Energy Profile

The Energy Profile of a Reaction

A chemical reaction involves a transfer of energy as reactants are converted into products. We illustrate this idea using an energy profile diagram.

A system always wants to be as stable and low energy as possible, which is why high potential energy species are very reactive and unstable. Another way to think of it is that the potential energy is how much energy is required to hold a molecule together. As we move along our reaction road, there is always a "hill" we have to clear, called the activation energy.

Even when the reactants are higher in energy, there is still a hill to clear. That is because it requires energy to break chemical bonds. The creation of bonds is what releases energy. The shape of an energy diagram will always be a hill for this reason, but the "steepness" of the hill and the length of the "downhill" portion is dependent on the energy of the reactants and products. In this article, we will be looking at the different kinds of energy diagrams and learning how to interpret them.

Energy Profile Diagram for Exothermic and Endothermic Reaction

An energy profile shows how the energy of a system changes as the reaction progresses. It is synonymous with an energy diagram/energy profile diagram.

As stated previously, the energy profile diagram follows the change in potential energy. The y-axis is the potential energy and the x-axis is the reaction coordinate/reaction progress. This represents the progress of the reaction from reactants to products. The reactants are labeled on the left and the products are labeled on the right. To get a clearer picture, we will be walking through the different types of energy diagrams.

By looking at an energy diagram, we can determine whether a reaction is exothermic or endothermic.

The diagrams for these types of reactions are shown below:

In an exothermic reaction (left) the reaction has a smaller activation energy since the system is going to a more stable state. In an endothermic reaction (right) more energy is required since the system is going to a less stable state. StudySmarter Original

Let's break down the different pieces here. The first is the activation energy (E_A). The "hill" is much higher for an endothermic reaction since this reaction is thermodynamically unfavored (i.e. the system is becoming less stable), whereas it is much lower for the thermodynamically favored exothermic reaction. The activation energy is measured from the energy level of the reactants to the "peak" of the curve.

The second piece is the difference in potential energy (ΔE). If ΔE>0, the reaction is endothermic since the potential energy of the system is increasing, while a negative ΔE is exothermic for the opposite reason. We measure this change where E_initial is the energy of the reactants and E_final is the energy of the products.

The difference in energy between the reactions can also be shown as a change in enthalpy (ΔH). Enthalpy is the part of the potential energy that can be converted into heat energy. The signs for the change in enthalpy is the same as for the change in potential energy. This also makes sense, since an exothermic reaction (literally meaning "outside heat") is releasing heat, so its heat energy is decreasing (ΔH < 0). However, an endothermic reaction ("inside heat") is gaining heat, so its heat energy is increasing (ΔH > 0). The last important section is the transition state.

$$[XY-Z]$$

Then,

$$XY + Z \rightarrow [XY-Z] \rightarrow XZ + Y$$

The reaction is climbing up that hill so that the reactants can react to produce products. This "peak" is when they have finally gotten enough energy to cause a reaction. The energy then decreases because this is the point where the bonds are being formed, which releases energy. The energy also decreases because the transition state breaks apart into products.

Catalyst energy profile

Another type of energy profile is the one for reactions with a catalyst.

A catalyst speeds up a reaction by lowering the activation energy necessary for the reaction to occur. It does this by providing the reaction with an alternate reaction pathway (i.e. the reactants react with the catalyst to get to the same products). The energy of the reactants and products doesn't change, it is the pathway energy that does.

Another example of a catalyst is an enzyme. An enzyme is a biological catalyst that works slightly differently from other catalysts. Enzymes will bind to the reactant (called a substrate) which forms an enzyme-substrate complex, this acts as our transition state. After binding, the enzyme will then release a new product and the enzyme will go on unchanged. An enzyme is like a mold that our substrate (think of a lump of clay) is placed in, so when it is removed from the mold, it has become something new. Like with other catalysts, enzymes lower the activation energy by providing another reaction pathway.

Energy Profile Diagram for Decomposition of Hydrogen Peroxide

To get a better idea of a catalyzed reaction, let's look at the decomposition of hydrogen peroxide.

Thecatalysts for the reaction create an alternate pathway so that the activation energy is lower. StudySmarter Original

The two things you'll notice about the catalyzed reaction diagram are that there are two humps instead of one, and there are 2 extra species present (Br and H ions). The two humps mean that there is an extra step in the reaction. Catalysts can lower the activation energy by providing an alternate pathway. Think of it like taking a detour to get to your destination faster. There are two catalysts present in this reaction mechanism, and they are the Br and H ions. These ions react with the initial reactant (hydrogen peroxide) to end up with the same products as without the catalysts. Like with taking a detour, the place you are leaving, and your destination doesn't change, but your route does.

Energy Profile Diagram for Multistep Reactions

The last type of energy profile we will cover is for multistep reactions. These are reactions that, like the same suggests, have several steps that proceed in order. Think of it like knocking down dominoes: As one falls, the other is knocked over. The catalyst reaction we saw before is an example of a multistep reaction.

Let's look at another multistep reaction mechanism:

$$NO + NO \rightarrow N_2O_2$$ $$N_2O_2 + H_2 \rightarrow H_2O + N_2O$$ $$N_2O + H_2 \rightarrow N_2 + H_2O$$ $$\text{Net reaction:}\,2NO + 2H_2 \rightarrow N_2 + 2H_2O$$

Species like N₂O₂ and N₂O are called intermediates. These are species that are both formed and consumed during the reaction mechanism, so they aren't shown in the net reaction. These are different from transition states, since transition states are the in-between period during a single reaction, while intermediates are actual products for that respective intermediate step. Let's look at the energy diagram for this mechanism:

Each step has its own activation energy, the largest activation energy is the one that dictates the activation energy for the whole reaction. StudySmarter Original

Each step has its own activation energy and transition state. The dips in the curve show where a product has been formed and is labeled. The overall speed of the reaction is dependent on the step with the largest activation energy (here it is step 2). This step is called the rate-determining step. Picture the reaction like as the flow of traffic: it can only go as fast as the slowest car in front, even if the cars before it can go faster.