How does oxidation work

And that intellectual tool is the idea of oxidation states. What the oxidation state is, even if you're in a situation where you have covalent bond, you say, well, look, I understand. Those are partial charges. These are covalent bonds. The electrons are being shared. But I don't like this partial stuff. I want to just assume hypothetically, what if these were ionic bonds? And you say, well, if these had to be ionic bonds, then the oxygen would nab the electrons from these pairs. And so the oxygen would have a fully negative charge, a negative 2 charge.

And the hydrogens would have a fully positive charge each. And so, if we were to write down the oxidation states for the atoms in the water molecule-- let's write that down, so H2O-- we would say that oxygen has an oxidation state of negative 2, and each hydrogen atom has an oxidation state of plus 1. And notice, the whole molecule is neutral, and these things cancel out with each other.

Positive 1, positive 1, that gets you to positive 2. Then you have negative 2. They cancel out. Now, the one thing, I keep saying this is negative 2, but I wrote the negative after it. If I wanted to write positive 1 as an oxidation state, I would actually write it as 1 positive, although you can assume that if someone just writes the positive. And this is just the convention, to write the sign after the number when we are writing actually ionic charges or oxidation states, because an oxidation state is nothing but a hypothetical ionic charge.

If you really had to-- if you were forced to assume these aren't covalent bonds, but these are ionic bonds. Once again, I want to stress. This is the reality. These are partial charges, the oxidation state, intellectual tool, that's forcing us to pretend like these are ionic bonds. And you might say well, this kind of makes sense right over here. This involved oxygen in some way.

That's why it's called oxidation states. And that's how I initially conceptualized it when I first learned about this. You say, well, look, each of these hydrogens lost an electron to oxygen. So it makes sense that we say that each hydrogen got oxidized, so hydrogen oxidized by oxygen. It makes sense that oxygen would oxidize something else. This got done to it. The charge was taken away by oxygen, so it got oxidized. Now, the other term on the other side of oxidized is reduced.

And the word "reduced" really comes from the idea that oxygen's charge has been reduced. So we could say, O, or we could say oxygen has been reduced by the hydrogens. And so there there's a temptation here to say, well, OK, this must always involve oxygen in some kind, because it seems to begin with the same words. Well, that is not the case. Let's take, for example, if this is an aqueous solution, hydrofluoric acid right over here.

You have a hydrogen covalently bonded to a fluorine. Now, just like we saw in water, fluorine is one of the most electronegative elements. It's going to hog the electrons in this covalent bond.

So this is going to have a partial negative charge here. This is going to have a partial negative charge here. And this is going to have a partially positive charge. But if we were going to think about it in terms of oxidation states, we would say when push comes to shove, if this had to be at an ionic bond and not a covalent bond, what would be the charges on each of these atoms? We can say, well, in that case, hydrogen would lose an electron, and it would have a full positive charge.

Rust, patina, fire, rancid food — they all have oxidation in common. The same property that makes the gas vital to most Earth-borne life — its unquenchable thirst for electrons — slowly kills the very life it supports.

Oxidation is the process in which one atom strips electrons from another, claiming them for its own. It is one side of redox type reactions. Reduction is the process via which an atom ceeds electrons to another. The term draws its name from oxygen because it was the first known oxidative element.

A good example of this traditional definition for oxidation can annoyingly display itself on the body of our cars: rust iron oxide. Most elements can be oxidized, given proper coaxing, in a variety of environments. Many can be made to oxidize their peers. Some flake and break apart when oxidized, others tend to become more resistant to further oxidation. The process comes in many forms and involves many players.

For purely theoretical approaches, half-reactions can be used to explain half of a redox reaction — be it the oxidation or the reduction component. These are pretty helpful in simplifying the whole process, to make it easier to teach or understand. But keep in mind the first line: in real life, oxidation and reduction always come together. But having a more inviting host nearby to move on to can draw it out.

Think of it as a marketplace. Smelters or blacksmiths, I guess? But the name stuck. It loses electrons during oxidation. Some atoms have a greater electron negativity so have a greater attraction for electrons. Oxygen is an example of an element that has a high electron negativity and draws electron density away from other elements. When another element has its election density draw away from the element the elements is said to have been oxidized. The process of electron density transfer is named oxidation after Oxygen which tends to take electron density from almost every other element.

How does oxidation work? Chemistry Electrochemistry Oxidation and Reduction Reactions. David Drayer.