I'm trying to figure out how to decide the oxidation states of some basic individual atoms in a compound/determine which is a reducing agent and which is an oxidizing agent, but I'm not sure.

Question

mendicant_bias · Answer

I'm supposed to determine what is a reducing agent and what is an oxidizing agent in $$\sf Si(s)+2Cl_{2}(g) \rightarrow SiCl_{4}(l)$$
What I don't understand is that in all previous problems, I was able to use common rules like Oxygen's oxidation state almost always being 2-, or anything in the first two groups on the periodic table having an identical oxidation number.

In this situation, both Si and Cl are on the same side of the periodic table, and I'm not aware of any special rules that govern how they behave with one another. How do I solve this? @Hoslos

mendicant_bias · Answer

I'm guessing this is maybe determined by which thing is more electronegative?

anonymous · Answer

Yes brother.
As I said before, every element has a 0 oxidation number, except when it is in a compound, the value is different from zero.
So in the LHS they are both 0 and after, Cl is -1, because generally, G7 elements in compounds have oxidation number as -1. All you have to do now is find for Si, making an equation, equating it to 0. So you have x-1=0, so x is 1. Therefore Si is 1.
Now, Let me give you a note: follow the acronyms OILR-RIGO
OILR - Oxidation is Loss of Electrons performed by Reducing agents.
RIGO - Reduction is the gain of electrons performed by Oxidizing agents.
Si has changed from 0 to +1, meaning it lost an electron to become positive, so it is a Reducing agent.
Cl changed from 0 to -1, meaning it gain an electron to become negatively charged, therefore it is an oxidizing agent.

mendicant_bias · Answer

Yeah, I understand the rest of that, it was just hinging entirely on the assumption that electronegativity determines things when everything else isn't clear. Thanks.

anonymous · Answer

Right.