Hi. Welcome back to General Chemistry. My
name is Chuck Wight, and today's lesson is
on Thermochemistry. We're going to start by
talking about energy, kinetic energy and potential
energy, which are traded off between one or
the other, and then we'll talk about chemical
energy and heat energy, and what those things
are. We'll talk a little bit about the First
Law of Thermodynamics, which is statement
about conservation of energy, and we'll also
talk about how to calculate the enthalpy for
a chemical reaction. Now, let's start by talking
about a bouncing ball, and we'll just assume
that we have a ball that weighs 0.1 kg. The
height of the bounce is 0.2 m, and the ball
has a volume of 0.0001 cubic meters, all in
SI units. Now at the top of its trajectory,
this ball has gravitational potential energy.
Its velocity, at least instantaneously, is
zero, and so we can calculate the amount of
gravitational energy it has at the top of
its trajectory by taking the mass times the
gravitational acceleration times the height,
and all that multiplied together is 0.196
J. That's the total amount of gravitational
potential energy the ball has at the very
top. At the bottom of its trajectory, it has
(just before it hits the ground) a lot of
kinetic energy of motion. By conservation
of energy, that has to be the same 0.196 J,
and that's equal to one-half m v-squared,
the kinetic energy, and so we can solve for
the velocity of the ball just before it hits
the ground, of 1.98 m/s. Now as the ball hits
the ground, it gets compressed, and the kinetic
energy gets converted to potential energy
of compression, and we can actually calculate
the amount of compression (it's about 2%),
and that compressive energy gets stored and
then released again, pushing the ball up back
to near its original height. Eventually, the
ball will cease to bounce and that's due to
friction and so forth, and all of that energy
that was traded off between gravitational
potential energy, kinetic energy and compressive
energy, will all get converted into heat.
Now let's talk a little bit about chemical
reactions. We're all familiar with the reaction
of hydrogen and oxygen, that can make water,
and that's a very exothermic reaction, so
it releases a lot of heat. Just in the same
way that we can trade off gravitational energy
and kinetic energy of the ball, we can trade
off chemical energy of the hydrogen and oxygen,
and as that gets turned into water, water
has a much lower chemical energy, so the difference
between those two things is released as heat
energy. Now the thing about chemical energy
is the kinetic and potential energy of the
electrons in the hydrogen and oxygen is higher
that that in the water. That is to say, the
chemical energy of the water is lower. So
that makes this particular reaction exothermic.
Now, what is heat energy, exactly?
Heat is when energy is distributed randomly
among all possible degrees of freedom of a
collection of atoms and molecules, that is
to say, translational energy of gas molecules,
for example, rotational energy, which is also
kinetic energy of gas molecules, and vibrational
energy of molecules, which is a combination
of kinetic and potential energy. When it's
all distributed randomly, that's known as
heat. So, how do we measure heat?
We. we know that we can tabulate the heat
capacity of things like water. Water has a
heat capacity of 4187 joules per kilogram
per Kelvin. So if we have a liter of water,
which weighs 1 kg, then it takes 4187 J of
energy to raise the temperature of that water
by one degree Kelvin. Typically, we measure
heat indirectly by measuring changes in temperature
and knowing the heat capacity of the thing
that we're heating. Now that brings us to
the First Law of Thermodynamics. The First
Law is an expression of conservation of energy,
and it says that the change in energy, Delta-E,
for any system, like a liter of water for
example, is equal to the sum of two terms:
q, which is the heat added to the system,
and w, which is the work done on the system.
So, we're all familiar with what heat is now,
but what is thermodynamic work?
Work is the non-heat energy added to, or taken
away from, the system. In chemistry, we're
mainly concerned with compression work, usually
on gases. So work is equal to the integral
of -PdV, which is an expression of when you
change the volume, as you decrease the volume
of a system by compressing it, then if you
integrate the pressure times the change in
volume, that's equal to the work that's done.
We're sometimes also interested in electrical
work, and that would be the integral of VdQ.
That would be the electrical potential, in
volts, times the difference in charge. So
now that brings us to calculating the enthalpy
changes for reactions, calculating the amount
of heat released or consumed in a chemical
reaction at constant pressure. Enthalpy is
very similar to energy, but it automatically
includes some terms associated with an expansion
and contraction of gases. The assumption here
is that we're working at constant pressure,
which is almost always the case in ordinary
chemical reactions. So now the enthalpy change
for a reaction, Delta-H of reaction, is equal
to the enthalpy of formation of the products
of the reaction minus the enthalpy of formation
of the reactants. So it helps to see an example.
Let's suppose that we have this reaction of
hydrogen and oxygen, which gives water. We
know that, by definition, for any element
in its standard state, the enthalpy of formation
is zero. So for oxygen, the enthalpy of formation
is zero and for hydrogen, the enthalpy of
formation is zero. But for water, we can look
up in tables that its enthalpy of formation
is -285.8 kJ per mole. So water has a much
lower enthalpy of formation, so we know that
as we transform hydrogen and oxygen to water
heat is going to be released in that chemical
reaction. How much will be released?
First of all we have to balance the chemical
reaction. The balanced reaction is 2 H2 plus
O2 gives 2 water molecules. So the Delta-H
of reaction will be the Delta-H of formation
of the products, in this case, water, 2 moles,
minus the Delta-H of formation of the reactants,
in this case 2 moles of hydrogen and one mole
of oxygen. So the Delta-H of reaction is twice
-285.8 kJ/mol for the water minus 2 times
zero for the hydrogen, minus zero for the
oxygen, and when you add all that up it turns
out to be -571.6 kJ/mol. A negative value
means the reaction is exothermic. If the reaction
turned out to have a positive Delta-H, then
that would be for an endothermic reaction.
Let's take another example for combustion
of methane. The balanced reaction for combustion
of methane is CH4 plus 2 O2 gives CO2 plus
2 H2O. The Delta-H, remember, is the Delta-H
of formation of the products minus the Delta-H
of the reactants, so if we write this out
for CO2, water, methane and oxygen, it turns
out that the Delta-H overall for this chemical
reaction is -890.5 kJ/mol of methane combusted.
So the combustion of methane is highly exothermic,
that's why it produces a flame and it's useful
for cooking. OK, next time we'll talk about
the quantum mechanics of atoms. We''ll talk
about wave-particle duality. We'll talk about
the deBroglie wavelength for quantum mechanical
sized particles. We'll talk about the Uncertainty
Principle that Heisenberg came up with, and
we'll talk a little bit about atomic orbitals.
We'll see you then.
