CAUTION: If you choose to attempt any of the experiments or procedures described on this site, you do so entirely at your own risk.
Reduction of Copper (II) with Thiosulfate
The "Yellow Snowstorm" Reaction
Sodium thiosulfate [Na2S2O3·5H2O] dissolves in water to give thiosulfate ion (S2O32-), a useful reducing agent. Copper (II) sulfate [CuSO4·5H2O] dissolves to give cupric ion (Cu2+); in the context of a redox reaction with thiosulfate, the cupric ion acts as an oxidizing agent. The vivid color of Cu2+ makes it easy to see when it's been reduced by the thiosulfate; the blue cupric ion becomes the colorless cuprous ion.
We can see from the reduction potentials that thiosulfate should reduce Cu2+ to Cu+, possibly all the way to copper metal. Does this happen in real life? We might expect a two-reagent redox system to give simple, predictable results. Here we do two experiments with copper (II) sulfate and sodium thiosulfate to see if the results are what we expect. The first one involves an equimolar ratio of copper sulfate to sodium thiosulfate; the second uses an excess of sodium thiosulfate.
Wear safety goggles at all times. Your laboratory must be inaccessible to children, animals, and untrained personnel. The reagents must not be ingested.
Methods & Observations:
Using clean volumetric flasks, prepare the following solutions:
A = 0.5 M copper (II) sulfate
B = 0.5 M sodium thiosulfate
You should prepare at least 10 mL of each. Store in tightly-sealed containers to prevent evaporation. We will conduct two experiments using these solutions.
Experiment 1. Equimolar Reaction
In a clean test tube, place 2 mL of A. In another clean test tube, place 2 mL of B; a 13x100 test tube is about the right size.
Pour the contents of one test tube into the other. Stir with a glass rod. Cover the test tube with polyolefin lab film or a loose-fitting stopper and let stand.
Notice the almost instantaneous color change. Upon mixing, the solution should assume a pale green color. Compare this with the color of cupric ion in solution. What happened?
After 15 minutes or so, a precipitate should begin to form. Is it homogeneous, or are there two or more types of precipitate evident?
The reaction appears complete by the next day. Note the yellow color. Did this reaction go as we expected from the redox potentials? (Note: Though the yellow bears a strong superficial resemblance to sulfur, 0 oxidation state, a reader has done some tests and searched the literature; it appears the yellow material is actually a Cu-thiosulfate complex).
Centrifuge the mixture to settle the precipitate. Pour off the liquid into a clean test tube. Save both for further tests.
Place some of the liquid in a spot plate well where the air can get to it. Let it sit for 24 hours. What happens? How about after 48 hours? With prolonged standing you may notice that a thin, dark layer settles out on top of the yellow layer.
Wash the precipitate several times with distilled water. Does any of the yellow material dissolve?
Let some of the washed precipitate sit in a spot plate well for 24 hours. Does it change visibly? What happens when you place a small amount of the dried powder in a capillary tube and heat it in the burner flame (cautiously)?
Experiment 2. Excess of Thiosulfate
This time we'll repeat the general procedure using 1 mL of cupric sulfate solution and 3 mL of sodium thiosulfate solution.
Do you notice the pale green color this time? After several hours, is there much of a sulfur precipitate?
Why did the reactions behave differently? In the case of equimolar amounts of Cu2+ and S2O32-, sulfur (or something resembling it) was produced in abundance. In the case of a large excess of S2O32- relative to the Cu2+, however, we get what appears to be either copper metal or maybe cuprous oxide, Cu2O. Could it instead be a copper sulfide such as Cu2S or CuS? It's hard to tell by looking at it; we might consider centrifuging, drying, and testing this material to see whether it's an oxide or a sulfide. (Update: someone else has done the same experiment, and they verified the dark precipitate as CuS).
The formation of the brown precipitate appears complete after two days. There also appears an iridescent coating on the inside of the test tube.
As the days run into weeks, the iridescent coating becomes thicker and darker. It takes on a blue-black appearance, typical of covellite (CuS) or other Cu sulfide mineral. To verify that it's a sulfide, it could be dried and then treated with a drop of nitric acid. A piece of lead acetate paper held near the sample will blacken in the presence of H2S (Toxic! Do the acid test outside or in a fume cupboard!)
Conclusions and Discussion:
The chemistry of the Cu2+ / S2O32- reaction is a little more complex than we might have guessed. No one seems to know exactly what's happening in the way of half-reactions; the picture is complicated greatly by complexes and side reactions. Changing the relative proportions of reactants seems to have a remarkable effect on the outcome.
In Experiment 1, at least some of the thiosulfate reduces at least some of the copper (II) to copper (I). An excess of thiosulfate, as in Experiment 2, appears to perform the reduction completely. The reduction of cupric to cuprous ion is fast. The formation of the yellow material (Experiment 1) or copper sulfide (Experiment 2) is slow.
Update: Jordan Vonnida, writing in from Queensland University of Technology, says that the yellow material in Experiment 1 is probably a copper (I)-thiosulfate complex, possibly with polymeric character. He says a literature search suggests the precipitate may be either [Cu(S2O3)3]5- or a polymeric complex having a tetrahedral Cu+ center.
Furthermore, he tested the yellow precipitate and found it not to give the characteristic melting habit of sulfur nor the blue flame when burned. The material also decomposed in HCl (or when heated) to give a gas presumed to be SO2. The gas reduced red-orange Cr(VI), as dichromate, to green Cr(III).
The reaction of copper (II) with excess thiosulfate, as in Experiment 2, shows almost complete loss of color (with perhaps a pale yellowish tinge remaining) for the first several hours. There is very little yellow precipitate formation. This suggests that, in Experiment 1, some unreduced Cu2+ (or perhaps a Cu2+ / thio complex) is required for whatever happens next. Since we didn't see the green color fading as more yellow precipitate settled out, we might guess the action to be catalytic.
The green color of the solution formed in Experiment 1 remains during the formation of the "sulfur" (?) precipitate, but it seems to go back to pale blue eventually. This is most visible after centrifuging down the precipitate. It could be that a complex ion is present but dissociates over time.
We could perform additional experiments to see exactly how much cupric ion remains in solution after all the yellow solid has precipitated.
Does sulfate (SO42-) remain unchanged in the reaction, or does it participate? A simple, qualitative test would be to add aqueous barium nitrate or barium chloride to the blue liquid to see if a precipitate forms. Any sulfate ions in solution would immediately combine with barium ions to make the highly insoluble BaSO4.
It might be instructive to repeat the experiment with cupric chloride instead of cupric sulfate. One could perform the experiments alongside one another to get an estimate of how fast the solid copper-thiosulfate complex is formed in either case. If none were formed in the cupric chloride experiment, we might have to modify our conclusions.
In Experiment 2, we can dry and weigh the CuS precipitate and compare it with the number of moles CuSO4 introduced.
The best possible outcome of this experiment has been achieved: at least one reader became seriously interested in finding out what was happening. He tried the experiment himself, tested some assumptions, and cleared up a mystery.
CR Scientific Catalog