Write 2-3 sentences describing and identifying the type of chemical reaction that occurs when Fe2+ reacts with potassium permanganate (KMnO4) and specifying the type of solution required for the reaction to occur.
b. using standard reduction potential table, identify the oxidation and reduction half-reactions for the reaction.
c. write the balanced net ionic equation for the reaction, and identify which substance is oxidized and which is reduced.
d. suppose at the endpoint of the reaction 0.030 moles of KMnO4 were added to the analyte. How many moles of Fe2+ were contained in the beaker?
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b. The reaction between Fe2+ and potassium permanganate (KMnO4) is a redox reaction. It can be classified as a single replacement reaction. The solution required for the reaction to occur is an acidic solution because the reaction takes place in an acidic medium.
c. To identify the oxidation and reduction half-reactions, you can use the standard reduction potential table. The half-reactions for this reaction can be determined as follows:
Oxidation Half-Reaction: Fe2+ -> Fe3+ + e-
Reduction Half-Reaction: MnO4- + 8H+ + 5e- -> Mn2+ + 4H2O
d. To write the balanced net ionic equation, combine the two half-reactions above:
5Fe2+ + MnO4- + 8H+ -> 5Fe3+ + Mn2+ + 4H2O
In this reaction, Fe2+ is oxidized to Fe3+ (losing an electron) while MnO4- is reduced to Mn2+ (gaining 5 electrons).
d. To determine the number of moles of Fe2+ contained in the beaker, we need to use the balanced equation. The stoichiometric ratio between KMnO4 and Fe2+ is 1:5. Since 0.030 moles of KMnO4 were added, the number of moles of Fe2+ can be calculated as:
0.030 moles (KMnO4) * (5 moles Fe2+ / 1 mole KMnO4) = 0.15 moles of Fe2+.