# Last edited on 2016-12-15 18:17:22 by stolfilocal

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Introduction

Copper(II) Chloride CuCl2 is a salt with some unusual properties. In
cold pure dilute solutions it is sky blue, like other more boring salts
of copper. In those conditions, the copper(II) ions (Cu++) are
surrounded by, and weakly bonded to, 8 molecules of water, forming the
charged complex Cu(H2O)6++; which is responsible for the blue color of
most copper salts in solution.

However, when hot or containing an excess of chloride ions, solutions of
CuCl2 becomes lime green; in more extreme conditions, they become yellow
and then rusty brown. This is atributed to the replacement of water by
chloride ions in the envelope of the copper ions, forming the negative
ions Cu2Cl6-- and Cu2Cl4-- instead of Cu(H2O)6++.

The solid too is rusty brown. This is a typical color for Copper(I)
salts, but is exceptional among the salts of copper(II).

Goals and methods

In this experiment we will make some copper(II) chloride from two very
common chemicals, copper(II) sulfate CuSO4 and calcium chloride CaCl2.

The plan exploits the fact that CuCl2, CaCl2, and CuSO4 are all quite
soluble in water, whereas CaSO4 is almost insoluble. Therefore, mixing
stoichometric solutions of CuSO4 and CaCl2 should result in a solution
of CuCl2 and a precipitate or CaSO4. This is a very simple type of
reaction, the staple of most chemistry sets.

After filtering out the CaSO4, our plan is to remove all the water by
evaporation, obtaining a solid that is mostly CuCl2. We then purify this
compound by exploiting the fact that it is soluble in ethanol (common
alcohol), wereas the main expected impurities are not.

The reagents and the other products

Copper sulfate is commonly used in agriculture as fungicide. It is also
a popular material for hobbyists as it easily produces large crystals of
an intense blue color.

Massive amounts of calcium chloride are used to melt ice on roads and
sidewalks and to control dust on dirt roads. It is also a household
chemical, used to absorb moisture from closets and in the preparation of
some foods.

Calcium sulfate, the solid byproduct, is also happens to be a very
common substance, known as gypsum or plaster of Paris. It is used to
make stucco decoration, drywall boards, molds and casts. We will not try
to use this byproduct, since it will be contaminated with small amounts
of the other compounds.

Hydrates

To get the right amounts of the reagents, and check the mass of the
result, we must take into account the fact that the solid forms of these
substances contain some water, loosely bound to the metal ion and/or
trapped in the crystal structure.

Solid calcium chloride is commonly available as a hydrate of variable
composition CaCl2(H2O){x}, where {x} can be 0, 1, 2, 4, or 6. The
transition from 6 to 4 happens at 30 C, and the transition from 4 to 2
happens at 60 C. In both transitions, the excluded water dissolves part
of the solid. (Below 30 C the hxahydrate continues to absorbs water from
the air to produce a concentrated solution.)

Copper sulfate is commonly available in solid form as the pentahydrate
CuSO4(H2O)5. There are two other hydrates with 3 and 1 water molecules
per CuSO4 unit, besides the anhydrous form; but they are formed only by
heating the pentahydrate above 100 C.

Note that hydration water contributes to the mass of the solid, and
contributes some water to the solution. That is, when dissolving 10
grams of CaCl2(H2O)6 in 100 ml of water, we are actually dissolving a
bit less than 10 grams of CaCl2 in a bit more than 100 ml of water.
Thus, when weighting a desired amount of a solid substance, one must
know which hydrate it is.

The calcium sulfate that precipitates below 100 C is the dihydrate
CaSO4(H2O)2, the spent form of plaster of Paris. (It loses some of its
water only when heated well above 100 C, when it becomes the
hemihydrate, which is plaster of Paris before use.) Note that the
precipitation of the dihydrate will remove some of the water from the
solution.

Copper(II) chloride has a dyhidrate, CuCl2(H2O)2, that loses the water
slightly above 100 C to become the anhydrous salt.

Reagent amounts

We arbitrarily choose to make 50 g of CuCl2. That will be enough for
other experiments that we plan to make later on. From that we can
compute how much of each reagent we need.

First, we convert those 50 g to mols of the formula CuCl2:

  [ Recipe ]

To make one mol of CuCl2 we need one mol of CuSO4 and one mol of CaCl2.
Converting back to grams, we get:

 [ Recipe ]

That will make also 1 mol of CaSO4.

Water amount

We will use the minimum amount of water needed to dissolve both the
reagents and all the copper chloride produced. The less water we use,
the less we will need to evaporate at the end.

It turns out that the water from the two solutions that remains in
liquid form will always be sufficient to dissolve all the CuCl2 formed.
Therefore, we will use as little water as possible in the two starting
solutions. Namely, we will use saturated solution of the two salts.

Preparing the calcium chloride solution

Since we do not know which hydrate of CaCl2 we have, we do not know how
much of the solid to use. We will get around this problem by exploting
the fact that all hydrates give the same solution, and the saturated
solution contains a know concentration of CaCl2.

Namely, we choose some temperature {Ts}, say 20 C, and look up the
solubility of CaCl2 at that temperature. That is, how many grams of
anhydrous CaCl2 can be dissolved in 100 ml (or 100 g) of water at that
temperature:

[ Recipe ]

(Note that the solubility value depends on which hydrate was used by
whoever measured it. If he did not use the anhydrous form, the number
must be adjusted.)

Then, from that solubility value, we work out how many grams of
saturated CaCl2 solution we need to use, in order to get the desired
amount of CaCl2:

 [ Recipe ]

Then we put a large enough amount of water in a beaker, at that chosen
temperature {Ts}, and we add CaCl2, until some of it will remain solid
even after lengthy stirring.

Then we will decant the right amount (by mass) of the clear solution
into a separate beaker, carefully leaving the remaining solid behind.

By this procedure, we know that the first beaker will contain the right
amount of CaCl2, no matter what hydrate we have.

   [ Recipe ]

Preparing the copper sulfate solution

The copper sulfate one can get in stores is almost always the
pentahydrate CuSO4(H2O)5, and the other hydrates do not form
spontaneously. Thus, the problem that we had with CaCl2 does not arise
here. We could prepare the CuSO4 solution by weighting the desired
amount of CuSO4(H2O)5. and adding as much water as needed to dissolve
it.

However, for this video we have chosen to use the same method we used
for the calcium chloride solution. Namely, we prepare a saturated
solution of CuSO4 in water at the chosen temperature {Ts}, and we
separate the right amount (by mass) of the liquid. Redoing the same
computations we did for CaCl2, we get

[ Recipe ]

We will use a little excess of this solution, for reasons to be
explained later.

Reaction

After weighting the desired amounts of the two saturated solutions, we
mix them by pouring in a large beaker. To prevent the salts from
crystallizing out before the mixing, the solutions can be heated at some
higher temperature, and/or diluted with some water. Obviously, that
would not affect the amount of CaCl2 and CuSO4 in the mix.

On mixing, a white precipitate of CaSO4(H2O)2 is formed immediately.
Depending on the amount of water present, the result may be a paste.
Heating the mixture to near boiling and letting it cool may help
increase the sizes of the particles (by dissolving and re-crystallizing
a bit of the calcium sulfate). That will speed up filtering.

Filtering

The mixture is then filtered through some coarse paper, such as a coffee
filter (or even a couple of sheets of paper towels, if they don't break
down when wet).

Most of the precipitate should remain on the paper, and a relatively
clear solution (the "main filtrate") should go through, containing
mainly CuCl2. (A bit of precipitate may get through, but that is no
problem; we will take care of it later.) The solution should be some
color between green and blue, depending on the concentration of the
solutions (that is, on the tepmerature {Ts} at which they were
prepared).

The precipitate on the filter should be washed with a bit of water
squirted over it and deep into it. The filtrate of this washing, that
contains significant amounts of CuCl2, should be added to the "main
filtrate" of the original mixture.

This washing should be repeated a couple of times, until the wash water
is a very light blue-green and the precipitate is mostly white all
through. We save the precipitate for later.

The main filtrate (including the water from the first washes) should
then be run again through a real filter paper, in order to remove the
small amounts of precipitate that it still contains. Again we wash the
precipitate a couple of times, with a bit of water, adding the wash
water to the "main filtrate". We should then have a few hundred ml of a
perfecly clear blue-green solution.

The precipitate

While our goal is the copper chloride, we will try to save the calcium
sulfate too, for future experiments. It should be more pure than the
gypsum that one can buy at hardware store, that may contain the cheaper
calcium carbonate, sand and silcates from rocks, etc.

So, we put all the precipitate (from both filtering steps) into a single
beaker, scraping and washing as much as we can off the filter papers. We
add about 100 ml of water, mix thoroughly, and repeat the coarse and
fine filtering steps.

In both steps we wash the precipitate several times while on the filter
paper. This time, however, we keep all the light blue filtrate in a
separate flask of "wash solution". If we added it to the main filtrate,
it add to it very little CuCl2 but a lot of water that would have to be
boiled off.

Once the precipitate is throughly washed snow-white, we scrape and wash
it off the filter paper. To remove the last amounts of CuCl2 from the
the precipitate, we spread it over a filter paper and several sheets of
towel paper. When it is firm we squirt a bit of water on top of it, to
wash it even more completey, changing the towels (which must be
discarded). finally we cover it with another disk of filter paper, more
paper towels, and some non-metallic waterproof weight, such as a plastic
container or a ceramic tile.

After an hour or so we remove the paper towels, leaving the two filter
papers, and we let the precipitate dry out completely in air for a day
or two. In the end we weight the precipitate and compare to the
predicted amount of CaSo4(H2O)2:

[ Recipe ]

Not bad!

Crystallization and purification

To extract the copper chloride from the main filtrate, we basically boil
off the water until only dry solid remain, and heat the solids above 110
C to turn any CuCl2(H2O)2 into anhydrous CuCl2.

Then we will redissolve the dry residue in ethanol and evaporate the
ethanol to crystallize the anhydrous CuCl2.

When the recipe is executed at 20 C, he filtrate will contain a few
grams of calcium sulfate in solution. To reduce that impurity in the
residue, we will do the evaporation in two steps. First we boil the
filtrate until just enough water is left to dissolve the CuCl2
completely at 20 C (step 1). Then we cool the solution to 20 C. That
will cause most of the CaSO4 to come out of solution. We filter again
the liquid and proceed to complete evaporation (step 2).

The brown residue should contain mostly the CuCl2, plus some excess
CuSO4, possibly hydrated, and some CaSO4 that remained in solution.

Recall that, in the basic recipe, used a small excess (5%) of copper
sulfate. One advantage of doing so is that the extra SO4-- ions will
drive a little more of the CaSO4 out of solution. However we will end up
with some CuSO4 in the solid residue of the evaporation.

Another advantage of using some excess of CuSO4 is that CuSO4 and CaSO4
are quite insoluble in ethanol, whereas CaCl2 is very soluble.
Therefore, for the purification step, it is useful to ensure that there
is no CaCl2 in the product that results from evaporation and
dehydrating.

False starts

In my first attempt, I tried to prepare the reagent solutions at 100 C
and carry out the reaction at that temperature, in order to minimize
the amount of water in the resulting CuCl2 solution. At that
temperature, only 46 and 21 ml of water are needed to dissolve the
CuSO4(H2O)5 and the CaCl2, respectively, and the final CuCl2 solution
has only 94 ml of water.

As in the recipe above, I prepared each solution separately at 100C,
with a little excess of the salt, and decanted the right mass of the
saturated liquid to a separate beaker. Since these solutions cooled a
bit after pouring, some of the salt crystallized out. I then re-heated
the two solutions until the salts dissolved again, before mixing them.

The reaction worked, but the amount of calcium sulfate was so large
compared to the amount of liquid that they formed a thick paste of a
beautful lime green color. I had to add more than 200 ml of water to
make the mixture fluid enough for the calcium sulfate to precipitate
out.

I also made a technical mistake that could have resulted in burns. When
re-heating the two solutions before mixing, I overheated the calcium
chloride one to about 120 C. (A solution of CaCl2 that is concentrated
at 100 C will boil only at 120 C or more.) Impatient, I added it to the
CuSO4 solution while still at 110 C. But a saturated solution of CuSO4
still boils at 100 C, so there was vigiorous boiling that could easily
have splattered the paste out of the beaker.