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Lesson Video: Energy Changes in Reactions Chemistry • 10th Grade

In this video, we will learn about different types of energy and see how they are involved in the making and breaking of chemical bonds, and in chemical reactions.


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.

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