Indicators of Cu(I) or Cu(II)

We performed an experiment in class where we placed a piece of copper wire in a solution of silver nitrate, we were tasked with predicted the mass of copper that should react and the mass of silver that should be formed. My question is, **what are indicators of Cu(I) and Cu(II) ions?**


$$\ce{2 AgNO3(aq) + Cu(s) -> Cu(NO3)2(aq) + 2 Ag(s)}$$


I need to know this in order to calculate reaction yield and efficiency (whether the product is $\ce{CuNO3}$ or $\ce{Cu(NO3)2}.$ 




There is very well-written chapter "*The chemistry of monovalent copper in aqueous solutions*" in *Advances in inorganic chemistry*, Volume 64 [1, pp. 220–223, DOI: [10.1016/B978-0-12-396462-5.00007-6](https://doi.org/10.1016/B978-0-12-396462-5.00007-6)] which extensively covers as to why copper(I) is unlikely to exist in aqueous solution and why nitrate is a poor ligand for the purpose of preserving monovalent copper:



> 
> The redox potentials of copper in acidic aqueous solutions are as follows:
> 
> 
> $$
> \begin{align}
> \ce{Cu+(aq) + e- &-> Cu^0(s)}, &\quad E^\circ &= \pu{0.521 V} \tag{1}\\
> \ce{Cu^2+(aq) + e- &-> Cu^+(s)}, &\quad E^\circ &= \pu{0.153 V} \tag{2}\\
> \ce{Cu^2+(aq) + 2 e- &-> Cu^0(s)}, &\quad E^\circ &= \pu{0.342 V} \tag{3}\\
> \end{align}
> $$
> 
> 
> From these values, it is clear that $\ce{Cu+(aq)}$ is unstable and
>  disproportionates in aqueous solutions:
> 
> 
> $$\ce{2Cu+(aq) <=> Cu^2+ + Cu^0(s)}, \quad K\_{1.4} = \pu{(1 - 1.7) × 10^{6} M-1} \label{rxn:4}\tag{4}$$
> 
> 
> However, if $\ce{Cu+(aq)}$ is formed in a clean aqueous medium, the disproportionation reaction has often a relatively long induction period as reaction \eqref{rxn:4a} is endothermic due to the $\ce{Cu^0(s)}$ lattice energy.
> 
> 
> $$\ce{2Cu+(aq) -> Cu^2+(aq) + Cu^0(aq)}\label{rxn:4a}\tag{4a}$$
> 
> 
> Furthermore, the concentration of $\ce{Cu^0(aq)}$ in nonacidic aqueous solutions is limited due to the low solubility of $\ce{CuOH}/\ce{Cu2O}:$
> 
> 
> $$\ce{Cu+(aq) + OH- <=> CuOH(s)/Cu2O(s)} \quad K\_\mathrm{SP} = \pu{1E-14 M-2} \label{rxn:5}\tag{5}$$
> 
> 
> Naturally, the concentration of Cu(I) species in aqueous solutions
>  can be increased by adding appropriate ligands.
> 
> 
> […]
> 
> 
> …the instability of $\ce{Cu+(aq)}$ stems from the difference between the solvation energies of the two cations and the second ionization potential of copper.
>  This is correct for all transition-metal ions, but as the ionization energies increase along the series faster than the solvation energies, the lower oxidation states become relatively more stable.
>  In order to shift reaction \eqref{rxn:4} to the left ligands which
>  preferentially bind to $\ce{Cu(I)},$ must be added to the solution.
> 
> 
> […]
> 
> 
> $\ce{Cu(I)L\_n}$ complexes can be stabilized in aqueous solutions via
>  employing ligands with one or more of the following properties:
> 
> 
> 1. Ligands that have a lower coordination number with $\ce{Cu(I)}$
>  complexes than with $\ce{Cu(II)}$ complexes; these $\ce{Cu(I)}$ complexes are stabilized due the entropy gain in reaction \eqref{rxn:7}:
> 
> 
> $$\ce{Cu(II)L\_m + e- -> Cu(I)L\_n + $(m - n)$ L} \label{rxn:7}\tag{7}$$
> 
> 
> 2. Ligands that are π acids and therefore prefer $\ce{Cu(I)}$ over $\ce{Cu(II)}$ and shift the redox potential of the $\ce{Cu(II)/Cu(I)}$ couples anodically.
> 3. Hydrophobic ligands, which often also increase the radius of
>  the complex, decrease the stabilization of $\ce{Cu(II)}$ by decreasing
>  its solvation energy.
> 4. Ligands which enforce a tetrahedral coordination sphere, or
>  one which approaches this geometry, on the central cation.
> 5. Ligands that are soft bases, for example, thiols. Thus, for example, the biological copper-trafficking proteins, the metallochaperone Atx1, and the copper-transporting P-type ATPases have a highly conserved CXXC metal binding motif which binds a single $\ce{Cu(I)}$ ion via a two-coordinate site consisting of two cysteines.
> 
> 
> All these types of ligands stabilize $\ce{Cu(I)}$ complexes by shifting the redox potentials of the $\ce{Cu(II)/Cu^0(s)}$ and $\ce{Cu(I)/Cu^0(s)}$ couples and thus shifting the equilibrium constant of reaction \eqref{rxn:4} to the
>  left.
>  However, ligands might stabilize $\ce{Cu(I)}$ complexes also
>  kinetically by inhibiting the disproportionation reaction by
>  inhibiting the formation of $\ce{Cu^0(s)}$ via blocking the approach of two copper atoms, this is, for example, the mechanism by which
>  some $\ce{Cu(I)}$ enzymes are stabilized.
> 
> 
> Naturally, all these ligands also increase the solubility of $\ce{Cu(I)}$ by competing with reaction \eqref{rxn:5}.
> 
> 
> 


### References


1. *Inorganic/Bioinorganic Reaction Mechanisms*, 1st ed.; Eldik, R. van, Ivanović-Burmazović, I., Eds.; Advances in inorganic chemistry; Academic Press: London; Waltham, MA, **2012**; Vol. 64. ISBN 978-0-12-396462-5.


 5

As noted in the comments, $\ce{Cu+}$ is not stable in aqueous solution. It tends to disproportionate to $\ce{Cu^0}$ and $\ce{Cu^{2+}}$. You can therefore assume that all but a negligible amount turned into $\ce{Cu^{2+}}$.


Source:
Holleman/Wiberg: Lehrbuch der Anorganischen Chemie, de Gruyter, 101. edition, 1995.


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