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I was balancing this reaction when I found an unusual compound had been formed. $$\ce{Cu(s) + HNO3(aq)-> Cu^{3+}(aq) + NO(g)}$$

My attempt:

Balancing oxidation number: $$\ce{Cu -> Cu^{3+} +3e^-}$$ $$\ce{3e^- + NO_3- -> NO}$$ Balancing hydrogen and oxygen atoms: $$\ce{3e^- + 4H^+ + NO3- -> NO + 2H2O}$$ Adding $\ce{3NO3-}$ on both sides to form compounds: $$\ce{3e^- + 4H+ + 4NO3- -> NO + 2H2O + 3NO3-}$$ Adding both half reactions $$\ce{Cu + 4H+ + 4NO3- -> NO + 2H2O + Cu^{3+} + 3NO3-}$$ $$\ce{Cu + 4HNO3 -> NO + 2H_2O + Cu(NO3)3}$$

So does this compound $\ce{Cu(NO3)3}$ actually exist? or have I done a mistake.

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    $\begingroup$ Could you clarify where the first equation came from. I'm guessing that you were trying to determine what the products would be for copper reacting with nitric acid and you developed the equation rather than trying to "balance it". $\endgroup$ – MaxW Nov 4 '15 at 3:30
  • $\begingroup$ I had a test yesterday and it was in the paper. $\endgroup$ – Rahul Gupta Nov 4 '15 at 3:41
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    $\begingroup$ Did some digging with Google. Evidently some Cu+3 compounds exist as an anion complex where copper has organic ligands. pubs.acs.org/doi/abs/10.1021/ja00244a020 The equation does balance. However $\ce{Cu(NO3)3}$ does not exist. $\endgroup$ – MaxW Nov 4 '15 at 3:59
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    $\begingroup$ Copper (iii) salts need a very basic anion such as fluoride to form , for this reason it mostly exists as an oxide or fluoride or something like that. I'm not sure whether the nitric acid would form the copper (iii) nitrate salt. $\endgroup$ – Technetium Nov 4 '15 at 4:06
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    $\begingroup$ Inorganic chemists have cooked up some weird chemicals. I'd be curious as to how copper (iii) nitrate could be prepared. $\endgroup$ – MaxW Nov 4 '15 at 4:51
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Copper(III) nitrate cannot be obtained from aqueous nitric acid, and likely doesn't exist.

Reaction carried under strongly oxidative conditions between $\ce{Cu(NO3)2}$ and fuming $\ce{HNO3}$ yields in nitrosyl copper(II) trinitrate $\ce{[NO+][Cu(NO3)3−]},$ [1, 2] sometimes written as adduct $\ce{Cu(NO3)2 · N2O4},$ which is contradictory to the crystal structure consisting of copper(II) centers octahedrally coordinated by nitrate groups and the voids in cross-linking mesh of octahedra are occupied by nitrosyl ions [3]:

Crystal structure of nitrosyl copper(II) trinitrate

Figure 1. Crystal structure of nitrosyl copper(II) trinitrate $\ce{[NO+][Cu(NO3)3−]}$ (ICSD-422033). Color code: $\color{#3050F8}{\Large\bullet}~\ce{N}$; $\color{#FF0D0D}{\Large\bullet}~\ce{O}$; $\color{#C88033}{\Large\bullet}~\ce{Cu}$.

According to Wiberg's Anorganische Chemie [4, p. 1707], oxidation of $\ce{Cu^2+}$ to $\ce{Cu^3+}$ is possible only with strong oxidation agents $(E^\circ(\ce{Cu^2+/CuO+}) = \pu{+1.8 V}),$ and copper(III) is highly unstable in water and has a tendency to oxidize the medium $(E^\circ(\ce{H2O/O2}) = \pu{+1.23 V}).$ The reduction potential or $\ce{Cu^3+}$ can, however, be reduced down to $\pu{+0.45 V}$ by coordinating it with appropriate strongly donating ligands (which nitrate really isn't). Examples of such complexes are:

  • $\ce{K3CuF6}$ (oxidant: fluorine $\ce{F2}$; $\ce{CuF3}$ is unknown);
  • $\ce{K5[Cu(IO6H)2]},$ $\ce{K5[Cu(TeO6H2)2]},$ $\ce{Na5[Cu(TeO6H2)2]}$ (oxidant: peroxodisulfate $\ce{S2O8^2-}$ in strongly alkaline medium);
  • $\ce{Na[Cu(OH)4]},$ $\ce{KCuO2},$ $\ce{Ba(CuO2)2}$ (oxidant: hypobromite $\ce{BrO-}$/oxygen $\ce{O2}$; $\ce{Cu2O3}$ is unknown).
  • Cu(III) complex of 2-ethoxy-5,10,15,20-tetrapentafluorophenyl-3,7-diaza-21,22-dicarbaporphyrin (Cu(III)-$\ce{N2CP}$, a doubly N-confused porphyrin) [5]
  • A variety of Cu(III) peptide complexes [6, pp. 213–223]

References

  1. Znamenkov, K. O.; Morozov, I. V.; Troyanov, S. I. Synthesis and Crystal Structure of Copper(II) Nitrato Complexes $\ce{NO[Cu(NO3)3]},$ $\ce{Na2[Cu(NO3)4]},$ and $\ce{Ag2[Cu(NO3)4]}.$ Russian Journal of Inorganic Chemistry 2004, 49 (2), 172–179. (in Russian)
  2. Gagelmann, S.; Rieß, K.; Wickleder, M. S. Metal Oxidation with $\ce{N2O5}$: The Nitrosylium Nitrates $\ce{(NO)Cu(NO3)3},$ $\ce{(NO)2[Zn­(NO3)4]}$ and $\ce{(NO)6[Ni4(NO3)12](NO3)2·(HNO3)}.$ Eur. J. Inorg. Chem. 2011, 2011 (33), 5160–5166. https://doi.org/10.1002/ejic.201100565.
  3. Blackwell, L. J.; King, T. J.; Morris, A. Crystal Structure of the Copper Nitrate–Dinitrogen Tetroxide Adduct $\ce{Cu(NO3)2 · N2O4}.$ J. Chem. Soc., Chem. Commun. 1973, No. 17, 644–644. https://doi.org/10.1039/C39730000644.
  4. Wiberg, E.; Wiberg, N.; Holleman, A. F. Anorganische Chemie. Nebengruppenelemente, Lanthanoide, Actinoide, Transactinoide, 103. Auflage.; De Gruyter: Berlin; Boston, 2017; Vol. 2. ISBN 978-3-11-049573-7. (in German)
  5. Furuta, H.; Maeda, H.; Osuka, A. Doubly N-Confused Porphyrin:  A New Complexing Agent Capable of Stabilizing Higher Oxidation States. J. Am. Chem. Soc. 2000, 122 (5), 803–807. https://doi.org/10.1021/ja992679g.
  6. Karlin, K. D.; Tyeklár, Z. Bioinorganic Chemistry of Copper; Springer Netherlands: Dordrecht, 1993.
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Copper(III) compounds indeed do exist, but are very strong oxidizers. It is safe to assume that they are not formed in water solutions unless some special ligand is present, such as periiodate.

I highly doubt that Cu(III) nitrate was ever obtained in water solutions.

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Yes, copper (III) nitrate does exist according to ChemSpider. Although I think it is quite rarely used. Copper (III) salts require a very basic anion such as oxide or fluoride to be produced and are usually found in these forms.

Reacting nitric acid with solid copper, under normal conditions, would most likely form the copper (II) nitrate salt as the nitrate anion may not be basic enough for copper (III) ions to form. Fuming nitric acid, possibly under specific conditions, or another nitrating agent must produce the copper (III) nitrate.

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