Chapter 6 Laboratory: Separating Mixtures 103
FIGURE 6-3:
Recrystallized copper sulfate
funnel, trying to make sure it wets all of the crystals in the
filter paper.
Carefully remove the filter paper and set the purified
copper sulfate crystals aside to dry.
When the filter paper and copper sulfate is completely
dry, weigh it and record the mass of the filter paper and
copper sulfate crystals in Table 6-3.
Transfer the purified copper sulfate crystals to
a labeled bottle, and retain them for use in later
laboratory sessions.
At this point, we have (assuming no losses) about 200 g of
pure copper sulfate crystals on the filter paper and about 50 g
of crude copper sulfate dissolved in about 150 mL of solution.
Rather than waste that 50 g of crude copper sulfate, we’ll use it to
produce some crude copper carbonate, which we’ll use in a later
laboratory session.
Copper sulfate is fairly soluble in water, as are sodium carbonate
and sodium sulfate, but copper carbonate is extremely insoluble
in water. We’ll take advantage of these differential solubilities
to produce, separate, and purify copper carbonate. By reacting
the waste copper sulfate solution with a solution of sodium
carbonate, we precipitate nearly all of the copper ions as the
insoluble carbonate salt, leaving sodium sulfate in solution.
- Weigh about 50 g of sodium carbonate heptahydrate
and transfer it to the empty 250 mL beaker.
- Use the graduated cylinder to measure about 100 mL
of warm tap water and transfer it to the beaker with
the sodium carbonate. Stir until the sodium carbonate
dissolves.
- Pour the 100 mL of sodium carbonate solution into the
beaker of copper sulfate solution, with stirring.
- Put a fresh piece of fan-folded filter paper into the filter
funnel and place the empty beaker beneath the funnel.
Swirl the contents of the beaker to keep the precipitate
suspended in the solution and pour the solution through
the filter paper.
- Rinse the precipitate three times with about 50 mL
of water each time. This step removes nearly all of the
sodium sulfate from the precipitate, as well as any
other soluble salts formed by trace contaminants in the
original waste solution.
- Carefully remove the filter paper and set the copper
carbonate aside to dry. If you have sufficient acetone,
you can do a final rinse with 25 mL to 50 mL of acetone
to remove nearly all of the water and allow the copper
carbonate to dry faster.
- Transfer the copper carbonate to a labeled bottle, and
retain it for use in a later laboratory session.
FIGURE 6-4:
Precipitated copper carbonate
SoUBILITL y vERSUS SoLvENT TEmpERATURE
Most solid water-soluble compounds are more soluble in
hot water than in cold. Gases and a few solid compounds
are more soluble in cold water than in hot water, a
phenomenon called retrograde solubility. For those
solid compounds, it’s possible to do a recrystallization
in reverse. Instead of making a saturated solution of the
compound in hot water and then cooling the solution to
cause crystals to form, you make a saturated solution of
the compound in cold water, and then heat the solution to
cause crystals to form.
The change in solubility with solvent temperature differs
from compound to compound. Some compounds, such
as the copper sulfate we use in this lab, are much more
soluble in hot water than in cold. But the solubility of many
compounds is little affected by solvent temperature.
For example, the solubility of sodium chloride in water is
355 g/L at 0°C, increasing only about 10% to 390 g/L at
100°C. Recrystallization is inefficient for purifying such
compounds, because most of the compound remains in
solution regardless of the temperature.
wHENC RUdE ISN’T
Ironically, the “crude” copper sulfate crystals we bought
turned out to be quite pure. The label included an assay
that listed “copper sulfate — 99.0%,” which is nearly
reagent grade purity. Dr. Paul Jones, one of my technical
reviewers, is a chemistry professor. I showed him the
two-pound bottle of copper sulfate I’d bought at a local
big-box DIY store. He examined the label, commented on
its high purity, and asked, “How much did this cost?” When
I told him I paid about $7.50 for the two-pound bottle and
added that at the same time I’d also paid about $5.00 for
a two-pound bottle of sodium hydroxide with an assay
of 100.0%, Paul commented, “I may have to start buying
some of my chemicals there.” Indeed.
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