can you store copper sulphate solution in a zinc pot?

Explanation

`a.` From Table `3.1` , the standard reduction potential of `Zn^(2+)|Zn` is `-0.76V` and standard oxidation potential will be `0.76 V`, whereas the standard reduction potential of `Cu^(2+)|Cu` is `0.34V` and standard oxidation potential will be `-0.34 V`.
Therefore ,` E^(c-)._(o x i dation(Zn|Zn^(2+))gtE^(c-)._(o x i d ation(Cu|Cu^(2+)).`
Hence, `Zn` is more reactive than `Cu.` Hence, it displaces `Cu` from `CuSO_(4)` solution as follows `:`
`Zn(s)+CuSO_(4)(aq)rarrZnSO_(4)(aq)+Cu(s)`
or
`Zn(s)+Cu^(2+)(aq)rarr Zn^(2+)(aq)+Cu(s)`
Thus, `Zn` reacts with `CuSO_(4)` solution. Hence,` CuSO_(4)` solution cannot be stored in `Zn` pot.
Alternatively
Find `E^(c-)._(cell)` . If it is positive, it means cell reaction will occur, and one cannot store the solution in the pot. If the standard oxidation potential of the metal behaving as pot is greater than the standard oxidation otential of the metal consisting of solution ,then `E^(c-)._(cell)` will be positive and one cannot store the solution in the pot.
`:. E^(c-)._(cell)=(E^(c-)._(reduction))_(cathode)-(E^(c-)._(reduction))_(anode)`
`=E^(c-).Cu^(2+)|Cu)-E^(c-)._(Zn^(2+)|Zn)`
`=0.34-(-0.76)=1.10V`
Therefore, a solution of `CuSO_(4)` cannot be stored in `Zn` pot.
`b.` From Table `3.1`, the standard reduction potential of `Ag^(o+)|Ag` is `0.80V` and standard reduction potential of `Cu^(2+)|Cu` is `0.34 V` .
`E^(c-)._(Ag^(o+)|Ag)gtE^(c-)._(Cu^(2+)|Cu)`
or
`E^(c-)._(Ag|Ag^(o+))ltE^(c-)._(Cu|Cu^(2+))`
Since the standard oxidatioin potential of the metal `(Ag)` behaving as pot is less than the standard oxidation potential of the metal consisting of solution `(CuSO_(4))`, the `E^(c-)._(cell)` will be negative, and one can store the solution `(CuSO_(4))` in the pot of `Ag`.
`E^(c-)._(cell)=(E^(c-)._(reduction))_(cathode)-(E^(c-)._(reduction))_(anode)`
`=E^(c-)._(Cu^(2+)|Cu)-E^(c-)._(Ag^(o+)|Ag)`
`=0.34-0.80=-0.46V`
Therefore, one can store `CuSO_(4)` in `Ag` pot.