Video Transcript
In this video, we will learn what
bond energy is, look at energy changes when reactants break apart and products form,
and use bond energies and reaction energies to calculate unknowns.
So firstly, a quick recap. What’s a chemical bond? A chemical bond is a strong, stable
attraction between two atomic-scale things. But what strong and stable actually
mean is a little difficult to define. But we generally include the
electron–proton interactions that occur when electron is around a nucleus, a
covalent bond between two atoms, the attraction between oppositely charged ions, and
the attraction between ions and delocalized electrons. Forces of attraction can be weak
like the London dispersion forces between molecules. We generally wouldn’t consider
these when we’re talking about chemical bonds.
We generally take the attraction
between electrons and nuclei for granted, leaving us with covalent bonds, ionic
bonds, and metallic bonds. All of these involve some kind of
attraction between electrons and the nuclei of atoms or ions. A chemical bond isn’t some kind of
physical rod that joins things together. It’s just a name we use when we see
that we need energy to separate things. You can think of bond energy like
the energy needed to lift a book above your head. The book is perfectly stable on the
floor where it is. But if you put effort into lifting
it, it will separate from the floor and stay above your head. We can call this book-lift
energy.
But what happens if you let go? It will naturally fall back down
again. This is because there’s a force of
attraction between the book and the earth. As the book falls, it will heat up
the air, and when it hits the floor, the floor will heat up and there’ll be a
thump. The total amount of energy released
will be the same as the book-lift energy. It’s similar for a chemical
bond. Here we have two ions next to each
other. One is positively charged, and one
is negatively charged. Since they’re oppositely charged,
there is a force of attraction between them.
If they are already stuck together,
you need energy to separate them. The energy could be thermal
energy. It could be light energy. It doesn’t really matter. What we care about is the
amount. So, we can put this energy in,
separating our two lines and converting this energy to chemical potential
energy. So, bond energy is a general term
for the amount of energy needed to break the bond, to separate the things that are
bonded. But we can also reverse the
situation and let our ions snap back together. This formation of a bond releases
energy of other types, like thermal, light, or sound. So, the bond energy is also equal
to the energy release when the bond is formed.
Remember that a bond energy is
always a positive amount of energy. We only ever see negative energies
from bonds when we’re considering which direction the energy is going. But we’ll come to that later. Forming one chemical bond releases
a minuscule amount of energy. If two nitrogen atoms come together
to form a nitrogen-nitrogen triple bond, only 0.00000000000000000156 joules of
energy is released to the surroundings. So it’s much easier to talk about
bond energies, otherwise known as bond enthalpies, per mole of bonds. If we take a mole of nitrogen gas
into molecules and break all their bonds, it will take about 942000 joules to do the
job. What we get in the end is two moles
of nitrogen atoms. Remember that one mole of atoms is
equivalent to an Avogadro’s number of atoms, which is about 6.022 times 10 to the
23.
So, going back to our bond energy,
we have a bond energy for the nitrogen-nitrogen triple bond of 942000 joules per
mole. But this is a little awkward to
say. So, instead of using joules per
mole, we use kilojoules per mole. This unit is ideal for the vast
majority of bonds whose energies lie in the range between 50 and 1000 kilojoules per
mole. A phrase you’ll hear a lot when
talking about bonds and bond energies is bond strength. It’s pretty simple; the greater the
energy input required to break a chemical bond, the stronger it is. So, the higher the bond energy, the
stronger the bond. Let’s look at a few different bonds
and compare their strength. From this point forward, we’ll only
be using the traditional definition of bond energy where we’re looking at the energy
required to separate the components.
Hydrogen atoms tend to form
hydrogen molecules containing a hydrogen-hydrogen single covalent bond. The bond energy for this bond is
432 kilojoules per mole. This means that in order to
separate one mole of hydrogen molecules into two moles of hydrogen atoms, we’ll need
to put in 432 kilojoules of energy. An oxygen molecule consists of two
oxygen atoms bonded by a double covalent bond. This bond is about 15 percent
stronger than the hydrogen-hydrogen bond. And then we have the nitrogen
molecule consisting of two nitrogen atoms bonded by a triple covalent bond. This bond is over twice as strong
as a hydrogen-hydrogen single bond.
You might have noticed a pattern;
the nitrogen-nitrogen triple bond is stronger than the oxygen-oxygen double bond,
which is stronger than the hydrogen-hydrogen single bond. You might think this pattern is
always true, that all triple bonds are stronger than all double bonds and all double
bonds are stronger than all single bonds. This isn’t always true, but it is
true in very many cases. For instance, one exception is that
a hydrogen-hydrogen single bond is actually stronger than a nitrogen-nitrogen double
bond. But if we don’t change the atoms,
we just change the number of bonds between them, it will always be true, of course
assuming the bond can form in the first place.
Let’s have a look at the single,
double, and triple bonds between carbon atoms. Just bear in mind we expect carbon
to have four bonds, so these bonds would be part of larger molecules. The typical bond energy for a
carbon-carbon single bond is 346 kilojoules per mole. The molar bond energy of the
carbon-carbon double bond is about double that of a single bond. And the bond energy of the triple
bond is about 2.5 times that of the single bond. So, if we’re dealing with the same
atoms, a single bond is weaker than the double bond, which is weaker than a triple
bond.
How does all this apply to
reactions? Generally, when reactants transform
into products, bonds have to be broken in the reactants, and new bonds are formed
when the products form. For instance, when we react
hydrogen and chlorine gas together, we break hydrogen-hydrogen single bonds and
chlorine-chlorine single bonds. Generally, we need to put energy in
to break these bonds, so we add energy to the system. Our product is hydrogen chloride
containing a hydrogen-chlorine single bond. And we get two molecules of this
for one molecule of hydrogen and one molecule of chlorine. Energy is released when the bonds
form in the products. So we get energy out of the
system.
When keeping track of energy going
in and coming out, we can use signs, plus and minus. Energy going in increases the
energy of the system, so the change in energy is positive. Energy being released from the
formation of bonds means energy is leaving the system, so the change in energy is
negative. Taken in isolation, we can describe
the breaking of bonds as an endothermic process because energy is coming from the
surroundings into the system, while the formation of bonds involves the release of
energy into the surroundings. So it’s an exothermic process. Whether the reaction overall is
endothermic or exothermic comes down to the balance of the two.
If we know how strong the bonds in
the reactants and products are, we can predict if the reaction will be exothermic or
endothermic. If the bonds in the reactants are
actually stronger than the bonds in the products, we’ll actually get less energy out
than we put in. The reaction can still happen, but
energy will be required from the surroundings, so the surroundings will generally
cool down when the reaction occurs. If the reverse is true and the
bonds in the products are stronger than the bonds in the reactants, we’ll get more
energy out than we put in and we’ll be dealing with an exothermic reaction. So let’s dig into the details of
our example where hydrogen and chlorine react to produce hydrogen chloride.
The bond energy of the
hydrogen-hydrogen bond is 432 kilojoules per mole. The chlorine-chlorine single bond
is quite a bit weaker, at only 240 kilojoules per mole. And the bond energy for HCl is 428
kilojoules per mole. The total energy required per mole
of reactants is 672 kilojoules per mole. And the total per mole of products
is 856 which is twice that of the molar bond energy for hydrogen chloride. The energy released when the bonds
in the products form is greater than the energy required to break the bonds in the
reactants. So we’re dealing with an exothermic
reaction.
We can calculate the total change
in energy or enthalpy by taking the total bond energy of the reactants, which is
energy we need to put in, and taking away the energy we get out from the formation
of bonds in the products. In this case, the difference in
bond energy between the reactants and the products is 184 kilojoules per mole. And we’re dealing with exothermic
reaction, so this is energy leaving the system. But let’s take away the term energy
and talk about enthalpy for a moment.
We can think of enthalpy as the
energy of the system we’re dealing with. It’s sometimes easier than talking
about just energy because we always know that we’re talking about a direction. If energy is leaving the system,
then the change in enthalpy is negative. And if energy is entering the
system, then the change in enthalpy is positive. So, we can simply substitute in
enthalpy where we had energy. We have bond enthalpies, and the
change in enthalpy of the reaction is equal to the total bond enthalpy of the
reactants minus the total bond enthalpy of the products. With this example, we’ve already
demonstrated how we would calculate an unknown reaction enthalpy using known bond
enthalpies.
We calculated the total bond
enthalpy of the reactants and the products and subtracted the energy released when
bonds form from the energy required when bonds are broken. The change in enthalpy of the
reaction is negative 184 kilojoules per mole. We’re dealing with an exothermic
reaction. But what if we don’t know one of
the bond enthalpies? Let’s say we don’t know what the
bond enthalpy is for the hydrogen-chlorine single bond. Let’s say that all we know is the
change in enthalpy of the reaction, the bond enthalpy of the hydrogen-hydrogen
single bond, and the bond enthalpy of the chlorine-chlorine single bond. While we’re working out the bond
enthalpy of the hydrogen-chlorine single bond, I’ll just call it 𝑥.
So I’ve substituted the change in
enthalpy due to reaction into our equation. Here, I’m omitting the units to
keep things a little bit cleaner. The total bond energy of the
reactants is one lot of the hydrogen-hydrogen bond enthalpy and one lot of the
chlorine-chlorine bond enthalpy. And the total bond enthalpy for our
products is twice the bond enthalpy for the hydrogen-chlorine single bond. Let’s rearrange by adding 184 to
both sides. And then adding two 𝑥 to both
sides gives us the 𝑥 on the left-hand side. We then get two 𝑥 is equal to 672
plus 184 which is equal to 856. And if we divide by two, we get the
bond enthalpy of the hydrogen-chlorine single bond, 428 kilojoules per mole.
When doing calculations like this,
it’s vital you keep track of the stoichiometric coefficients because you don’t want
to have half or double, for instance, the value of the bond enthalpy you’re actually
after. If it helps, you can write out each
bond individually and then add together the unknowns. Next, let’s put all this into
practice.
The diatomic molecule of
phosphorus, P2, contains a triple bond. The molecule is highly unstable and
rapidly converts to molecules containing only single bonds, such as the
pyramid-shaped molecule P4. The equation for this reaction is
given. This reaction is highly
exothermic. How many single bonds are present
in the P4 molecule?
We’re introduced to the diatomic
molecule, that’s a molecule containing two atoms, of phosphorus, meaning both atoms
are phosphorus atoms. In this molecule, there’s a triple
bond, meaning six shared electrons between the two phosphorus atoms. We’ve been given an equation where
we see two of these P2 molecules reacting to form one P4 molecule. And we’ve been told this reaction
is highly exothermic, which means much more energy is released because of the
reaction than is absorbed.
The first question is fairly
simple. We just need to count the number of
single bonds in a P4 molecule. Counting from the front, each
phosphorus atom has three bonds to each other phosphorus atom. And being careful not to count the
same bond twice, we can see overall there are six unique single bonds. For the next part, I’m going to
summarize some of the information because it’s not all essential to the answer. So we’re dealing with the reaction
of two P2 molecules reacting to form one P4 molecule. That’s a very exothermic
reaction.
Why is the conversion of P2 to P4
exothermic?
A full understanding and
explanation of why this reaction is exothermic is far beyond the scope of this
video. But we can use some simple
principles to understand where to look. When discussing energy and
reactions, we need to remember that bond breaking requires energy and bond formation
releases energy. An exothermic reaction releases
more energy than it requires. So in this case, we definitely know
that the total bond energy of our products, the P4 molecule, is greater than the
total bond energy of our two P2 molecules. So our answer to why is the
conversion of P2 to P4 exothermic is that two phosphorus-phosphorus triple bonds are
weaker than six phosphorus-phosphorus single bonds.
And we can simplify a little and
say that one phosphorus-phosphorus triple bond is weaker than three
phosphorus-phosphorus single bonds. I’ll stall our information away for
the next bit.
Which bar chart illustrates the
difference in bond energy between the single and triple bonds of phosphorus?
Let’s have a quick recap of what we
know. For covalent bonds between two
atoms, a triple bond is stronger than the double, and it’s stronger than the
single. Each graph has a bond energy for
the single bond and the triple bond. The higher the bar, the stronger
the bond. Bar chart (C) and (E) both have
higher bars for the PP single bond, so they can’t be true. And when we look at bar chart (B),
we can see the other piece of information we know from the previous parts isn’t
true. The PP triple bond is much greater
than three times the bond strength of the PP single bond. So this is not correct either.
This just leaves us (A) and (D),
where we see a much more substantial difference between the bond energy of the PP
triple bond and the PP single bond in bar chart (A). The difference in bar chart (D) is
too small. It’s only (A) that shows roughly
the relationship we’d expect between a single and triple bond strength.
So let’s finish off with the key
points. Bond energy is the energy required
to break a bond, and it’s usually given in kilojoules per mole. And the stronger the bond, the
higher the bond energy. Triple covalent bonds are stronger
than the equivalent double bonds, which are stronger than the equivalent single
bonds. And we can calculate the reaction
enthalpy by taking the total bond energy of the reactants and taking away the total
bond energy of the products. Bond breaking is endothermic, and
bond making is exothermic.
