Question

1. Balance the equation (smallest number of coefficients) for the reaction between iron (II) and permanganate ion. Remember that the equation must balance as to both mass and charge.

2. How many moles of electrons are transferred in the balanced equation?

3. Which species is oxidized? Which species is reduced?

5. What is the indicator for this reaction? Explain.

Answer 1

You unstated equation is Fe^2 + + MnO_4^- rightarrow Fe^3 + + Mn^2 + First identify the oxidation state of each atom Left hand side Fe = + 2 Mn = + 7 0 = -2 Right hand side Fe = + 3 Mn = + 2 So due to the electron transformer than is a change in the oxidation state of each atom Fe: + 2 rightarrow + 3 rightarrow change of electron + 1 Mn: + 7 rightarrow + 2 rightarrow change of electron -5 Total invented in oxidation number = total deleveraging oxidation

2) 5 electrons are involved in the reaction.

3) LEO = Loss of Electron is Oxidation

GER = Gain of Electron is Reduction

From the above equation Fe (Iron) changes/ loss is electron from Fe²⁺ to Fe³⁺, so Fe is Oxidized (LEO)

Mn (Manganese) changes/ gains is electron from Mn⁷⁺ to Mn²⁺, so Mn is Reduced (GER)

4) For Redox titration, MnO4^- permanganate no need of indicator.

In MnO4^- permanganate case it’s an intense purple colour changes slowly to Mn²⁺, colorless state. It’s a reduction of Mn⁷⁺ ion to Mn²⁺ ion, the solution changes from dark purple to faint pink at the equivalence point.