This post was made after reading Boffis' post in an old thread dating back to 2011; after trying his proposed method and obtaining better results than his, it’s with his kind
permission that I am now posting the improved method here. I believe his idea was excellent, but was hindered by the too quick addition of the
solutions, and the lack of mechanical stirring surely didn't help his cause. Additionally, the pH could be better kept in check with a dual system of
indicators, and the need to expose the salt to air during vacuum filtration can be much reduced by washing the product with Ethanol, and then in a
homemade vacuum system. Lastly, this procedure seems to yield better results on a much smaller scale. His idea was nonetheless amazingly simple, yet
effective.
Introduction
Copper(I) cyanide is a white to butter-cream-off-white solid that is almost insoluble in water (2.5 mg/L) and virtually insoluble in alcohol. It is a
much safer compound than alkali cyanides (which are very soluble and emit vapours of HCN on standing, and in neutral/acid solution), but it also has
more limited applications, mostly in the preparation of organic nitriles. Organic chemistry is not my field, so I’ll not dwelve into that.
A crude procedure for the preparation of CuCN involves the bare reaction between solutions of CuSO4 and NaCN: the reduction of Cu(II) to
Cu(I) is carried out by a portion of the cyanide, which is oxidated to the toxic gas cyanogen, (CN)2. This reaction should never be
attempted, as there are much safer versions that employ a reducing agent.
2 CuSO4 + 4 NaCN → 2 CuCN + (CN)2 + 2 Na2SO4
According to an earlier edition of the Vogel (1974), CuCN can be prepared with CuSO4, NaHSO3 and NaCN. In this reaction, the
bisulfite acts as the reducing agent towards Cu(II) and is oxidised to bisulfate. This reaction is not exactly OTC, as NaCN is an expensive, often
times monitored, toxic compound, and NaHSO3 is generally hard to come by at an affordable price.
This procedure was slightly improved in a latter edition (1988) by employing Na2S2O5, a cheap food additive. The next
improvement was made by Boffis (2011), by using Na2S2O5 and K4[Fe(CN)6]/NaOH. With this set of
reactants the price we pay to avoid the formation of HCN gas during the handling of alkali cyanides is the slow but bearable emission of
SO2 gas. The reactions below are heavily simplified, as they are summing up a complex system of equilibria.
Na2S2O5 + H2O ⇌ 2 “NaHSO3”
HSO3- + H+ ⇌ SO2 + H2O
K4[Fe(CN)6] + NaOH ⇌ K4[Fe(CN)5OH] + “NaCN”
[it is a much more complicated reaction, a ligand substitution with dissociative activation on the complex hexacyanoferrate(II) coordination centre]
Fe(II) stays in behind in the acidic solution as FeSO4.
Techniques
The pH of the reaction should be at about 5.5-6.0, and should be monitored with two sets of indicator paper: 3.8-5.4 (Bromophenol blue) and 5.5-9.0
(Bromothymol blue), available for a few cents on Aliexpress, or easily made at home. At optimal conditions, the upper set is yellow-green, and the
lower set is blue. If the pH increases too much, it is shown by the upper set of indicating paper; if the pH is dangerously acidic, it is revealed by
the lower set. Towards the end of the reaction, the pH will be < 3.8, due to the huge concentration in solvated SO2: the reaction is
complete, and as long as the amounts have been properly calculated and added, it is safe to digest the precipitated CuCN on boiling, as no
cyanide-based fumes will be released.
At no point, and for no reason, should the pH be hastily manipulated with strong acids.
If the solution is too alkaline, iron flocculates will ruin the product, and will take some time to redissolve again. If the solution suddely reaches
pH 9 or 10, and is acidified with an improvised amount of concentrated acid, there is the risk of making it dangerously acidic, with a huge emission
of SO2, and, possibly, HCN or (CN)2.
Controlling the pH also means controlling the safety and the redox chemistry of this reaction, and so the presence of iron.
Iron presence, assuming that there are no flocculates, can be lowered drastically by working with a cotton syringe, a simple and inexpensive technique
to filtrate small amounts of solutions without wasting chemicals and filter paper. This is excellent in a home lab, as many of our chemicals end up
being contaminated by dust, paper fibers, or other solid impurities, especialy if they come from craft shops. To avoid the sudden formation of
flocculates, all of the solutions in this experiment are added through a cotton filter, that both removes solid traces and ensures a dropwise
addition.
The cotton syringe can be set up with a plastic or glass syringe, without needle, and a small ball of cotton tightly packed on the bottom with the
piston. By holding the syringe on a clamp above the solution, you can now mix the two solutions in a dropwise manner, which is doubly convenient in
this preparation, as you get to both limit the flocculation of iron, and to keep the pH under constant control.
Lastly, as soon as the white product precipitates, the system is kept under an aluminium cover to reduce the exposure to light, and the subsequent
operations, especially when the product is free of reducing sulfites, are carried out in dim light. This may be a cyanide, but it’s still a Cu(I)
salt, and is to be treated as a light-sensitive compound, like CuCl, CuBr and CuI.
Three solutions are prepared.
a) in a 100 mL beaker, with a stir bar, CuSO4·5w (10 g) in cold water (80 mL).
b) in a 25 mL beaker, with a stir bar, NaOH (1.5 g) in cold water (9 mL).
c) in a 250 mL beaker, with a stir bar, K4FeCy6·3w (2.84 g) and K2S2O5 (2.31 g) in cold water
(40 mL). pH is around 5.4
Half of b) is added, dropwise via a cotton syringe, to c), under heavy stirring. The final pH should be about 7.5. Now, the copper solution a) is
added to the neutral reducing ferrocyanide under heavy stirring. The addition through a 20 mL cotton syringe allows us to add it dropwise in 5
subsequent runs, and lets us monitor the pH in real time with the two sets of indicating paper. The first addition forms a brick-red suspension, which
turns burgundy-brown. By the second and third additions, the suspension goes through a whole café of shades: coffee-brown, caffelatte, cappuccino,
and pinkish-macchiato-ochra. This is, incidentally, the same sequence of colours observed during the preparation of CuI.
During this time, the remaining NaOH is added with a pipette to keep the pH in check. By the second half of the additions, the NaOH is finished, and
the suspension is at a pH of 3.5-3.8. I decided to increase the pH to 5.4 with more NaOH, but this can prove risky, especially if done hastily, and
the constant emission of SO2 will keep the pH to about 3.8 regardless of your actions.
The beaker is covered with a watchglass and the hotplate is turned on with stirring. After a couple of minutes of stirring, the yellowish ochra
suspension turns to a dirty off-white and finally a milky white. The final pH is, again, 3.5-3.8, and the suspension is digested by boiling.
Once a satisfactorily white solution is observed (5’ of boiling), the beaker is taken off the hotplate and cooled with tap water. From this moment,
the beaker is kept under an aluminium foil cover.
After a couple of minutes enough curdy white solid has sedimented, and the suspension is decanted (150 mL). To the beaker is added some cold water (50
mL + 0.2 g K2S2O5) and the suspension is boiled again for 10’. The suspension is cooled again, decanted (60 mL).
The product is filtered on a Buchner funnel, thoroughly washed with water (at least 2x10 mL) and 96% ethanol (at least 3x5 mL).
The solid is left on the pump for 5’, covered under aluminium foil. The white powder is placed in an amber glass vial and dried under vacuum, away
from light, for at least an hour, powdered with a thin spatula and weighed constantly. When constant weight is reached, it is stored in a foil-wrapped
brown vial. Alternatively, the solid can be dried in the oven at 100-110°C. The last washings contain little to no iron to the SCN test, and only
give a green cloudiness on addition of NaOH. The product does not smell of SO2, and turns only very slightly pink after months of storage
away from light.
I started from 9.9779 g / 249.69 g/mol of CuSO4·5w = 39.961 mmol Cu(II).
I also used 2.8425 g / 422.388 g/mol of K4FeCy6·3w = 6.7275 mmol Fe(II) = 40.365 mmol CN-.
The limiting reactant is thus CuSO4·5w.
Theoretical yield: 39.961 mmol * 0.089563 g/mmol = 3.5790 g of CuCN.
Actual yield: 3.4819 g
% yield: 97.2%
In a follow-up post I'll show some qualitative spot tests.
[Edited on 7-10-2020 by valeg96]
[Edited on 7-10-2020 by valeg96]valeg96 - 7-10-2020 at 05:44
As you may have noticed, the original preparation was made back in april 2020. Now it's october, and I finally found the will to test it a it more
extensively.
Within 6 months, the solid has turned a pinkish shade, I'm not sure if because of very slight Fe(II) traces that oxdised to Fe(III) or because of some
degradation of the copper cyanide. Anyways, I've tested the compound with some classical reagents. Briefly:
NaOH: slowly forms a gelatinous precipitate of Cu(OH)2. No obvious presence of reddish-brown Fe(OH)3 or greenish
Fe(OH)2, but the solution retains a grayish shade.
5% ammonia: solution turns into a very slight blue shade, and the solid appears grayish. No obvious blue shade of
[Cu(NH3)4]2+ ions.
Ferricyanide: Solution is deep brick red with a precipitate. This could be cupric ferrocyanide Cu2[Fe(CN)6], the Cu(II) and the
Fe(II) being produced with a redox reaction.
Ferrocyanide: cream-pink solid, probably some Cu(I) ferrocyanide salt, the Cu(II) analogue Cu3[Fe(CN)6]2 being
greenish.
Ferrocyanide/HSO3-: white gelatinous precipitate. As above.
HCl 32%: yellowish solution of [CuCl4]2- anions.
Silver nitrate 2%/HNO3: White precipitate of AgCN.
KI: no result.
Na2S 2%: black precipitate, most likely Cu2S, but identical anyways to FeS and CuS.
KSCN: Clear solution, with some pinkish specks. Nothing close to the blood red soluble Fe(SCN)3, white CuSCN or black Cu(SCN)2.
2,2'-bipy in EtOH: grimy ochra precipitate, which doesn't match with the red soluble complex of Fe(II) [Fe(bipy)3]2+ or the
yellow soluble analogue of Fe(III). This could be a reddish coordination compound of Cu(I), bipy and CN like those found by Chesnut 1999.
dmg in 5% ammonia: brownish grimy precipitate that doesn't match with the bright red soluble Fe(II) Fe(dmg)2 complex.
I believe none of these tests are strikingly positive for iron. I'm sure some modern instrumental analysis could detect it, tough.
These tests are complicated by the fact that there are three analytically relevant species around: Cu(I), traces of Cu(II) and CN- ions, if we assume
Fe isn't around.
[Edited on 7-10-2020 by valeg96]
[Edited on 7-10-2020 by valeg96]
[Edited on 7-10-2020 by valeg96]
[Edited on 7-10-2020 by valeg96]Boffis - 15-10-2020 at 05:10
Really nice work and write-up Valeg96. I am glad to see you managed to get a nice white product something I never quite managed to do. Like you I
tested it for iron and couldn't find conclusive evidence that there was much more than a trace present. I think the faint pinkish colour in my cuprous
cyanide is due to co-precipitated cuprous oxide or possibly metallic copper. Cu+ tends to disproportionate to Cu++ and Cu under even mildly acid
conditions but the CN- ions stabilize the monovalent state by effectively removing it from the reaction solution.valeg96 - 19-10-2020 at 07:00
I think the faint pinkish colour in my cuprous cyanide is due to co-precipitated cuprous oxide or possibly metallic copper.
Sorry for the late reply, but I suspect that's the case as well, leaning more towards the Cu2O for me. I'm glad you've appreciated the preparation; it
was very hectic and I couldn't waste time to take good photos, but I think they are decent enough.
Now to find something to do with it... I'd like to find some coordination chemistry that uses CuCN without added alkali cyanides.
