It's all about who has it, who wants it, and what they're willing to do to get it.

Electrons are what make it possible for an atom to bond with other atoms to form molecules,

and when that happens a tremendous amount of energy can be

exchanged in the process.

Not all chemical reactions involve electrons changing hands.

Acid base reactions, you'll recall, are more about swapping protons but

because electrons are the real coin of the realm, in chemistry the most important reactions

that take place on Earth involve the transfer of one or more electrons from one

atom to another. These are redox reactions. Redox which is the portmanteau

of reduction oxidation.

Well what's up with those words? You know what reduction is: making less of

something and then oxidation maybe have something to do with Oxygen...well sometimes but not always.

These are actually super terrible choices for what is actually

happening in redox reactions but we are stuck with them. Reduction is when a

substance gains electrons. Yes it gains which is the opposite of what the word

reduce means, fantastic, and yes sometimes I want to punish the people who named these

things so inaccurately but they didn't know any better and they're all dead so we

can't do anything about it.

Protochemists would make pure metals by heating or smelting their ores and they

noticed during the smealting these substances would become lighter so I guess

it's not crazy that they decided to say that these substances were being reduced.

Our old French friend Antoine Lavoisier figured out that this was because Oxygen

gas was actually leaving the compound and making it lighter. What he didn't know

was the actually chemistry involved. Oxygen is unsuprisingly the quintessential

oxidizer. It pulls electrons off of one molecule to make itself more stable but if

you heat it up enough it gets all energetic. Today we understand that

oxidation and reduction are all about electron transfer so you might think that

we'd rename them and some chemist have tried producing terms like electronation,

and deelectronation but once a set of terms is decided upon and used for a while

it's pretty difficult to uncreate it, so we're stuck. To keep these seemingly

nonsensical names straight I rely on the phrase OILRIG: Oxidation is loss of

electrons, reduction is gain of electrons. We just got to know this stuff because

it's everywhere. When your cells convert sugar into energy so you can move, and

breathe and think that's redox. When plants photosynthesis sunlight into food

that's redox. The battery powering your laptop? Redox. Fire? Also redox.

Since electron swapping is the name of the game here when you study redox reactions

it's important, critical, absolutely essential to keep track of the electrons.

Think of them as dollars or pesos or pounds or euros. In any transaction, one

person is going to gain them and the other is going to lose them and to stay on top

of things you have to keep tabs on who's ahead and who's behind. Atoms are fond of

sharing electrons though forming covalent bonds so sometimes keeping track of where

they are and where they're going to end up isn't quite so simple. I like to think of

every covalent compound like a marriage. Though it's going to be a weird marriage

because it might be like six people in it. Sometimes the same person several times

without commitment and also no emotions. Don't think too much about it. Like in a

marriage where money gets shared, covalent compounds share electrons. The trick is figuring

out who gets the cash when the vows break. So we've created a useful little system

assigning electrons one hundred percent to atoms that are actually at the moment

sharing them. The number that we assign is the atoms oxidation state or oxidation

number. Even though we are of course aware of covalent bonds in the sharing of

electrons the processes are easier to follow if we imagine the atoms are already

splitting up the bank account as if they were in an ionic or non-sharing bond.

So an atom's oxidation number is basically what it's charge would be if it actually

owned all of it's electrons exclusively like the newly mented bachelors they may

become. So, to figure out those oxidation states or those oxidation numbers we have

some simple rules for some atoms. First, the oxidation number for any element by

itself whether it's monatomic or diatomic or polyatomic like an atom of Calcium or a

molecule of H2 or an even bigger molecule of Sulfur S8, the oxidation number is zero

Atoms by definition do not have a charge. If they had a charge they would be ions

and if they're sharing with themselves they share equally. Second, for a

monatomic ion, basically a charged atom, it's simply the size or number of it's

charge so the Iron two in Fe2 plus has an oxidation state of plus two, while a

Chloride ion is minus one. Third, oxygen which is unsuprisingly all over redox

chemistry almost always has an oxidation of negative two, unless it happens

to be in a peroxide molecule like Hydrogen Peroxide. Fourth, Hydrogen is plus one and

fifth, Flourine is negative one as are all the other halogens most of the time, pretty

much unless they're bonded to Flourine or Oxygen because Flourine and Oxygen are so

bad that they can make anybody's oxidation number positive if you know what I mean.

And those are the rules. Now the total of all the oxidation numbers of all the atoms

in a neutral compound will add up to zero. Like water with one Oxygen with a negative

two oxidation state and two Hydrogen at plus one and viola the neutral compound

has an oxidation number of zero. A polyatomic ion on the other hand has to

work out to have an oxidation state that matches it's charge. So SO42 minus the

Sulfate ion has four Oxygens for a total of negative eight but we don't have a rule

for Sulfur so I guess we just give up and walk away, who cares anymore. No, we use

like third grade algebra because we have to end up with an oxidation number of

negative two for the whole compound. We know that Sulfur in this particular

compound has an oxidation state of plus six but Sulfur's oxidation state isn't

always plus six and that's why we don't have a rule for Sulfur or a lot of other

elements for that matter because oxidation states of most elements change depending

on what they're bonded with. Now we can use this same logic to figure out what

happens when these compounds interact in redox reactions. Molecular divorce courts

of electrons changing hands, being haggled and treated with some players taking big

profits while others lose nearly everything. Let's start out with a simple

example: a chemical reaction that I believe has saved more lives than any

other in the history of chemistry created by a war criminal to blow people up during

World War One: the Haber Process. The Haber process removes the ultra stable

Nitrogen from the air and combines it with Hydrogen to form NH3 Ammonia for use in

bombs and also in fertilizer increasing the carrying capacity of the Earth by

billions. Nitrogen in the air exists as elemental diatomic Nitrogen and Hydrogen

likewise is also diatomic H2. So we know that starting out, all of the atoms have an

oxidation state of zero. The product of the reaction Ammonia is a neutral compound

with one Nitrogen and three Hydrogens. The Hydrogens each have an oxidation state

of plus one. Remember the rules, so nitrogen must have an oxidation state of minus

three. Nitrogen thus gained electrons. It's oxidation state went down and so it

was reduced. So at least we're talking about what oxidation states are doing.

The word reduced makes sense. Hydrogen lost electrons, the oxidation state went up and

it was oxidized. Now this is a pretty simple equation of balance but redox

equations can be a huge headache sometimes because of the number of the individual

atoms involved so we often have to balance them in half reactions. So even though we

don't really need to do the half reactions, because this is a pretty simple equation,

we're going to do them anyway just because it's an example that's simple to start with.

So start out with the Nitrogen getting reduced. We have N2 with an

oxidation state of zero becoming NH3 with an oxidation state of negative three.

First we balance the number of Nitrogens, the add the number of electrons we need to

have to have there be the same number of electrons on each side. Do the same with

the oxidation half of the reaction and then combine them to get your whole

reaction with electrons cancelling out. Now you ask, maybe that seemed like an

unnecessary step, but allow me to show you a more complicated example that will prove

how necessary it may be. In this flask is silver diamine. We're going to use

some redox chemistry to get the elemental silver out of it nice and clean and shiny

and it's not going to be no simple Haber process. The silver diamine is going to

react with an organic aldehyde any aldehyde actually. The business end of

the aldehyde is the CHO end the R in organic chemistry is a symbol for some

organic group of atoms and in this reaction those atoms don't matter.

The silver diamine reacts with the aldehyde and the Hydroxide creating

carboxylic acid, ammonia, and water. First let's assign ourselves some oxidation

states. The silver is in a complex with two neutral ammonias that are going to

remain unreacted throughout the equation so we can treat them like a single species

with an oxidation state of zero. Since the silver diamine has a charge of plus

one and the ammonia's don't affect that Silver's oxidation state must also be plus

one. The aldehyde has one Hydrogen at plus one and one Oxygen at minus two but

is neutral over all so the Carbon must be plus one as well. The hyrdoxide ion is

simple minus two for the Oxygen and plus one for the Hydrogen and an overall charge

in thus oxidation state of minus one. On the reactant side Silver is now atomic so

it's oxidation state is zero. The carboxylic acid has two Oxygens and one

Hydrogen so the carbon now has an oxidation state of plus three. NH3 remains

at zero and the Hydrogen and Oxygen of water also haven't changed oxidation

states. So, silvers oxidation state decreased or was reduced from plus one to

zero while carbon was oxidized from plus one to plus three. Half reaction time.

Silver was reduced gaining one electron forming elemental silver and ammonia from

silver diamine. The aldehyde was oxidized forming carboxylic acid and requiring two

electrons. With the help of those electrons we know that at the very least

we have to double the reduction half of the equation entirely in order to get the

right number of electrons on both sides. we do that and oh God that's good stuff

and we combine them together for a perfectly balanced redox equation.

And now watch me take those electrons and turn them into money and there you

have it folks that is pure silver coating the inside of the flask.

Thank you for watching this episode of Crash Course Chemistry.

If you were paying attention you learned that any reaction where electrons move

around from atom to atom is a redox reaction. That oxidation is the laws of

electrons and that reduction is the gain of electrons and an oxidation numbers

are assigned to take part in reactions in order to keep track of what their

electrons are up to. You also learned a few simple tricks to help figure out what

an atoms oxidation state is and you got a little practice figuring out how to assign

oxidation states and balance oxidation reaction with two examples. One pretty

simple and another a little less so. This episode of Crash Course Chemistry was

written by Kim Krieger and myself. Our script editor was Blake de Pastino.

Our chemistry consultant is Dr. Heiko Langner and a troop of chemistry teachers

also advised and edited this one so thanks very much to all of them. This episode

was filmed edited and directed by Nicholas Jenkins our sound designer and script

supervisor is Michael Aranda and our graphics team is Thought Cafe.