Which of the following species will be able to reduce #Cu^(2+)# to #Cu# ? (A). #I_2# (B). #Ni# (C). #Al^(3+)# (D). #Ag# ?

(B).

Standard electrode potential #(E^(0))# values can be used as a guide to predicting the outcome of redox reactions.

They should be listed most negative to most positive. The relevant values are listed here:

#Al^(3+)+3erightleftharpoonsAl# #E^(0)=-1.66"V"#

#Ni^(2+)+2erightleftharpoonsNi# #E^(0)=-0.25"V"#

#Cu^(2+)+2erightleftharpoonsCu# #E^(0)=+0.34"V"#

#I_2+2erightleftharpoons2I^(-)# #E^(0)=+0.54"V"#

#Ag^(+)+erightleftharpoonsAg# #E^(0)=+0.8"V"#

#F_2+2erightleftharpoons2F^(-)# #E^(0)=+2.87"V"#

The most powerful oxidisers have the greatest ability to take in electrons. They have the most +ve #E^(0)# values and are found towards the bottom left of the table.

In the same way the most powerful reducers have the greatest ability to push out electrons. They have the most -ve #E^(0)# values and are found towards the top right of the table.

A useful rule is "Bottom left will oxidise top right".

Or "Top right will reduce bottom left"

So which of the species will be able to reduce #Cu^(2+)# to #Cu# ?

• anything above and on the right.

This leaves us with #Ni# i.e (B)

n.b This refers to standard conditions. By altering these we can often drive the reaction in the direction we want - particularly if the #E^(0)# values are close together.

The reaction may also be hampered by a high activation energy.