Write a net ionic equation for the following: Acidified solutions of sodium dichromate and copper (I) bromide are mixed?

I know that the #Cr_2O_7^{2-}# will reduce to #Cr^{3+}#, but what isn't immediately obvious to me is why #Cu^+# is oxidized into #Cu^{2+}# how do we know necessarily that the copper oxidizes, and that it specifically oxidizes to copper(II).

There was also another problem where #Fe^{2+}# was oxidized to #Fe^{3+}#. In general, how do we know whether or not the metals oxidize, and what specifically they oxidize to.

1 Answer
Jan 8, 2018

Well, usually these redox equations follow standard protocols....

Explanation:

#Fe(II+)# is oxidized to #Fe(III+)#.....#Cu(+I)# is oxidized to #Cu(+II)#...at A level, it is usually made clear as to the individual redox process that occur. At undergraduate level you would have to be mindful of standard redox processes.

Here, we know that red-orange dichromate ion, a potent oxidant, is reduced to green #Cr^(3+)#....(there was an old song about #"the orange and the green"#, it referred to sectarian division in Northern Ireland)...

#Cr_2O_7^(2-) +14H^+ + 6e^(-)rarr 2Cr^(3+) + 7H_2O(l)# #(i)#

#underbrace(Cr(VI+))_"red-orange"rarr#

#underbrace(Cr(III+))_"green"#

And for every reduction there is a corresponding oxidation...

#Cu^+ rarrCu^(2+) + e^(-)# #(ii)#

And so for the final redox process we take... #(i)+6xx(ii)#

#Cr_2O_7^(2-) +14H^+ + 6Cu^(+)rarr 2Cr^(3+) + 6Cu^(2+) +7H_2O(l)#

....the which I think is balanced with respect to mass and charge...