Homework 14.3.

[Cobie_Allen_1H]

Hey guys! I was working on the homework and I honestly was just really confused on how to balance redox equations. Could you help out?

Question 14.3. states " Balance each of the following skeletal equations by using oxidation and reduction half-reactions. All the reactions take place in acidic solution. Identify the oxidizing agent and reducing agent in each reaction.

b. MnO₄^-(aq)+ H₂SO₃(aq) --> Mn²⁺(aq)+HSO₄^-(aq)"

Thanks so much.

[Maggie Bui 1H]

The first step should being figuring out what's oxidized and what's reduced. The straightforward way is to go through each reactant and its corresponding product and see how each element in it changes oxidation states, but since we know that transition metals commonly change oxidation states, I think it'd be a good idea to check those first if there are any in the equation.

In this case, there is one: manganese (Mn). We can calculate the initial oxidation of manganese by looking at the ion's overall charge and the other elements in the ion. Oxygen has an oxidation state of 2- and there are four of them, so 4 x 2- is 8-. The overall ion's charge is 1-, so take that into account: (8-) - (1-) = 7-. This must be balanced by manganese to ensure that the ion's overall charge is 1-, so the manganese has an oxidation state of 7+. We see that the manganese on the products side has a charge of 2+, so we see manganese go from 7+ to 2+, a gain in electrons, so it is a reduction.

Repeat this process to find what's oxidized; you're looking for something that changes oxidation states (increasing oxidation states to be specific). After that, you can separate the equation into redox half-reactions, balance them, and then recombine them for a final answer. I recommend reading and doing the examples on pages 562 to 569 in our textbook for the specifics, but that is essentially what you do.

[Cobie_Allen_1H]

Wow, thank you so much! That helped a lot!