Lesson Video: Bond Energy | Nagwa Lesson Video: Bond Energy | Nagwa

Lesson Video: Bond Energy Chemistry

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.

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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.

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