In a general chemistry lab protocol (green chemistry, ISBN-13: 978-1482230208), students determine the content of ascorbic acid in a vitamin C tablet. There is an ascorbic acid standard in a buret, and there is also some ferric ammonium sulfate solution at known concentration, i.e. iron(III). The indicator is salicylic acid, which forms a colored complex with iron(III), but not iron(II).
The suggested protocol is a back titration where you first add a known excess amount of ferric salt to the unknown amount of ascorbic acid and add the indicator, giving a colored solution. Then, you titrate back with the ascorbic acid standard. The endpoint is the disappearance of color.
Is there are reason for the back titration in this case? In other back titrations, there is a step after adding an excess of reactant (for calcium carbonate, you add an excess of HCl and then drive out the carbon dioxide with heat, making the back titration a strong acid/strong base reaction). For the ascorbic acid reacting with ferric salt, I wonder if the endpoint is easier to see, or maybe the kinetics of complex formation motivate this (it says to let the solution sit before starting the back titration).