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Half-reactions 12,073 views
Half reactions are used for balancing oxidation - reduction reactions that occur in aqueous solutions. The sum of the oxidation and reduction half reactions forms the overall equation for the full reaction. To add half reactions, it is necessary for the number of electrons lost to equal the number of electrons gained.
Hi guys. So let's continue talking about oxidation reduction reactions, but discuss kind of a more specific sub-heading of that known as half reactions. So half reactions are a method for balancing oxidation reduction reactions that occur in aqueous solution. So remember before oxidation reduction, one cannot happen without the other. So if you have oxidation, you have to have reduction and vice versa.
So here in the half reactions, one of the half reaction will illustrate reduction and in illustrating reduction, that means the electrons denoted here as an e- will show up on the reactant side. If I was writing a hypothetical equation, sorry, reaction, I would say a plus an electron gives me a-. And that will illustrate that a has been reduced.
In the other half reaction, for oxidation, the electrons are going to show up on the product side. So that means if I had b, it would go to b+ plus an electron and if I added these two equations up, I would get a+b on my reactant side forma product ab, where it's an ionic species and a has a minus charge and b has a plus charge. Right. So overall in that, zero charge.
The potential pitfall when doing these problems is going to be that you need to remember that the number of electrons lost by the oxidised species must equal to the number of electrons that are gained by the reduced species. So if we go back here again, we'll see that I lost one electron and, sorry. That I gained one electron on my reactant side here from my reduced species and lost one electron here on my product side for my oxidised species. So let's do a quick example.
Before you even delve into this, I want you to remember and put it at the forefront of your mind, OILRIG. So that you remember the oxidised species is losing electrons and the reduced species is gaining electrons.
So here we have on our reactant side tin 2+ aqueous plus two iron with a 3+ charge, also aqueous forming tin 4+ it's aqueous plus two iron 2+ aqueous. Okay, so this is the equation that we've been given and we need to write the two half reactions. So basically indicate which of your reactants is being oxidised and which is being reduced.
So here if we look at tin, it starts as 2+ and it goes to 4+ which means it must have gained two electrons. Right? So that means the electrons show up on our product side and tin is being reduced. Sorry, oxidised. So here our iron, we have two iron 3+ to start 2 iron 2+ to end with which means on our reactant side we have a 6+ charge and on our product side we have a 4+ charge which means that we must have gained these two electrons here on the reactant side meaning that our electrons are reactants and iron has been reduced. And that is basically half reactions.