Indicators of Cu(I) or Cu(II)
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] 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:
  
  $$
\begin{align}
\ce{Cu+(aq) + e- &amp;-&gt; Cu^0(s)}, &amp;\quad E^\circ &amp;= \pu{0.521 V} \tag{1}\\
\ce{Cu^2+(aq) + e- &amp;-&gt; Cu^+(s)}, &amp;\quad E^\circ &amp;= \pu{0.153 V} \tag{2}\\
\ce{Cu^2+(aq) + 2 e- &amp;-&gt; Cu^0(s)}, &amp;\quad E^\circ &amp;= \pu{0.342 V} \tag{3}\\
\end{align}
$$
  
  From these values, it is clear that $\ce{Cu+(aq)}$ is unstable and
  disproportionates in aqueous solutions:
  
  $$\ce{2Cu+(aq) &lt;=&gt; Cu^2+ + Cu^0(s)}, \quad K_{1.4} = \pu{(1 - 1.7) × 10^{6} M-1} \label{rxn:4}\tag{4}$$
  
  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.
  
  $$\ce{2Cu+(aq) -&gt; Cu^2+(aq) + Cu^0(aq)}\label{rxn:4a}\tag{4a}$$
  
  Furthermore, the concentration of $\ce{Cu^0(aq)}$ in nonacidic aqueous solutions is limited due to the low solubility of $\ce{CuOH}/\ce{Cu2O}:$
  
  $$\ce{Cu+(aq) + OH- &lt;=&gt; CuOH(s)/Cu2O(s)} \quad K_\mathrm{SP} = \pu{1E-14 M-2} \label{rxn:5}\tag{5}$$
  
  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:
  
  
  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}:
  
  
  $$\ce{Cu(II)L_m + e- -&gt; Cu(I)L_n + $(m - n)$ L} \label{rxn:7}\tag{7}$$
  
  
  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.
  Hydrophobic ligands, which often also increase the radius of
  the complex, decrease the stabilization of $\ce{Cu(II)}$ by decreasing
  its solvation energy.
  Ligands which enforce a tetrahedral coordination sphere, or
  one which approaches this geometry, on the central cation.
  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


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.