Video Transcript
In this video, we will learn about
different types of energy and see how they’re involved in the making and breaking of
chemical bonds and in chemical reactions.
First off, let’s have a look at
energy. Energy is a bit of a funny
thing. The word “energy” is like
currency. You can’t ever just have currency
in your pocket. It has to be currency of a certain
type, like dollars, pounds, or yen, or something else. Even though these are currencies,
there is no one currency, just like there isn’t one energy. But we do use the word “energy”
when we’re talking about multiple types of energy at once or when we’re not sure
what types we’re dealing with.
But all types of energy can be
measured in joules, which is the internationally accepted standard unit of
energy. It takes about 4.2 joules of energy
to raise the temperature of one milliliter, that’s one cubic centimeter, of water by
one degree Celsius. Unlike real currencies, energy
cannot be created or destroyed. However, energy in one form can be
converted to another form as long as the total amount doesn’t change. In currency terms, you can think of
this as the value never changing, even though you exchange your dollars for yen.
There are lots of types and groups
of energy, but there are too many for this video. So I’m only going to look at the
ones that are particularly useful for chemists.
There’s nuclear energy, which comes
from protons and neutrons binding together. And then there’s what we call
chemical energy, which is what we get from the attraction of electrons and protons
and all the things it leads to: ionic, covalent, and metallic bonding. At the root of all this is the
electrostatic attraction between positive and negative charges.
We could also put electrical energy
under the bracket of chemical energy, although it’s often separated out. And then there’s kinetic energy,
which is the energy an object has when it’s moving relative to another object. We can think about kinetic energy
of large objects just like we can think about the kinetic energy of individual
particles.
And then we have what we call
thermal energy, commonly called heat. This can involve kinetic energy,
rotations, or vibrations. The more thermal energy something
has, the hotter it is. Then we have energy in the form of
packets or photons of light. Then we have sound energy, which
involves a lot of kinetic energy. And then there’s gravitational
energy from the force of attraction between objects with mass. It’s the force of gravity between
an object and the Earth that allows us to measure the object’s mass. And then there’s the magnetic
energy that we see between poles of a magnet. In some systems, magnetic energy
does contribute to chemical energy, but we won’t be looking at the details of
that.
But I am going to use an analogy
for magnetism that will help you understand chemical energy. Here, we have two bar magnets. If we hold them in place, they will
attract each other and there will be magnetic potential energy. If we let go, this magnetic
potential energy first transforms into kinetic energy as the magnets accelerate and
move toward each other. When they collide, that built-up
kinetic energy is converted into thermal energy of the air and the magnets and sound
energy. And in the whole process, the total
energy hasn’t changed; only types of energy have been interconverted.
The exact same principle applies
with chemical reactions. When electrons and protons come
together, chemical potential energy due to the electrostatic attraction between them
is converted into kinetic energy as they accelerate and move toward each other. When they meet, all this kinetic
energy needs to be converted to another form of energy. Otherwise, they’ll just fly apart
again. Generally, they’ll collide with
something else and give up some of their energy. With collision after collision, the
energy is dispersed as chaotic kinetic energy, which is what we call thermal energy
or heat, although in some cases chemical reactions also produce other forms of
energy.
Now, let’s see how this applies to
making and breaking chemical bonds. Here, we can see two atoms forming
a single covalent bond. When chemical bonds form, chemical
energy, a type of potential energy, is converted into thermal energy, sound, light,
et cetera. Here, we can see the same bond
being broken. We can also do the reverse and
introduce energy to break bonds.
When we make bonds, the energy
released by converting chemical energy to other forms of energy has to be
transferred to the surroundings. But breaking bonds requires energy
coming from the surroundings into the system. Often, people talk about chemical
bonds storing energy, but that’s not correct. When chemical bonds form,
nonchemical forms of energy are released. And when chemical bonds are broken,
other forms of energy are converted into chemical potential energy. This is like lifting a book from
the floor to above your head. The energy you expend to lift the
book has been converted into gravitational potential energy. Potential energy is just primed to
be released by the formation of bonds or by dropping the book.
When you think about a chemical
system, it might be overwhelming to consider all the forms of energy in there:
chemical potential energy, thermal energy, excited electrons, light. But chemists have a way of lumping
them all together so it’s easier to talk about.
Quite often, it doesn’t matter so
much what form the energy is in. It just matters how much of it
there is. The word for the sum of all the
types of energy in a chemical system is the enthalpy, which we give the symbol
𝐻. We generally leave out certain
forms of energy, like gravitation and nuclear energy, for convenience. For the vast majority of chemical
systems, the chemical energy and the thermal energy make up the majority of the
enthalpy, while light and sound are relatively insignificant.
So if we keep it simple, we can
treat the enthalpy of the system as the chemical potential energy plus the thermal
energy. While there are other factors
involved, generally speaking, systems tend to a position where they have lower
enthalpy. So imagine a system that has more
enthalpy than the surroundings. We generally expect the enthalpies
to equalize by transfer of energy. We can think of this as adding or
subtracting enthalpy. And the enthalpy changes usually
involve thermal energy moving between the system and the surroundings. This is why enthalpy is often
described as the heat content of a system because even though a lot of the enthalpy
is from chemical potential energy, we see enthalpy moving as heat.
So all that’s been quite
theoretical up to now. Let’s have a look at an
example.
This is the decomposition of
calcium carbonate, producing calcium oxide and carbon dioxide gas. This reaction forms the backbone of
the construction industry because calcium oxide is a primary ingredient in
cement. And it’s produced from limestone,
which is mostly calcium carbonate.
If we looked at the enthalpy of the
system after heat has been transferred to and from the surroundings, we’d find
something interesting. The enthalpy of the products is
actually higher than the enthalpy of the reactants. We actually need to introduce 178
kilojoules of energy per mole of calcium carbonate we want to decompose. This is because the bonds in the
products are actually weaker than the bonds in the reactants. The reaction is partly driven by
the carbon dioxide gas, which simply escapes, stopping it reacting again with a
calcium oxide and forming the more stable calcium carbonate again.
But we could make this even more
simple. We don’t really care about the
exact enthalpy of the reactants and the products. All we really care about is the
difference in enthalpy or enthalpy change. This is given the symbol Δ𝐻, where
the Δ simply indicates change in what follows.
If the change in enthalpy is
positive, then the enthalpy of the system has increased. We call this type of reaction an
endothermic reaction because energy is being absorbed from the surroundings, which
generally reduces its temperature. We see exactly the opposite happen
when we do something like burn a fuel like methane or butane. In these reactions, the bonds in
the products are generally stronger than the bonds in the reactants. So the enthalpy of the products is
lower. This means that the change in
enthalpy of the system is negative. A negative change in enthalpy
indicates that the enthalpy of the system has decreased because of the reaction and
because energy has been lost to the surroundings. This is what we call an exothermic
reaction, because generally this type of reaction causes the temperature of the
surroundings to rise.
You may see endothermic reactions
written like this, where energy is actually a term in the equation. If it’s on the left, then we’re
inputting energy. It’s an endothermic process. But if we see energy on the
right-hand side, that means the formation of the products releases energy. And we’re dealing with a negative
enthalpy change and an exothermic reaction.
So, going back to our example, you
may see the enthalpy change written as energy in on the left-hand side for the
decomposition of calcium carbonate. This means that the enthalpy change
is positive 178 kilojoules per mole of calcium carbonate because the stoichiometric
coefficient next to calcium carbonate is one. And we can read each coefficient in
moles. And the whole equation becomes one
mole of calcium carbonate plus 178 kilojoules react to form one mole of calcium
oxide plus one mole of carbon dioxide.
The next thing we’re going to look
at are the conditions we use when we measure enthalpy changes. It’s really difficult to measure
enthalpy directly because there are so many factors involved. Instead, we measure enthalpy
changes generally by saying how much heat comes out of or goes into a specific
reaction. But enthalpy changes are not always
consistent. They can change depending on the
conditions. This is why scientists agreed on a
few set conditions so that enthalpy changes were all measured in the same way.
When these conditions have been
used, this character is added to indicate the measurement was made under standard
conditions. It’s usually added to the top right
of the relevant measurement, for instance, Δ𝐻 ⦵.
Standard conditions are one bar of
pressure, although all the textbooks may use one atmosphere. One atmosphere is equivalent to
1.01325 bar. When measuring under standard
conditions, it’s vitally important to state the temperature. But standard conditions are usually
taken to be 25 degrees Celsius or 298.15 kelvin. This is commonly rounded to just
298 kelvin. And when dealing with measurements
involving solutes, they should be at one-molar concentration.
These conditions mean that standard
enthalpy changes must be measured under constant pressure. If the pressure changes, other
energy changes need to be factored in. And a lot of measurements under
standard conditions assume that chemicals are in their standard states. But what does that mean?
Since enthalpies are hard to
measure directly, scientists came up with a special system, where every chemical has
a state that is its zero point for enthalpy changes. It’s more complicated than it
sounds. We basically chose a point to
measure against, and we chose it for each chemical and each element. Rather than measuring all the
possible enthalpies of every single chemical, we just have standard values that we
reference either being above or below them. We then add up the differences for
the whole reaction.
For most chemicals, the standard
state is simple. It’s the typical state of that
chemical at one bar of pressure and 25 degrees Celsius, although for chemicals that
aren’t stable under those conditions the reference state has to be chosen
individually. For example, the standard state for
water, H2O, is liquid because water is a liquid at one bar and 25 degrees
Celsius. For some elements, the choice of
standard state isn’t that simple because some elements have more than one form or
allotrope. For these elements, you just have
to remember what the standard state is.
Carbon has many allotropes,
including graphite, diamond, and fullerenes. And they all have slightly
different enthalpies. So for carbon, the standard state
is chosen as graphite. In equations, strictly speaking, we
should write this as carbon in a solid graphite form because carbon solid is not
specific enough. There are other elements with
allotropes, but carbon is the one that most commonly pops up. For most of the other elements, the
standard state is simply the only state we find it in normally under those
conditions, like hydrogen gas as H2, oxygen as O2, and sodium metal in its solid
form. So after all that, I think it’s
time for some practice.
Which of these statements does not
describe the conservation of energy in a chemical reaction? (A) Energy is neither created nor
destroyed during a chemical reaction. (B) If the energy of a system
increases, then the energy of the surroundings decreases by the exact same
amount. (C) Energy can only be transferred
from one form to another. (D) If the energy of a system
decreases, then the energy of the surroundings increases by the exact same
amount. Or (E) the energy contained in the
bonds of reactant molecules always equals the energy contained in the bonds of
product molecules.
The question’s asking us to
identify the one statement out of the five that does not describe the conservation
of energy in a chemical reaction. The law of conservation of energy
can be stated in the form that energy cannot be created or destroyed. But energy can be converted between
its various forms, like thermal energy, sound, light, and chemical potential
energy.
The first statement says that
energy is neither created nor destroyed during a chemical reaction. This is perfectly true. And in fact, it doesn’t even need
to apply to chemical reactions. It’s a universal law. So since this does describe the
conservation of energy in a chemical reaction, it’s not a correct answer.
The second statement says that if
the energy of a system increases, then the energy of the surroundings decreases by
the exact same amount. What this statement is suggesting
is if the system energy goes up, then it must be drawing energy from the
surroundings. So the energy of the surroundings
must go down. This is an application of the law
of conservation of energy. The total energy must be
constant. So this statement is true. The only things in the entire
universe are the system and the surroundings. And the total energy of the
universe must be constant. We can substitute the term energy
and enthalpy when it pertains to the system, and it’s still true.
The third statement says that
energy can only be transferred from one form to another, for instance, the
conversion of electrical energy into heat and light in a tungsten bulb. This squares with our conservation
law perfectly. Energy can only be transferred or
converted from one form to another.
Statement (D) says that if the
energy of a system decreases, then the energy of the surroundings increases by the
exact same amount. This is just statement (B) the
other way around. And we can see that it’s perfectly
true, which leaves us with (E) the energy contained in the bonds of reactant
molecules always equals the energy contained in the bonds of product molecules. We know this is false because we
often see chemical bond energy changing throughout a reaction. In our example, when we form
calcium oxide and carbon dioxide from calcium carbonate, the products have weaker
bonds than the reactants.
So on to the key points, energy
cannot be created or destroyed, only converted between types. The energy of a chemical system is
known as the enthalpy, with the symbol 𝐻. Forming bonds decreases the
enthalpy overall, and breaking bonds increases it. Enthalpy changes occur when
reactions happen. And enthalpy changes are commonly
measured under standard conditions, indicated with the ⦵ character: one bar of
pressure, 25 degrees Celsius, one-molar solutions, and standard states for
chemicals.
