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This article incorporates, in modified form, material from Illustrated Guide to Home Chemistry Experiments: All Lab, No Lecture.


In this lab, we’ll use a procedure called recrystallization to purify crude copper sulfate. Crude copper sulfate is a mixture of copper sulfate with various impurities that may include copper carbonate, copper oxides, and other copper compounds. You can obtain crude copper sulfate at a hardware or lawn and garden store where it’s sold as a root killer or pond treatment.


Many chemicals besides copper sulfate are available inexpensively in impure forms, such as technical and practical grades. Although these chemicals may be insufficiently pure for general lab use, many of them can be purified to the equivalent of lab grade, or even reagent grade, by using recrystallization.


When Crude 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. 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 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.

Mary Chervenak comments:

Paul has also started purchasing sodium bicarbonate from the grocery store, because it’s much, much cheaper and just as pure (if not purer, since it’s routinely used in baking) than the containers he buys from his usual laboratory chemical suppliers!

Successful recrystallization depends on two factors. First, crystals are the purest form of a chemical, because there’s no room for impurities in the crystalline lattice. As crystals grow in a solution of an impure chemical, the impurities remain in solution. Second, some chemicals are much more soluble in hot water than in cold. For example, the solubility of copper sulfate pentahydrate, the chemical we’ll purify in this lab, is 203.3g/100mL at 100 °C, versus only 31.6g/100mL at 0 °C.

In other words, a saturated solution of copper sulfate at 100 °C contains about six times more copper sulfate than the same amount of solution at 0 °C. If we saturate boiling water with impure copper sulfate and then cool that solution to 0 °C, about 5/6 of the copper sulfate crystallizes in pure form, leaving the impurities in solution (along with about 1/6 of the original copper sulfate).

A supersaturated solution is one that contains more solute per volume of solution than the solution would ordinarily contain at that temperature. Supersaturation normally occurs when the temperature of a saturated solution is gradually reduced. Rather than precipitating out, the excess solute remains in solution until some event occurs that causes the excess solute to precipitate suddenly. That event may be as minor as dropping a tiny crystal of the solute into the solution or even simply tapping the container. This is called nucleation. Crystals produced from a supersaturated solution are not necessarily pure because the rapid crystallization may trap impurities within the crystal structure.

Solubility 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.

Once the crystals have formed, they can be separated from the supernatant liquid by filtration or simply by decanting off the liquid. In either case, some of the contaminated liquid remains mixed with the purified copper sulfate crystals. You remove this last bit of contamination by rinsing the crystals with a small amount of ice-cold water or another solvent that is miscible with water, such as acetone, leaving only purified copper sulfate crystals.

We’ll use acetone for the final rinse, because copper sulfate is almost insoluble in acetone. If we used ice-cold water, some of our purified copper sulfate would dissolve in the rinse water, lowering our final yield. Using acetone flushes away impurities without dissolving the copper sulfate crystals. Volatile organic solvents like acetone also evaporate faster and more completely than water.

Required Equipment and Supplies

  • goggles, gloves, and protective clothing
  • balance and weighing papers (optional, but recommended)
  • hotplate or kitchen stove burner
  • beaker, 250 mL
  • cylinder, graduated, 100 mL
  • flask, Erlenmeyer, 250 mL
  • stirring rod
  • funnel
  • filter paper (or bleached white coffee filters)
  • kitchen tongs or oven mitt(s)
  • freezer
  • copper sulfate pentahydrate (~ 100 g crude crystals)
  • acetone (a few mL)
  • sodium carbonate heptahydrate (~ 20 g)

All of the specialty lab equipment and chemicals needed for this and other
lab sessions are available individually from Maker Shed or other laboratory
supplies vendors. Maker Shed also offers customized laboratory kits at special
prices, including the Basic Laboratory Equipment Kit, the Laboratory Hardware Kit, the Volumetric Glassware Kit, the Core Chemicals Kit, and the
Supplemental Chemicals Kit.


Copper sulfate is moderately toxic. Handle the hot liquids used in this experiment with extreme care. Acetone is flammable. Use extreme care, avoid open flames, and have a fire extinguisher handy. Wear splash goggles, gloves, and protective clothing.

Substitutions and Modifications

  • If you do not have a balance, you can obtain a reasonably accurate mass of fine copper sulfate pentahydrate crystals using a kitchen measuring tablespoon. We made five separate weighings of one level tablespoon (15 mL) of fine copper sulfate crystals and observed an average mass of 19.01 g per tablespoon, or about 1.27 g/mL. That density is considerably lower than the known actual density of 2.28 g/mL for copper(II) sulfate pentahydrate. That difference exists because there is considerable air space separating the crystals in the measuring spoon. This volumetric method is usable only if your copper(II) sulfate pentahydrate is in the form of fine crystals. If you have large crystals or chunks, you can still recrystallize them, albeit non-quantitatively, by adding copper(II) sulfate to a volume of boiling water until no more will dissolve (note that copper(II) sulfate dissolves slowly, so be patient to make sure the solution is in fact saturated). Decant that hot saturated solution into a second beaker, add water (about 20% of the volume you started with) and again bring the solution to a boil. We don’t want a truly saturated solution of copper(II) sulfate, because a boiling saturated solution will start to crystallize out as soon as you pour it into the funnel. Adding some water and reheating to a boil leaves the solution sufficiently unsaturated to allow it to be filtered without premature crystallization.
  • If your hotplate has heating coil rather than a flat heating surface, use a piece of metal screen, a large tin can lid, or a burner cover between the coil and beaker rather than putting the beaker in direct contact with the burner.
  • The size of funnel (and filter paper) you need depends on the scale of the experiment. When I ran this experiment while writing the book, I actually purified 250 g of copper(II) sulfate in one pass, so I needed a relatively large funnel to capture all of the crystals in one filtration. I used a large, heat-resistant plastic funnel from Wal*Mart, with a bleached white coffee filter instead of actual filter paper.
  • The quantity of copper sulfate used in this laboratory assumes that you really do want to purify a few ounces of copper sulfate to use in the later labs in this book. If you want only to demonstrate the principle of recrystallization, you can reduce the quantity of copper sulfate (and water) and the size of the containers proportionately. A few grams of copper(II) sulfate in a half test tube of water works well to illustrate the principle of recrystallization.
  • You can substitute 20 g of crude washing soda or 9.2 g of anhydrous sodium carbonate (soda ash) for the 20 g of sodium carbonate heptahydrate. If you alter the mass of copper(II) sulfate, you can alter the mass of sodium carbonate proportionately. Actually, you don’t need to know the mass of the sodium carbonate, which is used to precipitate out excess copper(II) sulfate as copper(II) carbonate. You can simply add sufficient sodium carbonate to the waste solution of copper(II) sulfate to precipitate out all of the copper(II) ions, leaving the supernatant liquid colorless. Any excess sodium carbonate (along with the sodium sulfate formed in the reaction that precipitates the copper(II) carbonate) will be washed away when you wash the insoluble copper(II) carbonate.
  • We will cool the solution in the freezer, which raises an obvious safety issue. Make sure when you place the flask with the hot copper(II) sulfate solution in the freezer that it does not spill or otherwise contaminate any of the food items in the freezer. Label the solution to make absolutely certain it cannot be confused with food items. (This should go without saying; labeling containers is good lab technique, and you should never store any product in an unlabeled container.) To be completely safe, you can cool the flask in a coffee can or other container of appropriate size using ice. Place a bed of crushed ice or small chunks of ice on the bottom of the larger container, place the flask on top of the ice, and then surround the flask with additional ice and allow it to cool.


  1. If you have not already done so, put on your splash goggles, gloves, and protective clothing.
  2. Place a weighing paper on the balance pan and tare the balance to read 0.00 g. Add crude copper sulfate pentahydrate until the balance indicates about 100 g. Record the mass of the copper sulfate to 0.01 g in Table 6-3. If the maximum capacity of your balance is less than 100 g, weigh multiple samples until you accumulate about 100 g total. Transfer the copper sulfate to the 250 mL beaker.
  3. Set your balance to read 0.00 g and weigh a piece of filter paper. Record the mass of the filter paper to 0.01 g in Table 6-3, and use a pencil to write the mass on the filter paper itself.
  4. Fan-fold (flute) the massed filter paper and set it aside for later use.
  5. Set up your filter funnel with another piece of fan-folded filter paper over the 250 mL flask.
  6. Use the 100 mL graduated cylinder to measure 60 mL of hot tap water, and add it to the copper sulfate in the 250 mL beaker.
  7. Place the beaker on the hotplate and bring the water to a gentle boil. Stir the solution until all of the crude copper sulfate has dissolved. Any solid matter that refuses to dissolve is an insoluble impurity and can be ignored.

Paul Jones comments

I wouldn’t say ignored. As described in the following step, any insoluble residue should be trapped via hot filtration and discarded. Don’t add more water to try to dissolve it. Also, hot filtration works best if you heat the receiving vessel up prior to filtration. Use a little boiling water to heat the filter paper and funnel. This reduces crystallization on the paper.

  1. Using the tongs or oven mitts to protect your hands, pour the hot copper sulfate solution through the filter paper as quickly as possible. Some copper sulfate may recrystallize when it contacts the cooler filter paper, funnel, or receiving flask, but it should go back into solution as you continue to pour the hot solution through the funnel. When the beaker is empty, allow it to cool, rinse it thoroughly, and dry it.
  2. 9. As the solution in the flask cools, copper sulfate will begin crystallizing. Slow cooling produces larger crystals, and fast cooling smaller crystals. Smaller crystals are easier to handle and weigh, so we want to produce crystals as small as possible. Our goal, then, is to cool the solution as fast as possible from the boiling point to near the freezing point. To do that, place the flask of hot solution in the freezer or an ice bath.

Paul Jones comments

Yes, but it’s much cooler to grow big crystals by letting the solution sit to cool slowly without stirring, and then you get to watch them grow. You could even shoot a time-lapse video.

  1. As the solution cools, stir it periodically. While you’re waiting for the solution to cool completely, remove and discard the used filter paper from the filter funnel, rinse the funnel clean, and put in the fan-folded filter paper that you weighed earlier.
  2. When the solution has cooled completely, use the stirring rod to scrape free any crystals that have grown on the sides and bottom of the flask. Swirl the contents of the flask gently to keep the crystals suspended in the solution, and quickly pour the solution and suspended crystals into the 250 mL beaker. If some crystals remain in the flask, decant a bit of the liquid from the beaker back into the flask, swirl to suspend the crystals, and quickly pour the solution and crystals back into the beaker.
  3. Rinse and drain the flask. Set the funnel with the pre-weighed filter paper on top of the flask, swirl the contents of the beaker to suspend the crystals, and pour the liquid and crystals through the filter funnel. (If the funnel forms a tight seal with the lip of the flask, filtration may proceed very slowly. You can speed things up by placing a small wedge of paper or a paperclip between the funnel and the flask to break the seal.)
  4. Use a few mL of acetone to rinse any remaining crystals from the beaker and pour the acetone through the filter funnel, trying to make sure it wets all of the crystals in the filter paper.
  5. Carefully remove the filter paper and set the purified copper sulfate crystals aside to dry.
  6. 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.
  7. Transfer the purified copper sulfate crystals to a labeled bottle, and retain them for use in later laboratory sessions.

Figure 6-3

Figure 6-3. Recrystallized copper sulfate

At this point, we have (assuming no losses) about 80 g of pure copper sulfate crystals on the filter paper and about 20 g of crude copper sulfate dissolved in about 60 mL of solution. Rather than waste that 20 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.

  1. Weigh about 20 g of sodium carbonate heptahydrate and transfer it to the empty 250 mL beaker.
  2. Use the graduated cylinder to measure about 40 mL of warm tap water and transfer it to the beaker with the sodium carbonate. Stir or swirl the beaker contents until the sodium carbonate dissolves.
  3. Pour the filtrate (copper sulfate solution) from the flask into the beaker of sodium carbonate solution, with stirring. The two solutions react immediately to form a precipitate of insoluble bluish-green copper(II) carbonate. Rinse the flask.
  4. Put a fresh piece of fan-folded filter paper into the filter funnel and place the funnel atop the empty 250 mL flask. Swirl the contents of the beaker to keep the precipitate suspended in the solution and pour the solution through the filter paper.
  5. Rinse the precipitate three times with about 25 mL of water each time. This 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.
  6. Carefully remove the filter paper and set the copper carbonate aside to dry. If you have sufficient acetone on hand, 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.
  7. Transfer the copper carbonate to a labeled bottle, and retain it for use in a later laboratory session.

Figure 6-4

Figure 6-4. Precipitated copper carbonate

Table 6-3. Recrystallization of copper sulfate – observed and calculated data

Item Data
A. crude copper sulfate ___.___ g
B. filter paper ___.___ g
C. filter paper plus purified copper sulfate ___.___ g
D. purified copper sulfate (C – B) ___.___ g


The filtrate from the copper(II) carbonate contains only sodium sulfate and a small amount of copper. You can flush it down the sink with plenty of water. The used filter paper can be disposed of with household waste.

Review Questions

Q1: The solubility of copper sulfate pentahydrate is 2033 g/L at 100 °C, which means that 305 g should dissolve in 150 mL of water at 100 °C. When we originally did this experiment, why do you suppose we used only 250 g of crude copper sulfate rather than 305 g?

Q2: We obtained about an 80% yield (roughly 200 g of purified copper sulfate from 250 g of crude copper sulfate). We might have improved the percent yield somewhat by dissolving more than 250 g of crude copper sulfate initially, but at best that method would still result in about a sixth of the crude copper sulfate going to waste. What method might we use to obtain much higher percent yields?

Q3: A chemist requires copper sulfate of at least 99.9% purity for a particular procedure. She recrystallizes crude copper sulfate as we have done, but analysis shows that her purified copper sulfate is only 99.4% pure. What might she do to obtain copper sulfate of 99.9% or higher purity?

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  1. prerna says:

    it is important tu study the process of re