Why can't nitrate ion in aluminium nitrate solution oxidise copper when a copper spoon is used to stir the solution? I have been recently studying electrochemistry and I came across the electrochemical series. A question was asked saying what will happen if a copper spoon is used to stir an aluminium nitrate solution,the answer was no reaction will take place as Aluminium is more stable in its ionic form and copper cannot reduce aluminium ion, but can't nitrate ion which is above copper in the spectrochemical series oxidize copper and be transformed into NO which is more stable than Cu. Why doesn't it happen?
Minor correction: It is not the spectrochemical series but rather the electrode potential tables. Those tabulated values are given under standard conditions
You are indeed correct in thinking that nitrate ion should oxidize copper. It is a game of pH. If you have very low pH and nitrate ions, the following reaction takes place (in the absence of other ions)...
$$\ce{Cu(s) + 4 HNO3(aq) -&gt; Cu(NO3)2(aq) + 2 NO2(g) + 2 H2O(l)}$$