HS Chemistry - Heat in Reactions

Exothermic & Endothermic Reactions

Overview of The Page

This page will cover:

• What are exothermic and endothermic reactions?
• What is a chemical reaction's activation energy?

An exothermic reaction releases heat, while an endothermic reaction absorbs heat.

More specifically:

• In an exothermic reaction, heat energy is released. We say heat energy is evolved $as in, it evolves out of the system$. There is an overall decrease of energy in the system, shown by a negative enthalpy change value $a negative ∆H value$.

• In an endothermic reaction, heat energy is absorbed. There is an overall increase of energy in the system, shown by a positive enthalpy change value $a positive ∆H value$.

But what's meant by overall change of energy?

On the last subpage, the following reaction was given:

H + Cl → HCl

And in this reaction, the reactants have a greater energy than the products, because the atoms are separate in the reactants $and therefore more reactive, and have more energy$, while in the product, they've bonded into a molecule $therefore, they're less reactive and have less energy$.

This reaction sees an overall decrease in energy - it's an exothermic reaction.

So far, we know that atoms have less energy when bonded in a molecule than when they're separate. The product in the above reaction has less energy than the reactants. Since energy can't be created or destroyed, the excess energy is released, or evolved, as heat. Combining raw atoms to form a molecule is an exothermic reaction.

The opposite should hold true as well. Breaking a molecule into its separate atoms should require energy to be added in. In the following reaction:

HCl → H + Cl

1 molecule of Hydrochloric acid is broken up into 1 Hydrogen atom and 1 Chlorine atom. In this reaction, the reactant has less energy than the products, as the atoms have more energy when separate than when bonded together.

This reaction sees an overall increase in energy - it's an endothermic reaction. The additional energy that entered the system was absorbed from the surroundings.

A More Complex Reaction

It's rare to have atoms existing by themselves, except when undergoing a reaction $the exception tends to be monatomic molecules, although there are other exceptions as well$. A more common reaction, and one we might experimentally see, would be:

H₂ + Cl₂ → 2HCl

1 molecule of Hydrogen gas reacts with 1 molecule of Chlorine gas to form 2 molecules of Hydrochloric acid. The molecule of Hydrogen gas splits into two Hydrogen atoms, and the molecule of Chlorine gas splits into two Chlorine atoms. Then, each Hydrogen atom bonds with a Chlorine atom to form a molecule of Hydrochloric acid.

This reaction shows multiple changes inside the system:

1. H₂ → H + H
2. Cl₂ → Cl + Cl
3. 2H + 2Cl → 2HCl

The first two absorb energy, as the atoms in the molecule are separated. But the third one releases energy, as atoms are combined into molecules.

It's unlikely that the energy absorbed in the first two changes will be balanced out by the energy released in the third change. As such, there will be an overall change in the system's energy before the changes and after the changes. Either the system will have more energy before all the changes than after all those changes $in which case the overall reaction will be exothermic$, or the system will have more energy after the changes than before $in which case the overall reaction will be endothermic$.

This overall change in the system's energy is called the enthalpy change. Depending on the nature of the change, we can further specify what kind of enthalpy change is happening:

• If it's the enthalpy change of a reaction, it's called the enthalpy change of reaction $∆H~r~$.

• If it's the enthalpy change in forming a molecule, where one mole of a molecule is being formed from its elements in their most stable form $like 1/2 H~2~ \+ 1/2 F~2~ → HF$, it's called the enthalpy change of formation $∆H~f~$.

  • When we say "elements in their most stable form", that means the form that has the least energy. For example, a lone Oxygen atom $O$, oxygen gas $O~2~$ and ozone $O~3~$ all consist of only oxygen atoms, but O₂ is the most stable form of oxygen, as it has the least energy. So if we had to show the reaction of formation for water, it would be H₂ + 1/2 O₂ → H₂O, not H₂ + O → H₂O or H₂ + 1/3 O₃ → H₂O.

• If it's the enthalpy change in burning a substance or caused by burning a substance, it's called the enthalpy change of combustion $∆H~C~$.

• If it's the enthalpy change in a neutralization reaction $something that won't be covered in this Unit$, it's called the enthalpy change of neutralization $∆H~neut~$.

Let's go back to our reaction:

H₂ + C₂ → 2HCl

Is this an endothermic reaction or an exothermic reaction? We can find out by calculating the enthalpy change of the reaction by subtracting the enthalpy of the reactants from the enthalpy of the products.

• Enthalpy of reactants $H~2~ \+ Cl~2~$ = -436 kJ + -242 kJ = -678 kJ

• Enthalpy of products $2HCl$ = 2 × -431 kJ = -862 kJ

The enthalpies of both the reactants and the products are derived from the bond energies and heats of formation of the molecules, which were taken from a table. They are both negative because the molecules on both the reactant and product sides have less energy than if the atoms were separate. Bond energies and heats of formation are different for different molecules.

Subtracting the reactants from the products will give us -862 - $\-678$ = -184 kJ. Thus, the reaction H₂ + Cl₂ → 2HCl is exothermic, and it releases 184 kJ.

Activation Energy

The 2HCl have a lower combined energy than the H₂ and Cl₂. Atoms will take the most stable form they can. So why don't the H₂ and Cl₂ automatically turn into 2HCl?

In order for the H₂ and Cl₂ to form 2HCl, they first need to be broken up into 2 Hydrogen atoms and 2 Chlorine atoms. Breaking up a molecule into atoms requires energy. Therefore, we need some initial energy source $provided as heat$ in order to break up the H₂ and Cl₂. This is called the activation energy. If this activation energy isn't provided, then the reaction won't occur, even if the final products have a lower energy than the initial products. In that case, the mixture is kinetically stable, but energetically unstable.

In an exothermic reaction, the heat energy produced when the atoms form the products after being separated is more than enough to provide the activation energy for the next set of reactant molecules. In an endothermic reaction, however, it isn't, which is why endothermic reactions need a continuous supply of heat in order to continue occurring $until the supply of one of the reactants is exhausted$.

The potential energy diagram for the reaction is going to be:

Where:

• E_a is the activation energy.

• The energy released when the products are formed is shown by the large drop, signaling the decrease in the system's enthalpy from the activated complex.

• The overall enthalpy change $∆H~r~$ is the difference between the energy level of the reactants and the energy level of the products. It's calculated by subtracting the energy level of the reactants from the energy level of the products. In this example, it is equal to -184 kJ $184 kJ of energy is released, or evolved$.

There's some space left between the graph and the bottom axis; that's because we don't want to imply that the energy in the system has hit zero. There's still some energy in the system, we just want to show how it has changed. Therefore, the diagram isn't drawn to scale.

There's two other things to note:

1. The arrows showing the energy that's released when the products are formed are just there to help further explain the diagram. It's not actually a part of it.
2. When drawing a potential energy diagram, we don't need to include most of the information shown in the one above. We can just show:

But it's nice to show the extra information.

The potential energy diagram also has a peak, labeled the activated complex, which is when the H₂ and Cl₂ have broken up into 2H and 2Cl. At this moment, the atoms are separate, and the system has more energy than it did at either the beginning or the end of the reaction.

Once the activated complex is reached, the separated atoms bond to form the products, and release energy. The energy of the system $its enthalpy$ drops back down as the separated atoms, wanting to be stable, bond in ways that achieve the lowest energy level.

But doesn't that mean that all reactions should be exothermic? How can the separated atoms in the activated complex form molecules that have a higher energy level than the original reactants? After all, in that case, the reactants have a lower energy level, so according to what's been covered in this subpage so far, they should just form the reactants again.

Endothermic reactions, which see an overall increase in the system's energy (and which have a positive enthalpy change) happen because of something known as entropy, but that won't be covered here.

In summary, for the reaction H₂ + Cl₂ → 2HCl:

• There's only one reaction here: a molecule of hydrogen gas reacts with a molecule of chlorine gas to form 2 molecules of hydrochloric acid.

• However, there are three things happening here:

  1. The molecule of Hydrogen gas splits into two Hydrogen atoms.

  2. The molecule of Chlorine gas splits into two Chlorine atoms.

  3. Each of the two Hydrogen atoms bonds with a Chlorine atom, forming a molecule of Hydrochloric acid. Two molecules of Hydrochloric acid are formed.

• Each of the changes to the system has an enthalpy change. Combining them all gives the total enthalpy change of the reaction, which tells us if the reaction is exothermic $if the total enthalpy change is positive$ or if it is endothermic $if the total enthalpy change is negative$

• Even though this reaction $the one given above as an example, H~2~ + Cl~2~ → 2HCl$ is an exothermic one, an initial input of energy, known as the activation energy, is required to initiate it.