- [Voiceover] Today we are going to be talking
about the Reaction Quotient, Q.
In this video, I'm going to go over,
how you calculate Q and how you use it.
We're gonna start with and example reaction
between sulfur dioxide,
S02 gas,
which will react with oxygen gas,
and this is a reversible reaction
that makes
sulfur trioxide
or SO3.
We should make sure this is a balanced reaction.
We have two sulfur dioxdes reacting with one O2
to give
2S03.
At equilibrium we can calculate
the equilibrium constant, Kc.
So, at equilibrium
we know the concentrations
should be constant because the rate of the
forward and backward reactions are the same.
And if we plug those concentrations in
to this expression, we will get Kc.
So,
Kc is
the product concentration raised
to the second power so that's from
this stoichiometric coefficient.
And then our reactant concentrations,
so S02
squared
and the concentration of O2.
So we know at some temperature,
if you plug in the equilibrium concentrations,
Kc is equal to 4.3.
But what if we're interested in looking
at the reaction and it's not
at equilibrium yet or maybe we just don't know
if it's at equilibrium.
In that case,
when you're not sure it's
at equilibrium or really at any point
in your reaction or any time,
we can calculate the Reaction Quotient, Q.
So Qc
is equal to
the concentration of our product
squared,
so the concentration of the product
raised to
the stoichiometric coefficient
times the reactant concentrations,
also raised to their stoichiometric coefficients.
So S02 squared
and O2.
So you may be wondering at this point,
what's the difference?
The equation for Qc and Kc will
always look exactly the same
and the main difference is when you use them.
The equilibrium constant K you calculate
only with the equilibrium concentrations.
So the c means everything
is in terms of the molar concentration.
And for the reaction quotient, Q,
again everything is in terms of molar
concentration but we can calculate it
with any concentrations and we don't
have to be at equilibrium.
Molar concentration...
So let's calculate this
for a set of example concentrations.
At some point in our reaction we have
the following concentrations.
We have 0.10
molar
S02,
0.30
molar
O2,
and 3.5
molar of our product.
So if we plug these numbers into
our expression for Qc,
we get
3.5 molar
squared
in the numerator
and 0.10 squared
times 0.3
in the denominator.
So if I plug this into my calculator,
I get that Qc
with this set of concentrations
is 4,083.
So now we know how to calculate Qc.
So next we're going to talk about
what it tells you.
So there are three possible scenarios.
So when Q
is equal to K,
that tells us we're at equilibrium.
So, if at any point you're not sure
if your concentrations are the
equilibrium concentrations,
you can calculate Q
and check if it's equal to K.
And in this case, it's not.
So the other two possibilities are
that Q is greater than K,
which is the case here.
Or Q can be less than K.
So let's go through both of those possibilities.
We can draw all of the possible values of Q
on a number line
or a Q line.
So Q can have values anywhere
from zero
to infinity.
When you have no product
your numerator is zero
and Q is equal to zero.
So that tells us
Q equals zero
when you have all reactants
and no products.
And then if you have no reactants left
and all products,
we have zero in the denominator
and that gives us a Q value of infinity.
So that means at Q equals infinity,
we have all products.
Then we have a bunch of values in between.
And I'm gonna just write some intermediate values
in here
but the actual intermediate values here
aren't super important.
We're mostly gonna wanna compare the relative
values of our Q and K.
So Q here is equal to 4,083,
which I will place
right around here.
So that's Qc
and our K
in yellow
is 4.3.
So we'll place that right around here.
So we can see that
our Q is larger than K
and it's closer to having all products.
At the concentrations we have up here,
we have way more products than we should
at equilibrium.
So our reaction is gonna try to adjust
the concentrations to get to equilibrium.
And what that means in terms of our number line
is that our concentrations are gonna shift
so that Q can get closer to K.
Since our shift is to the right,
and it's moving towards all reactants,
our reaction is going to favor reactants
to get to equilibrium.
So when Q is greater than K, like here,
we're going to favor
reactants...
reactants.
And then the last scenario when Q is less than K
our reaction will
favor products.
And we can show that also on our number line.
If we had a different set of concentrations,
where Q was less than K,
which I will show
using this color here.
If we had, say, a Q value around here
then our shift would be
to the right
towards making more products
and therefore that would mean our reaction
is gonna try to reach equilibrium
by favoring the forward reaction.
So that's how you calculate Q and how you use it
to see how the reaction concentrations
will shift to get to equilibrium.
In our next video, we'll go over
an example problem using Q and trying
to figure out how the reactant concentrations
will shift for another reaction.