Question #5d695

This is an example of a redox reaction that is not spontaneous. The silver metal is a reducing agent, but a very weak one. This means it can cause to reduction of other substances (more about them later), by undergoing oxidation itself. However, being a weak reducing agent means it is difficult to oxidize.

The oxidation of silver metal is written

#Ag rarrAg^+ + e^-#

The electrode potential that goes with this process is #+0.80 V# (the V stands for volts of electric potential. You look this up it a table, like the one here: http://www.chemeddl.org/services/moodle/media/QBank/GenChem/Tables/EStandardTable.htm).

For this to happen, however, we require an oxidizing agent - one that is quite strong, because Ag is difficult to oxidize. The only oxidizing agents in this question are lead ion, #Pb^(2+)# and nitrate ion, #NO_3^-#. Note that an oxidizing agent is a substance that can be reduced.

Neither of these is sufficiently strong to cause the oxidation of the silver metal. Here is the proof:

The reduction of lead ion is given by:

#Pb^(2+)+2e^-# # rarr Pb#

But, the potential for this is #-0.13 V#

In order for a spontaneous reaction to occur, the value obtained by subtracting the potential for the oxidation #(+0.80 V)# from that of the reduction #(-0.13 V)# must be a positive number.

In this case, it is

#-0.13V - (+0.80V) = -0.93V#

a non-spontaneous reaction.

In a similar way, it can be shown that the nitrate ion is also not able to cause silver to oxidize (unless you provide a large concentration of #H^+# ions, in which case you are really working with nitric acid - a very powerful oxidizer).