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Summary
Core Theme
Redox reactions are chemical reactions characterized by a change in oxidation numbers of elements, involving the transfer of electrons. Understanding these reactions requires defining oxidation numbers and identifying oxidizing and reducing agents.
hi everyone Welcome to our next video
this video is going to be on redox
reactions our learning outcomes are to
be able to Define redox reactions and be
able to classify reactions as being
redox we're going to be able to compute
the oxidation States or oxidation number
for elements in compounds and then we're
going to be able to determine in a
reaction what are the reducing and
oxidizing agents
so in terms of defining what an
oxidation reduction or a redox reaction
is our definition is that it's a
reaction involving a change in oxidation
number for one or more reactant elements
so to start with we will need to be able
to identify what the oxidation number is
and we'll get to that in a little bit
but some more simplistic redox reactions
we can see involving the transfer of
electrons between species so an example
of this is the reaction given here on
the slide between sodium and chlorine so
we have sodium solid and chlorine gas
which then react to give us the ionic
compound sodium chloride
now this reaction involves a transfer of
electrons we can actually split the
reaction into two separate reactions
which we're going to call half reactions
we can write a reaction involving sodium
so here we have the reaction where we
have the sodium solid loses its
electrons to become the positively
charged ion sodium ion and then we also
can write a reaction for chlorine in
which chlorine gas
gains electrons to become the chloride
ion CL minus
so these are our two half reactions for
now when we can do this when we can
separate a reaction into its two half
reactions we'll be able to identify the
two parts of our redox reaction so we
can call this redox it really stands for
oxidation and reduction so this tells us
that one of our half reactions is going
to be oxidation and the other is going
to be reduction
um we'll learn here that oxidation is
the loss of electrons so if we look at
the reactions listed here we can see
that sodium right starts as sodium solid
becomes the ion and is losing two
electrons in this process so this is the
loss of electrons so this must be our
now reduction is defined as the gain of
electrons and we can see that happening
in our second half reaction right the
chlorine gas gains the electrons so that
it becomes the chloride ion so this is
so we can identify the half reaction for
oxidation we can identify the half
reaction for reduction another way for
us to talk about what's happening in
these reactions is to say that the
sodium is oxidized right so um
the process of the oxidation reaction
right sodium is being oxidized and the
chlorine gas the cl2 is reduced
now there's a lot going on here and it's
often hard to remember the the
definitions for oxidation reduction be
able to identify these there's a
mnemonic device and it goes as this so
we've got Leo
so Leo the Lion says gur uh the parts
that I've capitalized are what are going
to help you to remember what oxidation
and reduction are
right so l e o for Leo loss of electrons
is oxidation and then the gur stands for
so g e r
so Leo the Lion says girl tells us that
loss of electrons is oxidation and gain
all right so we've learned how to
identify the oxidation reaction the
reduction reaction there's another way
for us to talk about what's happening in
these reactions we've written them as
our two separate half reactions but
really these two half reactions have to
be happening together right and when we
think about the process the sodium is
losing its electrons so that so that
those electrons can be given to the
chlorine and so they have to happen
together there's this relationship these
couldn't happen in isolation
so there's a relationship between the
sodium and the chlorine
so we can say that in an active sense a
reducing agent is the substance that
brings about reduction of another
substance right so we know that the
chlorine is being reduced right the
chlorine is being reduced the sodium is
the species that brings about that
process right the sodium is the reducing agent
agent
it's the substance that's bringing about
the reduction of the chlorine and in the
process we see that the sodium is itself
oxidized right it goes from sodium solid
to sodium ions
the oxidizing agent is the substance
that brings about oxidation of another
substance and in the process becomes
reduced so we can see that the chlorine
right is the species that is reduced so
because it is taking those electrons in
its reduction right it is essentially
helping to oxidize the sodium so our
chlorine gas is our oxidizing agent
so the sodium functions as the reducing
agent because it provides the electrons
for chlorine
and then the chlorine functions as our
oxidizing agent because it's there to
remove the electrons from the sodium
okay so these are just different ways
for us to talk about the process that's
now that was an example where there was
a clear transfer of electrons but there
are plenty of redox reactions where we
don't see that full transfer
so this is an example where we have a
reaction between hydrogen gas and
chlorine gas the product of this
reaction is a covalent compound so in
our previous example we had an ionic
compound so we knew that there were
charged species right so that there was
the full transfer of electrons so in
this case this isn't true
um but we can assign what we call an
oxidation number to each atom involved
in the reaction and the oxidation number
of an element in a compound is what's
written here the charge its atoms would
possess if the compound was ionic so
we're sort of giving a hypothetical
charge to the species involved in the
reaction assuming that there would be
some full transfer of electrons instead
of those electrons being shared right we
know that in covalent compounds
electrons are shared so this is a
hypothetical charge that we are
assigning to the species
so there's going to be some rules
involved with assigning oxidation number
the first states that the oxidation
number of an atom in its Elemental form
is going to be zero so we've seen some
examples of these already in the
equations that we've looked at or the
reactions that we've looked at right so
sodium solid is an element so this would
have an oxidation number of zero
chlorine gas
right is Elemental so that would be an
oxidation number of zero
number two the oxidation number of a
monatomic ion is equal to the ions
charge so if we had sodium ions
right this would have an oxidation
number of plus one if we had chloride
ions this would be -1 if we had calcium
ions right this would have an oxidation
number of plus two
you get the idea
all right so now we get a little bit
more involved there's oxidation numbers
that are usually assigned for certain
elements so we're going to start with
hydrogen we can see that hydrogen is
going to mostly be a plus one when we
are combining it with non-metals but it
would have a an oxidation number of
minus one when it's combined with metals
so some examples here if we are
combining with non-metals we could have
something like methane CH4 in this case
the hydrogen would have an oxidation
number of plus one but if it's combined
with Metals we say that the oxidation
number is -1 so this would be something
like lithium hydride right we know
lithium is a metal
so that would have a different oxidation number
number
next we move to oxygen now oxygen is
usually going to have an oxidation
number of minus two in most compounds
but there are some cases where we would
see a different oxidation number
the first example are going to be
peroxide so this would be something like
peroxide which is H2O2 right so the
oxygen here has an oxidation number of
minus one instead of minus two
also there are compounds called
superoxides where we would have an
oxidation number of minus one-half right
but these are going to be rare
and then we would see positive values
for oxygen when it is combined with
fluorine so an example of this would be
a compound like dioxygen
difluoride so O2 F2
um but again in most cases we're going
to see the minus 2 oxidation number for oxygen
oxygen
and then lastly we have roles for
halogens so if we're looking at a
compound involving fluorine fluorine is
always going to have an oxidation number
of minus one
but our other halogens could see other values
values
so usually these are going to be -1 for
all of our other halogens except when
they are combined with oxygen or other
halogens so these would be some exceptions
and then our final rule is that the sum
of our oxidation number for all atoms in
a molecule or a polyatomic ion has to
equal the charge on the molecule or the
ion so I had written CH4 before right if
we find the oxidation numbers for the
carbon and the four hydrogens and we add
them up we can see that this is a
neutral molecule so that the that sum of
the oxidation numbers should be zero but
if we had an ion say we had sulfate
right then the oxidation numbers added
together should give rise to the total
so now let's do some examples using the
guidelines that we've learned
um so we have this compound right
dihydrogen sulfide we're going to be
able to assign the oxidation numbers for
the two species in our compound and
because we have a certain number of each
of these atoms contributing to the
molecule we'll have to think about the
total contribution to charge in order to
make sure that this is going to overall
be neutral so I'll show you what I mean
so we know that we said that hydrogen
should have an oxidation number of plus
one when it is in a compound with
non-metals so sulfur is a non-metal so
we assume that this should have a plus
one charge now we see that we have two
hydrogens in this compound which means
that the total contribution to charge by
the hydrogens should be plus two so I
simply multiply the oxidation number for
each of the hydrogens by that subscript
to get plus two
now I know that this is a neutral
molecule so I know that when I take the
sum of these two boxes that it should
equal zero because this is neutral so
that tells me that the sulfur should
have a minus two total contribution to charge
charge
and since I only have one sulfur that
minus two is also my oxidation number
for that individual sulfur let's try
another example so here we have
sulfite so so32 minus we can start by
assigning our oxidation number for
oxygen right we know that oxygen is
usually minus two and sulfur it isn't
one of those exceptions that we saw on
our rules
so if oxygen has an oxidation number of
minus two and we have three oxygens the
total contribution to charge would be
minus two times three or minus six
now we know that the total charge on
this ion is going to have to equal minus
two so when we add these two boxes it
should give us minus two
so that tells me that the total
contribution to charge by the sulfur
should be plus four right so when I add
4 and take away 6 that's going to give
me minus 2.
and looking then to identify the
oxidation number for sulfur I can see
that there's only one sulfur so it
should have the same number as the total
contribution to charge so this would
also be plus four so my oxidation number
for The sulfur and this ion is going to
be plus four and then each of the
oxidation each of the oxygens is going
all right let's do another example we
have sodium sulfate
um and here we have the ability to look
at this together as an ionic compound or
we'd be able to separate it into its
cation and anion to perhaps more easily
identify the oxidation numbers
but we saw in our rules is that if we
have a monatomic ion something like
sodium then we can simply say that it's
oxidation number is equal to its charge
and we know that sodium has a plus one
charge when it's an ion and so that
should be our oxidation number we see
that we've got two sodium ions so the
sodium's total contribution to charge
would be plus one times two or overall
plus two
now if I want to start to look at the
sulfur and the oxygen I know that my
rules told me that oxygen should
normally be a minus two so I can assign
that oxidation number and since I have
four oxygens my total contribution to
charge is going to be 4 times minus 2 or
minus 8.
and then we want to make sure that the
total sum of my total contribution to
charge should equal My overall charge on
the molecule right as an ionic compound
the charges are balanced here so overall
this should be neutral or zero which
tells me that the sulfur's total
contribution to charge should be a
number that allows these to add up to
zero and we can see that that would be
an oxidation number of plus six
I only have one sulfur so the sulfur
carries this number as its oxidation
number all right so we've learned quite
a bit in this video about redox
reactions to start off we learned that
redox reactions are those reactions in
which the elements undergo a change in
oxidation number so if we want to
identify whether a reaction is an
oxidation reaction it's helpful for us
to First determine what those oxidation
numbers are and then be able to see if
one of them is changing right and then
that would tell us that we have a redox reaction
reaction
so earlier we considered this reaction
as a potential redox reaction so let's
look at it and kind of go through all
we've learned about redox so far
so the first question is is it a redox
reaction well we just talked about how
it's helpful for us to assign our
oxidation numbers and then if we see
that a particular element is changing
its oxidation number then that tells us
we have something that's a redox so
let's first assign all of our oxidation numbers
numbers
so here we have hydrogen H2 this is
hydrogen in its Elemental form so this
should have an oxidation number of zero
chlorine is also in its Elemental form
right and so this should also have an
oxidation number of zero
when we look at HCL we should be able to um
um
to know that hydrogen should have the
plus one oxidation number because it is
bonded to a non-metal and therefore the
chlorine should have a minus one in
order to make this compound neutral
so we do see that there is there are
changing oxidation numbers so hydrogen
is going from zero to plus one and
chlorine is going from zero to minus one
so is this a redox reaction the answer
is yes
so now we want to identify what is being
oxidized and what is being reduced we
said earlier that a loss of electrons is oxidation
oxidation
we can see that hydrogen is the species
that would be losing potentially its
electrons as it goes from a zero
oxidation number two plus one right so
it would have had to lose an electron in
order to increase its oxidation number
so we can see that H2 is the species
that is oxidized
um what is being reduced well we can see
that the chlorine's oxidation number is
reduced right it goes from 0 to -1 as it
potentially gains those electrons so
here we have chlorine cl2 is being reduced
reduced
now what is being what is the reducing
agent right so the reducing agent is the
species that is going to give up the
electrons for
um for the reduction process right so we
know that the hydrogen is what is losing
the electrons it's giving up its
electrons to the chlorine and so our
reducing agent
is the hydrogen H2 so the species being
oxidized is the reducing agent and then
this tells us that the oxidizing agent
is the chlorine gas right it is
receiving the electrons that the
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