November 2012

An unusual and simply prepared compound of copper

Lewis Hogarth (now at graduate school) Department of Chemistry, Ohio Dominican University, 1216 Sunbury Road, Columbus OH

Introduction

At one time or another, all chemistry students have seen the colourful salts of copper(II), most commonly blue and green ones. Many students have investigated the reactions of copper(II) with a variety of common reagents as an exercise in qualitative inorganic analysis. However, few have investigated the chemistry of copper’s other oxidation state, copper(I); the only salts of copper(I) usually encountered are the oxide (red copper oxide) or one of its halides. The latter are insoluble and white in colour — not terribly interesting. The copper(I) compound described here is different!

In this short article instructions are given for the preparation of an unusual salt of copper(I). Some follow-up analyses are suggested, which I give to my freshman class as a laboratory exercise. There appear to be few references in the literature to this compound,1 and so I give the experiment to students as a simple research project. They are required to suggest the types of chemical reactions occurring throughout their investigation and to write chemical equations. Students develop an appreciation of reaction chemistry through this activity; we spend some time in lecture going over the results.

Experimental

Prepare a waterbath by filling a 250 mL beaker with about 150 mL of water and heating it to about 40°C. While the temperature is stabilizing, weigh into a 15 mm x 150 mm test tube 0.90 g of crushed sodium thiosulfate pentahydrate crystals, and add ~3-4 mL of distilled water. Into a similarly sized test tube weigh 0.50 g of crushed cupric sulfate pentahydrate crystals and add 3-4 mL of distilled water. Immerse the test tubes in the waterbath and allow them to equilibrate to the same temperature. During this process, gently agitate the samples to facilitate the dissolution process. When both salts have dissolved, add the sodium thiosulfate solution to the aqueous copper sulfate. Shake the mixture and re-immerse the tube in the ~40°C water bath for about 15 minutes. During this time several qualitative observations should be made and recorded, and a small suction filtration apparatus prepared (a Hirsch funnel or something similar). When it appears that precipitation is complete, suction filter the solid. Wash the precipitate with a few rounds of ~5 mL distilled water, and finish with a few rounds of 3 to 5 mL acetone. Continue the suction drying of the crystals for a few minutes and then transfer the precipitate to a tared screw-capped bottle and reweigh. Record the weights and the yield in grams.

Notes

It is very important that the product be dry and water-free. If not, colour changes — fluorescent canary yellow to a green colour — will be readily observed even as the material is being bottled. Suction and washing the complex with acetone serves to remove the water. My conclusion is that the compound will oxidise in moist conditions although I cannot verify this.
Take note of any qualitative observations made during the preparation of this copper complex. Note the yellow colour of the material in the photograph on page 9 of this issue.
Perform the same tests on the reactants as on the product for comparison.

Testing the compound

Suspend a small sample (~50 mg) of the compound in distilled water. Add dropwise a solution of iodine in alcohol (~0.005 M). Continue adding the iodine until no further change is noted and there is a slight excess of iodine, and then place the test tube in a rack. Perform the other tests and return to this one later.
In a dry test tube gently heat ~0.1 g of the compound with a Bunsen burner. Note any colour changes and if possible, test for the evolution of gases or for water vapour. The black material produced can be seen in the photograph on page 9.
To ~50 mg of the compound suspended in 1-2 mL of distilled water add 2-3 mL of dilute nitric acid (~2 M). Observe the suspension at room temperature and when it is gently heated.

As mentioned above, the sodium thiosulfate and copper sulfate solutions should be tested in the same fashion as the compound, concurrently with the tests on the compound, so that accurate comparisons may be made.

Questions to be considered

Identify the types of reaction taking place, i.e., oxidation/reduction, precipitation, etc.
Identify the reaction products by their properties, mostly colour.
Compare the reaction results for the compound with those of the two starting materials.

Results

1. Proposed reactions

2Cu²⁺ + 2S₂O₃^2- → S₄O₆^2- + 2Cu⁺ (a)

2Cu⁺ + S₂O₃^2- → Cu₂S₂O₃ (b)

Cu₂S₂O₃ + Na₂S₂O₃ → Cu₂S₂O₃•Na₂S₂O₃ (c)

During the addition of the thiosulfate solution to the copper solution the perhaps unexpected reduction of copper(II) to copper(I) occurs. Copper(I) is unstable in water, tending to disproportionate into copper(0) and copper(II). But in this case, because of the excess thiosulfate present, the copper(I) forms an insoluble salt, Cu₂S₂O₃, which, in turn, couples with the excess sodium thiosulfate to produce the fairly insoluble yellow compound isolated on the filter pad.

2. Reaction with iodine solution

The testing with iodine initially results in the brown colour of I₂ disappearing. Two further observations are that the excess iodine eventually disappears and a white precipitate is produced.

The iodine reacts in the usual way with the thiosulfate ion to produce colourless iodide ions and tetrathionate ions as shown:

I₂ + 2S₂O₃^2- → S₄O₆^2- + 2I^- (d)

Presumably, the iodine reacts first with the thiosulfate ion contained in sodium thiosulfate, followed by the reaction with thiosulfate derived from copper(I) thiosulfate, owing to the easier ionisation of sodium thiosulfate. This may produce unattached copper(I) ions which will react with the now large supply of iodide ions to produce copper(I) iodide and appear as a hazy white precipitate.

Cu⁺ + I^- → CuI (e)

3. Heating the material in a dry test tube

On heating the dry material, a change to an olive green colour occurs, and on prolonged heating a further change to dark brown or even black results. Also, a gas with a sharp odour will be present (acidic: pH paper; reducing: potassium dichromate paper).

During this reaction stage, the thiosulfate unit likely acts classically and produces sulfur dioxide gas which is acidic to pH paper and reducing to potassium dichromate paper.

Cr₂O₇^2- + 3SO₂ + 2H⁺ → 2Cr³⁺ + 3SO₄^2- + H₂O (f)

The thiosulfate unit also produces elemental sulfur when it is heated, and it is likely that this combines with the copper(I) to initially produce copper(I) sulfide, which on continued heating is converted to copper(II) sulfide. (The presence of sulfur was confirmed by the authors as it could be extracted from the heat-blackened product using carbon disulfide and subsequently crystallised out. I do not have a reference for this testing technique but remember it clearly from early in my career due to the unpleasant odour of carbon disulfide.)

4. Treating the precipitate with dilute nitric acid

When treated with cold, dilute nitric acid the suspension appears to get cloudier. This is likely due to the thiosulfate unit converting to sulfate and producing elemental sulfur. As the thiosulfate units are converted this releases copper(I) ions, and these are oxidised to copper(II) ions, which impart the blue colour to the solution when it is heated and releases nitrogen dioxide.

Proposed reaction:

2H+ + 3S₂O₃^2- → 2SO₄^2- + 4S + H₂O (g)

Cu+ + NO₃^- + 2H⁺ → Cu²⁺ + NO₂ + H₂O (h)

Conclusions

These qualitative experiments give only part of the story. Several students proceeded further and chemically identified some of the reaction products. An ingenious student wanted to test for the oxide and sulfide of copper after various stages of heating the solid, as well as testing the nitric acid precipitate to determine if it was elemental sulfur!

I have found that this type of experiment draws the students out and provides an excellent topic for discussion. It affords them the opportunity to devise hypotheses based on their own observations.

Reference

W.G. Palmer, Experimental Inorganic Chemistry, Cambridge University Press, 1954, pages 375-377.
(I found the compound in this reference first but a fairly extensive search of Chemical Abstracts produced little or nothing about this compound. Searching reference material before ca. 1900 produced little more than information about the preparation and analysis of this compound).

[Editor’s notes:

We came across a similar experiment online at CR Scientific LLC. The lab describes the reduction of copper(II) with thiosulfate. It might be worth reading if you are considering doing this lab.
During the proofreading process, several alternative reactions were proposed for reaction (g). For example, one proofreader suggested a simple acid-catalyzed decomposition of thiosulfate to sulfite and sulfur.

S₂O₃^2- → SO₃^2- + S

Small vial containing black substance beside small flask containing yellow substance. This illustrates the strength of this type of experiment. It allows students to have the opportunity to propose and discuss chemical reactions based on qualitative testing and past chemical knowledge. This leads to the question, “what could we do next to determine which reaction(s) might be happening?” This captures the scientific method much better than following a lab recipe.]