What is the balanced net ionic equation for the reaction between the permanganate ion and the ferrous ion in acid solution. What is the number of moles of ferrous ion that could be oxidized by 1.2×10^-2 moles of permanganate ion?

Honestly don't understand a single thing this question is asking me . All help needed!

The two half equations are

Fe^2+ ==> Fe^3+ + e
MnO4^- + 8H^+ + 5e ==> 4H2O + Mn^2+
Multiply equation 1 by 5 and equation 2 by 1 and add them. Cancel the electrons and you are left with the net ionic equation.

Use the coefficients in the balanced equation to convert 1.2E-2 mols MnO4^- to mols Fe^2+.

Don't worry, I'm here to help! Let's break down the question step-by-step.

1. Balanced Net Ionic Equation:
To determine the balanced net ionic equation, we need to first write the balanced regular equation for the reaction. From the given information, we know that the reactants are the permanganate ion (MnO4-) and the ferrous ion (Fe2+), both in acid solution.

The balanced regular equation for this reaction is:
MnO4- + Fe2+ + H+ -> Mn2+ + Fe3+ + H2O

Now, we can write the balanced net ionic equation by eliminating any spectator ions. In this case, the H+ ions are spectator ions because they appear on both sides of the equation. Therefore, we can eliminate them to get the net ionic equation:

MnO4- + Fe2+ -> Mn2+ + Fe3+

So, the balanced net ionic equation for the reaction between the permanganate ion and the ferrous ion in acid solution is MnO4- + Fe2+ -> Mn2+ + Fe3+.

2. Number of Moles:
To determine the number of moles of ferrous ion (Fe2+) that could be oxidized by 1.2×10^-2 moles of permanganate ion (MnO4-), we need to use the stoichiometry of the balanced equation.

From the balanced equation, we can see that one mole of permanganate ion (MnO4-) reacts with five moles of ferrous ion (Fe2+). Therefore, the ratio is 1:5.

Now, we can use this ratio to calculate the moles of ferrous ion:
Moles of ferrous ion = Moles of permanganate ion x (5/1) = 1.2×10^-2 x 5 = 6×10^-2 moles.

Therefore, 1.2×10^-2 moles of permanganate ion could oxidize 6×10^-2 moles of ferrous ion.

I hope this helps! Let me know if you have any further questions.

No worries! I'm here to help break it down for you.

First, let's start with the balanced net ionic equation for the reaction between the permanganate ion (MnO4-) and the ferrous ion (Fe2+) in acid solution. A net ionic equation only includes the species that participate in the reaction, while eliminating any spectator ions.

Steps to find the balanced net ionic equation:
1. Write the balanced molecular equation for the reaction:
MnO4- + Fe2+ → Mn2+ + Fe3+

2. Split the soluble ionic compounds into their respective ions:
MnO4- + 8 H+ + 5 Fe2+ → Mn2+ + 4 H2O + 5 Fe3+

3. Eliminate the spectator ions (ions that do not participate in the reaction):
MnO4- + 8 H+ + 5 Fe2+ → Mn2+ + 4 H2O + 5 Fe3+

So, the balanced net ionic equation for the reaction between the permanganate ion and the ferrous ion in acid solution is MnO4- + 8 H+ + 5 Fe2+ → Mn2+ + 4 H2O + 5 Fe3+.

Now, let's move on to the second part of the question, which asks for the number of moles of ferrous ion that could be oxidized by 1.2×10^-2 moles of permanganate ion. To answer this, we can use the stoichiometry, which relates the moles of one substance to the moles of another substance in a balanced equation.

From the balanced net ionic equation, we see that the stoichiometric ratio between permanganate ion (MnO4-) and ferrous ion (Fe2+) is 1:5. This means that for every 1 mole of MnO4-, we need 5 moles of Fe2+.

Given that we have 1.2×10^-2 moles of MnO4-, we can calculate the number of moles of Fe2+ using the stoichiometric ratio:

(1.2×10^-2 moles MnO4-) * (5 moles Fe2+ / 1 mole MnO4-) = 6×10^-2 moles Fe2+

Therefore, 1.2×10^-2 moles of permanganate ion could oxidize 6×10^-2 moles of ferrous ion.