crystal structure. It's best to assume that touching ferrous sulfate crystals will cause chemical burns.  Any residual H2SO4 present should not appreciably enter the air at normal temperatures;  sulfuric acid has a very low vapor pressure at 25 °C.  If the crystals are kept in a tightly sealed container, any H2SO4 that does evaporate will recondense in the container.  Open it in a well-ventilated area.   
Be careful with chemical solutions near electrical appliances such as hot plates.  The hot plate should be the sealed-element variety made for lab use.


Sulfuric Acid (H2SO4), about 30-40% concentration
Steel wool, #00 or 000 fineness
Safety goggles
Rubber or vinyl gloves
Lab apron & lab coat
Beaker, glass
Steel pan or glass dish, reserved only for chemistry
Petri dish or crystallizing dish, glass
Laboratory oven or sealed-element hot plate
Ice, crushed, in a shallow container
Funnel and filter paper; better yet, a vacuum filtration setup
pH paper or pH meter
Desiccator with calcium chloride drying pellets (optional)

Methods & Observations:
Again, keep those safety goggles on at all times.  During the acetone step, exclude all possible sources of spark, flame, or live heating element.
The size of beaker selected depends on how much ferrous sulfate you wish to make.  A good, all-around beaker size is 100 or 250 mL.

I.  Preparation of FeSO4 solution.

For best results the lab temperature should be kept below 25°C.  Avoid temperature fluctuations as much as possible, except where indicated in this procedure.
Degrease the steel wool by immersing it in acetone for half an hour.  Remove it from the acetone and let it dry in a well-ventilated lab where nobody can disturb it.  During this whole time there must be no ignition sources nearby.  Do not attempt to evaporate the acetone with a heat source.  Even an electric lamp could cause ignition. 
Place the glass beaker in the center of a metal pan or wide glass dish that's reserved just for lab use.  The reaction can produce minor spattering as hydrogen bubbles to the surface of the acid.
Place the degreased, dry steel wool in the glass beaker and pour in enough 30-40% sulfuric acid to cover it completely.  Don't use concentrated acid.  If the steel wool is not fully submerged, carefully push it down with a glass rod.
The sulfuric acid will begin to dissolve the steel, producing hydrogen gas.  Over the course of several hours the steel wool will gradually disappear.  Carefully add more steel wool.  Repeat this a few times.  Reddish-brown, insoluble ferric compounds will form if you add too much steel wool;  just add some dilute sulfuric acid if this happens.  The pH of the solution should be acidic at any given time, otherwise the ferrous ions will oxidize to the ferric state. 

II.  Filtering
Green crystals should have started forming in the solution and settling to the bottom after just a few hours.  They may become tangled up in any steel wool that's still left; don't bother trying to free these.
When you're satisfied with the amount of ferrous sulfate that has formed, add excess water that's been acidified to a pH between 2 and 4 using sulfuric acid.  Add enough of this acidified water to re-dissolve all the green crystals that have settled out.  If the solution turns brown, add just enough sulfuric acid to make it green again.
Filter this solution through filter paper.  Discard the solids and the paper;  neutralize their acidity with some dilute sodium carbonate solution.  Save the filtrate.  It should now be free of steel wool pieces, carbon (from the steel), rust, and other solids.  Keep the solution in a covered container to minimize contact with atmospheric oxygen.
If the pH is not kept low enough, ferrous sulfate will oxidize to ferric sulfate on standing.  Normally, atmospheric oxygen changes Fe2+ to Fe3+  quite readily.  This reaction is reversible, however, by lowering the pH.  Below 4 or so, ferrous ion is heavily favored over ferric, with the concentration of Fe3+ becoming vanishingly small at pH 1-2.  Aqueous FeSO4 in this pH range is stable for days, even with much air contact.

III.  Evaporation

The filtered solution is put in a shallow, glass container such as a petri dish or crystallizing dish.  Place this on the hot plate and turn it to the lowest setting;  even better, use a small, laboratory oven.   Don't use a household-type oven;  these are not safe for general lab use.
Slowly heat the solution to about 80 to 90 C; do not boil it.  Hold it at this temperature until about half the liquid has evaporated.  Don't breathe the vapors given off;  although they will be mostly water, they will contain some sulfuric acid.  The color of the solution may change to yellowish, but don't let this discourage you.  On cooling it becomes green again.
Allow the solution to cool to room temperature and place the shallow dish on crushed ice.  Leave it there for at least an hour.  Don't let it sink into the water as the ice melts.
Remove the dish from the cold and place it on the lab bench in your locked laboratory (remember, the liquid contains sulfuric acid!).  Let it stand for 24 hours.

Green crystals should form.  If not, the cause is probably one of the following:
1. You didn't evaporate the solution down far enough.
2.  The pH was too high.
3.  The ambient temperature was too high.
4.  Your sulfuric acid was contaminated with something that oxidized Fe2+.

Decant or pipette off the yellowish liquid that remains after the crystals have formed.  It is strong sulfuric acid, so be careful with it.  Use a dropper to add it slowly into a cold, dilute solution of NaOH, sodium carbonate, or ammonia.
Do not handle ferrous sulfate with your fingers.  Use a plastic or steel spatula.

Crystallizing mostly complete
Above:  Crystals to 1 cm or larger were readily attainable, starting with only a 10 mL micro beaker's volume of dilute sulfuric acid and some steel wool.

IV.  Washing

If possible, have a previously-prepared batch of ferrous sulfate crystals which  are dissolved to saturation in cold (1 to 4 °C) distilled water.  Use this solution, prepared and chilled shortly before use, to wash the crystals you've grown.
If you've prepared FeSO4 for the first time and don't have another batch, just use distilled water cooled to just above freezing.  There will be some minor losses as some of the crystals go into solution.
As the washing removes the strong sulfuric acid from the surface of the crystals, the atmosphere will unfortunately oxidize them more readily.   Larger crystals are more desirable;  less FeSO4 will oxidize.  Conversely, don't crush or powder the crystals, since it will increase the surface area.

V.  Drying; Final steps

If no desiccator is available, dry the ferrous sulfate crystals in air at the lowest relative humidity available (e.g., not outside on a foggy day).  It's preferable to do the drying in a desiccator having a some calcium chloride drying pellets in the bottom.  Obviously, the ferrous sulfate shouldn't contact the desiccant material.   Let them dry just enough to get the excess water off the crystals, then put them in a tightly-sealed container where moisture can't get to them. 
The washed and air-dried crystals of ferrous sulfate can be saved for future experiments as they are, or some / all of the moisture in the crystals can be driven off. 
Placing them in a vacuum desiccator containing some calcium chloride can dehydrate the crystals, depending on time and on vacuum strength.  It's uncertain what state of hydration this treatment leaves;  evidently, predicting this is more involved than it would seem.  The dried compound may be the monohydrate (FeSO4.H2O) or the tetrahydrate (FeSO4.4H2O), or even a mixture containing both.  Mitchell (1984) explores the subject in depth.

The desiccated ferrous sulfate
Some FeSO4.7H2O crystals were placed in a vacuum desiccator containing CaCl2 pellets as a drying agent.  A vacuum of 28 inches Hg was drawn on the desiccator, and the stopcock was then closed. 

The evacuated desiccator was left for about 10 hours.  The crystals lost their color, becoming dry and crumbly.

If ferrous sulfate is heated to dehydrate it, the heat mustn't be too strong, or they will decompose to make the toxic and corrosive sulfur trioxide (SO3).
According to The Merck Index (1983), heating FeSO4.7H2O to 90°C will leave ferrous sulfate monohydrate.  The same source (1983) lists the synonym "exsiccated ferrous sulfate" for the monohydrate;  the term appears in older texts and formularies, especially with the spelling "sulphate".  Some sulfuric acid remnants from solution may also begin to come off during heating to the monohydrate, so there must be adequate ventilation.
The monohydrate will give up its last H2O by heating to 300°C, preferably with as little air contacting the sulfate as is practical. (Never heat anything in a completely sealed vessel.)   Much of the remaining sulfuric acid bound to the crystals will also come off, so the heating must be done under a laboratory fume cupboard or outdoors.  In the latter case, wear a respirator that's rated by the manufacturer for sulfuric acid vapors, just in case the wind shifts your way.
FeSO4 decomposes at around 480°C to give off SO3 (dangerous!!).  To ensure a healthy margin of safety, avoid heating it above 350°C.


Now that we've prepared ferrous sulfate, it can be used as a reducing agent, a catalyst, an ingredient for Fenton's reagent, or a reactant for preparing a number of other compounds.  For example, mixing equimolar solutions of FeSO4 and ammonium sulfate will yield ferrous ammonium sulfate, also known as Mohr's salt.  This can be crystallized out and is much more resistant to air oxidation than FeSO4 (McGraw-Hill Encyclopedia of Chemistry, 1993).  Of course, Mohr's salt is not interchangeable with FeSO4 for every application.
Since so much of chemistry hinges directly on oxidation and reduction, it's easy to imagine how useful a reducing agent would be.  For example, ferrous ion will reduce Ag+ to metallic silver in the following manner:

Fe2+ (aq) + Ag+ (aq) = Fe3+ (aq) + Ag (s)

Some experiments will call for a solution of ferrous sulfate at a pH where there's not enough acid to stabilize it.  In other words, the FeSO4 is simply dissolved in water without adding the sulfuric acid.  This kind of solution won't last long in the presence of air, but it's a more effective reducing agent than acidic FeSO4.  If an experiment calls for the non-acidic variety, it should be prepared immediately prior to use.  Use de-gassed water, keep the solution in a closed container with very little air space, and keep it on ice.
Ferrous sulfate is incompatible with certain compounds;  see the Merck Index entry or MSDS for it.  You may read about some experiments here and there which, in apparent contradiction, call for mixing these same compounds together.  It's often a matter of the conditions.  For instance, there are well-established experiments that involve iron and silver salts together in dilute solution.  Under other circumstances, though, the resulting silver particles might pose a combustion or reactivity hazard. 
Ferrous sulfate is also incompatible with lead acetate, so keep these compounds separate if you've prepared both.  Whenever there's any doubt about chemical compatibility, just stay with the established procedures.

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Copyright:  The Society for Amateur Scientists has permission to reprint this article in full.   The article otherwise remains copyright of CR Scientific and may not be copied, reproduced, mirrored, or distributed without prior written permission (click here for contact info). 

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Works Cited:

CRC Handbook of Chemistry and Physics, 69th Edition.  Boca Raton, Florida: CRC Press, 1989.

McGraw-Hill Encyclopedia of Chemistry, 2nd Edition.  New York:  McGraw-Hill, Inc., 1993.

Merck Index, 10th Edition.  Rahway, New Jersey: Merck and Company, Inc., 1983.

Mitchell, A.G.  "The Preparation and Characterization of Ferrous Sulphate Hydrates". J. Pharm. Pharmacol. 36:506-510 (1984).

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