Lab Preparation of Ferrous Sulfate
WARNING: This procedure
involves working with sulfuric acid.
choose to attempt it, you do so entirely at your
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.
sulfate's systematic name is iron (II) sulfate, with the "II" referring
to the +2 valence state of iron. Iron (III) would instead refer
+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.
important to decide how much sulfuric acid
impurity is tolerable in the crystals you're going to grow; this
most important during the crystal growing step. The lowest
values translate to more acid in the crystals.
Higher pH values, on the other hand, mean there will be greater
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.
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.
FeSO4 • 7H2O growing from a
sulfuric acid -
ferrous sulfate solution (click for larger image).
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
acid solutions are corrosive to
skin, eyes, and clothing.
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.
sulfate is not nearly as toxic as mercury, chromium, lead, or
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
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.
crystals you will grow in this experiment will probably contain
sulfuric acid trapped in their crystal structure. It's best to assume
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
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
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
Rubber or vinyl gloves
Lab apron & lab coat
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:
those safety goggles on at all times. During the acetone step,
exclude all possible sources of spark, flame, or live heating element.
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
best results the lab temperature should be kept below 25°C.
temperature fluctuations as much as possible, except where indicated in
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
time there must be no ignition
sources nearby. Do not
to evaporate the acetone with a heat source. Even an electric
could cause ignition.
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.
the degreased, dry steel wool in the glass beaker and pour in enough
acid to cover it completely. Don't use concentrated acid.
If the steel wool is not fully submerged, carefully push it down with a
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
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.
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.
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
have settled out. If the solution turns brown, add just enough
sulfuric acid to make it green again.
this solution through filter paper. Discard the solids and the
paper; neutralize their acidity with some dilute sodium carbonate
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
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
filtered solution is put in a shallow, glass container such as a petri
crystallizing dish. Place this on the hot plate and turn it to
lowest setting; even better, use a small, laboratory
oven. Don't use a household-type oven; these are not
safe for general lab use.
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
be mostly water, they will contain some sulfuric acid. The color
the solution may change to yellowish, but don't let this
discourage you. On cooling it becomes green again.
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.
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.
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
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.
not handle ferrous sulfate with your fingers. Use a plastic or
1 cm or larger were readily attainable, starting with only a 10 mL
beaker's volume of dilute sulfuric acid and some steel wool.
possible, have a previously-prepared batch of ferrous sulfate
crystals which are dissolved to saturation in cold (1 to 4
water. Use this solution, prepared and chilled shortly before
use, to wash the crystals you've
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.
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
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
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.
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.
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
or the tetrahydrate (FeSO4.4H2O),
or even a mixture containing both. Mitchell (1984) explores the
subject in depth.
|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
and the stopcock was then closed.
The evacuated desiccator was left for about 10 hours. The
crystals lost their color, becoming dry and crumbly.
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
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
acid remnants from solution may also begin to come off during heating
to the monohydrate, so there must
be adequate ventilation.
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
vapors, just in case the
wind shifts your way.
decomposes at around 480°C to give off SO3
(dangerous!!). To ensure a healthy margin of safety, avoid
heating it above 350°C.
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
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.
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
the following manner:
Fe2+ (aq) + Ag+
(aq) = Fe3+ (aq) + Ag (s)
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.
of solution won't last long in the presence of air, but it's a more
effective reducing agent than acidic FeSO4. If an
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.
sulfate is incompatible with
certain compounds; see the Merck
Index entry or MSDS for it. You may read about some
experiments here and
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
combustion or reactivity hazard.
sulfate is also incompatible with lead acetate, so
these compounds separate if you've prepared both. Whenever
there's any doubt about chemical compatibility, just stay with the
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Legal Notice: In order to use this website
or any information
CRC Handbook of Chemistry and
Physics, 69th Edition. Boca Raton, Florida: CRC Press,
McGraw-Hill Encyclopedia of
Chemistry, 2nd Edition. New York: McGraw-Hill, Inc.,
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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