Fe2+ is oxidized to Fe3+ and CrO4is reduced to Cr3+ using an unknown sample containing Fe2+ and 0.04322M K2CrO4 solution in a redox titration. I first need to balance the formula, but I don't know what to do with the potassium.

Answer 1

The simple thing to do is to ignore it. It's a spectator ion. The rxn is between the chromate ion and the ferrous (FeII) ion.

Answer 2

The actual question is: A sample is analyzed to determine its iron content (as Fe2+0 via a redox titration with potassium chromate as the titrant. In the titration Fe2+ is oxidized to Fe3+ and CrO4(2-) is reduced to Cr3+. What is the percent by mass of iron in the sample if 0.9087 g of the sample required 45.68 mL of a 0.04322 M K2CrO4 solution to reach the endpoint?

I figured out moles of the titrant. I know that I need to balance the reaction by using the half-reaction method. I just don't know what the the second half looks like. I have:
Fe2+ goes to Fe3+ and e-
(some number of e-) and 8H+ and KCrO4 goes to Cr3+ and 4H2O and (2K??)

So I'm a little lost as to what to do next...

Thanks!