>> Welcome to a screencast that will provide an introduction to oxidation numbers.
Let's start by looking at a couple of chemical reactions.
This one is between magnesium and oxygen.
It's makes magnesium oxide.
This is an ionic compound, and we can explicitly show the charges
of the magnesium and oxide ions in this compound.
And when we do that it's reasonably easy to see
that the magnesium loses electrons to make Mg plus 2.
The oxygen gains electrons to make O minus 2,
and this is a redox reaction, oxidation and reduction.
Here's another example, zinc metal reacts with hydrochloric acid
to make zinc chloride and hydrogen gas.
And, once again, if we explicitly show the charges of the substances involved
in a reaction then we can see that the zinc loses electrons, it's oxidized.
The hydrogen ion gains electrons and it's reduced.
Again, we have an oxidation reduction reaction.
And in general, of course, if a substance is oxidized it's going to lose electrons
which means its charge increases.
When a substance is reduced it gains electrons and its charge decreases.
Now, in examples like the ones we just did they're pretty easy to deal
with because we're dealing with atoms and ions in ionic compounds
and the charges are well defined, and it's very easy to track gain and loss
of electrons and change of charge.
But if we have molecular compounds then things aren't quite as simple.
We don't have well defined charges even though we still have oxidation reduction.
So in molecular compounds even though oxidation reduction occurs, electron transfer is harder
to track, it's not as clear cut as when we have only atoms and ions.
And so what chemists have done is we've developed a concept called oxidation number.
And this is something we use to describe and to track what's going on in terms of oxidation
and reduction in terms of gain and loss of electrons even in molecular compounds
where the charges aren't that clear cut.
And how this works is what we do is we assign to each element in a compound the charge
that the atoms would have if the compound were ionic.
So note this is kind of an important thing to be aware of, we treat the compounds as ionic
in terms of how we assign oxidation numbers even though they aren't ionic.
So this is convention we use, it's not what's actually happening.
These charges don't really reflect charges of ions.
Now, of course, the stuff we've talked about before is still true.
Oxidation does still correspond to losing electrons which means the oxidation number
or the charge increases, and reduction corresponds to gaining electrons
or decreasing the charge or decreasing the oxidation number.
And, again, remember if this seems a little bit backwards it's
because electrons are negatively charged.
So less electrons means less negative charge, higher positive charge.
And more electrons means lower or more negative charge.
Okay, now here are the guidelines that we use to assign oxidation numbers.
For a given species its oxidation number is assigned in the following way,
and we'll do several examples here of types of substances and specific examples.
So, first of all, if we're dealing with an element --
sorry, an atom in its elemental state, the oxidation number is 0.
Now, remember when we're talking about oxidation number,
we're essentially saying what is the charge on the substance even if we're doing sort
of a pretend charge by treating it as ionic when it's not really ionic.
But in this case, atom and elemental state, you're dealing with the element carbon.
It's oxidation number is 0.
Even if it's diatomic, so hydrogen, H2,
each of the atoms in there have an oxidation number of 0.
And our convention is we write oxidation numbers above the symbol for the element.
Okay, monatomic ion, the oxidation number is the charge of the ion.
So if we are dealing with Ag, silver, a positively charged ion,
then the oxidation number is plus 1.
If we're dealing with a sulfide 2 minus ion, then the oxidation number is minus 2.
So in a molecular compound an atom gets the charge that it would have
if it were a monatomic ion even if it's not an ion.
And we'll look at several examples of this.
For example, hydrogen is plus 1 when combined with non-metals,
minus 1 when combined with metals.
Oxygen is usually minus 2, but it's minus 1 in a few exceptions known as peroxide.
And then it can have various positive numbers when combine with fluorine.
So let's do an example here.
H2O is a molecular compound.
It consists of hydrogen and oxygen.
These aren't ions, these are molecular compounds covalently bonded.
But we treat them as if they would be ions
in which case H would be a plus 1 ion, and O would be a minus 2 ion.
And in water note that the oxidation number of oxygen is minus 2,
oxidation number of hydrogen is plus 1.
But there's two hydrogens so that gets a total charge of plus 2
which balances off the minus 2 charge from the oxygen.
Now, here's an example of the exception.
H2O2 is a peroxide, hydrogen peroxide in fact.
And notice if we had an oxidation number of plus 1 for hydrogen and minus 2
for oxygen, that wouldn't add up to 0.
And so in these cases, this is an exception, hydrogen has an oxidation number of plus 1,
and oxygen has an oxidation number of minus 1.
Notice this is not that common.
You're not going to run across too many peroxides,
but if you do then that's an exception to the minus 2 for oxygen.
The other rule that we'll use is that halogens are always negative 1 if it's fluorine,
and then usually negative 1 in other instances with other --
for other halogens unless it's combined with oxygen or some other halogen as well.
And note these are the charges that the ion normally forms.
Just like O minus 2 is the normal for oxygen, hydrogen plus 1 is the normal for hydrogen.
And as has been alluded to or mentioned a couple times so far, the sum of the oxidation numbers
of all the atoms in a molecule or a polyatomic ion equals the charge on the molecule or ion.
So we can do a couple examples of this.
We could do something like maybe the sulfate ion,
SO4 minus 2 charge, and we have sulfur and oxygen.
A rule is that oxygen is usually negative 2 unless it's combined with --
unless it's peroxide which we don't have or it's combined fluorine which we also don't have.
So oxygen is minus 2, that's what we say its oxidation number is.
But then notice we have 4 oxygens so that's a total charge for the oxygens of minus 8.
And then whatever the oxidation number or charge for sulfur will be it has to add
up with the minus 8 to do to the oxygens to give minus 2 total charge on the ion.
And so, of course, the number that works there is plus 6.
So we'd say the oxidation number on the sulfur is plus 6.
Okay, one more, how about the compound OF2.
And we'll notice here we have two rules, oxygen usually minus 2, fluorine always minus 1.
So if we did oxygen when it's usually minus 2 and fluorine what it is as an ion minus 1,
it's not going to be able to add up to 0.
But fluorine trumps oxygen so to speak, so that has an oxidation number of minus 1.
That means since there's 2 fluorines --
sorry, there's 2 fluorines so that's a negative 2 total charge, and this has to add
up to 0 overall for the compound, that means that oxygen has to be a plus 2 to make it work.
So note oxidation numbers can be different for a compound --
so for an element in compound under different situations.
Oxygen is not always minus 2, not even always minus 1.
It can be plus 2 as well.
All right, so let's do an example of a reaction that might be familiar to you,
the reaction of copper to oxide with methane gas to make copper, carbon monoxide and water.
And, first of all, let's assign oxidation numbers.
And we'll just go from left to right.
In copper oxide we expect the oxygen to have an oxidation number of minus 2.
That means the copper has a oxidation number of plus 2.
And these, of course, just correspond to their charges as ions.
In the methane we don't in general know what the oxidation number of carbon will be,
but we predict since it's a non-metal with a non-metal that hydrogen will be plus 1.
Since there's 4 hydrogens that means that the carbon has to be negative 4
for its oxidation number to balance off the charge and have it all add up to 0.
And notice we're writing oxidation numbers for one of whatever element we're talking
about in this case, hydrogen and carbon.
We write them above the element symbol, and the total has to add up to in a compound 0.
But the charges or the oxidation numbers are for individual atoms.
Okay, copper is the element so charge or oxidation number is 0.
Carbon monoxide we expect oxygen to have an oxidation number of minus 2 most of the time.
This is not an exception, therefore carbon is plus 2.
And then in water, we did this one previously, hydrogen's plus 1, oxygen's minus 2,
and that adds up to zero as expected.
Okay, so what we note now is that the copper's oxidation number went from plus 2 to 0.
And the carbon's oxidation number went from negative 4 to plus 2.
And the other substances, the oxygen and the hydrogen, their oxidation numbers didn't change.
So the substances that have their oxidation numbers change are the ones
that are oxidized and reduced.
For carbon the oxidation number increased from negative 4 to plus 2.
So increase here means becoming more positive.
It does not mean magnitude.
And so increase in oxidation number corresponds to oxidation
so we say copper is -- sorry, carbon is oxidized.
For copper the oxidation number decreases from plus 2 to 0.
And so we'd say the copper is reduced.
A thing to note as well, the oxidation number for carbon changed by 6,
from negative 4 to plus 2 that's an increase of 6.
The oxidation number for copper only went down by 2, from plus 2 to 0 is a decrease of 2.
But notice there's 3 copper involved in the reaction, and so 3 times 2 is 6.
And you'll always get total oxidation number change being the same value in terms
of magnitude for the increase in the oxidation and the decrease in the reduction
if you consider all the atoms involved.
Now, going back to something we've seen before,
the substance that is oxidized is the reducing agent,
and the substance that's reduced is the oxidizing agent.
So copper to oxide, CuO, is the oxidizing agent.
And methane, CH4 is the reducing agent.
The agents are going to be the whole substance involved in the process.
And that is it for our introduction to oxidation numbers.