In a particular redox reaction, MnO2 is oxidized to MnO4– and Ag is reduced to Ag. Complete and balance the equation for this reaction in acidic solution. Phases are optional.
Sure, here's my attempt at balancing the equation for you:
2MnO2 + 4Ag → 2MnO4– + 4Ag
Now, isn't it interesting how manganese dioxide becomes permanganate and silver gets reduced to silver? It's like MnO2 said to Ag, "Hey Ag, I don't like being gray anymore. I wanna be purple permanganate instead!" And Ag was like, "Sure, MnO2! I'll lighten up and become shiny silver for you!"
Acidic solution? I guess they felt the need for some added excitement and drama in their reaction. Who am I to judge? Maybe they wanted a front-row seat to the acid party!
Hope this way of balancing makes it a little more fun for you!
To balance the redox reaction, we need to follow these steps:
Step 1: Assign oxidation numbers to all the elements in the reaction.
In MnO2, the oxidation number of Mn is +4. In MnO4–, the oxidation number of Mn is +7. In Ag, the oxidation number is 0. In Ag+, the oxidation number of Ag is +1.
Step 2: Write the half-reactions for both the oxidation and reduction processes.
Oxidation half-reaction: MnO2 -> MnO4–
Reduction half-reaction: Ag+ -> Ag
Step 3: Balance the atoms in each half-reaction.
For the oxidation half-reaction, we have:
MnO2 -> MnO4–
For the reduction half-reaction, we have:
Ag+ -> Ag
Step 4: Balance the oxygen atoms by adding water molecules.
For the oxidation half-reaction, we need to add water to the right side:
MnO2 + H2O -> MnO4–
Step 5: Balance the hydrogen atoms by adding hydrogen ions (H+).
For the oxidation half-reaction, we need to add hydrogen ions to the left side:
MnO2 + 4H+ + H2O -> MnO4–
Step 6: Balance the charges by adding electrons (e–).
For the oxidation half-reaction, we need to add 4 electrons to the left side to balance the charges:
MnO2 + 4H+ + H2O + 4e– -> MnO4–
Step 7: Balance the charges in the reduction half-reaction.
The reduction half-reaction already has the correct charges, so no adjustment is needed.
Step 8: Multiply each half-reaction by the appropriate factor to equalize the number of electrons transferred.
To equalize the number of electrons, we need to multiply the reduction half-reaction by 4:
4Ag+ -> 4Ag
So, the balanced half-reactions are:
Oxidation: MnO2 + 4H+ + H2O + 4e– -> MnO4–
Reduction: 4Ag+ -> 4Ag
Step 9: Combine the two half-reactions and cancel out the electrons.
To combine the half-reactions, we need to multiply them by appropriate factors so that the number of electrons cancels out:
4(MnO2 + 4H+ + H2O + 4e–) -> 4(MnO4–)
4(4Ag+ -> 4Ag)
Simplified, the balanced equation in acidic solution is:
4MnO2 + 16H+ -> 4MnO4– + 4Ag
To balance redox reactions, it is necessary to follow a few steps:
Step 1: Determine the oxidation states of the elements involved.
In this case, we are given that MnO2 is oxidized to MnO4–, which means that the oxidation state of manganese (Mn) increases from +4 to +7. Similarly, Ag is reduced to Ag, which means that the oxidation state of silver (Ag) remains the same at 0.
Step 2: Identify the species that are undergoing oxidation and reduction.
In this reaction, MnO2 is undergoing oxidation, and Ag is undergoing reduction.
Step 3: Separate the reaction into two half-reactions.
We can write the oxidation and reduction half-reactions as follows:
Oxidation Half-reaction: MnO2 → MnO4– (oxidation of Mn)
Reduction Half-reaction: Ag+ → Ag (reduction of Ag)
Step 4: Balance the atoms in each half-reaction.
In the oxidation half-reaction, there are two Mn atoms on the left side, so we need to add a coefficient of 2 in front of the MnO4– ion:
2MnO2 → 2MnO4–
In the reduction half-reaction, there is already one Ag atom on the left side, so it is already balanced.
Step 5: Balance the charges in each half-reaction by adding electrons (e–).
In the oxidation half-reaction, the MnO2 molecule has no charge, and the MnO4– ion has a charge of –1. Since each Mn atom gains 3 electrons in oxidation, we need to add 6 electrons (3 electrons per Mn atom) to the left side:
2MnO2 + 6e– → 2MnO4–
In the reduction half-reaction, the Ag+ ion has a charge of +1, and the Ag atom has no charge. Therefore, no electrons are needed in this half-reaction.
Step 6: Balance the number of electrons transferred in both half-reactions.
To make the number of electrons equal in both half-reactions, we need to multiply the reduction half-reaction by 3:
3Ag+ → 3Ag
Step 7: Combine the half-reactions.
By combining the oxidation and reduction half-reactions, the electrons cancel out:
2MnO2 + 3Ag+ → 2MnO4– + 3Ag
Step 8: Balance the remaining elements and charges.
In this equation, all the elements and charges are already balanced, so the final balanced equation for the reaction in acidic solution is:
2MnO2 + 3Ag+ → 2MnO4– + 3Ag
You didn't put a + charge on one of the Ag. It has to be on the left side
2H2O + MnO2 ==> MnO4^- + 3e + 4H^+
Ag^+ + e ==> Ag^+
Those are the two BALANCED half reactions. Multiply equation 1 by 1 and equation 2 by 3 and add them.