Endothermic vs. exothermic reactions (article) | Khan Academy (2024)

Excited but a bit confused, Sam and Julie run to their chemistry teacher. Sam asks, “Teacher, why did my flask turn cold after adding the salt to water, while Julie’s flask turned hot?”

The teacher replies: “That’s because you were given two different salts. One of your salts generated an endothermic reaction with water, while the other salt generated an exothermic reaction with water. Let me first reveal the identity of your salts: Salt A is ammonium nitrate (${\text{NH}}_4{\text{NO}}_3$‍) and Salt B is calcium chloride ($\text{Ca}{\text{Cl}}_2$‍)."

Now, Sam and Julie are curious about the difference between an endothermic and an exothermic reaction.

Consider the reaction mixture—salt plus water—as the system and the flask as the surrounding.

In Sam’s case, when ammonium nitrate was dissolved in water, the system absorbed heat from the surrounding, the flask, and thus the flask felt cold. This is an example of an endothermic reaction. In Julie’s case, when calcium chloride was dissolved in water, the system released heat into the surroundings, the flask, and thus the flask felt hot. This is an example of an exothermic reaction.

The reaction going on in Sam’s flask can be represented as:

You can see, heat is absorbed during the above reaction, lowering the temperature of the reaction mixture, and thus the reaction flask feels cold.

The reaction going on in Julie’s flask can be represented as:

In this case, heat is released during the reaction, elevating the temperature of the reaction mixture, and thus Julie’s reaction flask feels hot.

The teacher’s final comment to Sam and Julie about this experiment is, “When trying to classify a reaction as exothermic or endothermic, watch how the temperature of the surrounding—in this case, the flask—changes. An exothermic process releases heat, causing the temperature of the immediate surroundings to rise. An endothermic process absorbs heat and cools the surroundings.”

Based on the above definition, let's pick a few examples from our daily lives and categorize them as endothermic or exothermic.

1) Photosynthesis: Plants absorb heat energy from sunlight to convert carbon dioxide and water into glucose and oxygen.

2) Cooking an egg: Heat energy is absorbed from the pan to cook the egg.

1) Combustion: The burning of carbon-containing compounds uses oxygen, from air, and produces carbon dioxide, water, and lots of heat. For example, combustion of methane (${\text{CH}}_4$‍) can be represented as follows:

2) Rain: Condensation of water vapor into rain releasing energy in the form of heat is an example of an exothermic process.

In any chemical reaction, chemical bonds are either broken or formed. And the rule of thumb is "When chemical bonds are formed, heat is released, and when chemical bonds are broken, heat is absorbed." Molecules inherently want to stay together, so formation of chemical bonds between molecules requires less energy as compared to breaking bonds between molecules, which requires more energy and results in heat being absorbed from the surroundings.

Enthalpy of a reaction is defined as the heat energy change ($\Delta H$‍) that takes place when reactants go to products. If heat is absorbed during the reaction, $\Delta H$‍ is positive; if heat is released, then $\Delta H$‍ is negative.

$\text{ΔH value negative --> energy released --> exothermic reaction}$‍$\text{ΔH value positive --> energy absorbed --> endothermic reaction}$‍

$∆H=\sum ∆H\text{(bonds broken in reactants)}-\sum ∆H\text{(bonds made in products)}$‍

Let's understand this through an example. We can calculate the enthalpy change ($\Delta H$‍) for the following reaction:

We know that the bond energy—in kilojoules or kJ—for $H_2$‍, $F_2$‍, and $HF$‍ are $436$‍, $158$‍ and $568$‍ kJ/mole respectively.

Let’s first figure out what’s happening in this particular reaction. Looking at the chemical reaction, it’s clear that one mole of $H-H$‍ and one mole of $F-F$‍ bonds are being broken to generate two moles of $H-F$‍ bonds. Breaking of bonds requires absorption of energy, while formation of bonds releases energy.

To break one mole of $H_2$‍, energy absorbed is $436$‍ kJ.

To break one mole of $F_2$‍, energy absorbed is 158 kJ.

To form two moles of $HF$‍, energy released is 2 X (568) kJ.

So applying the equation, $∆H=\sum ∆H_{(bondsbrokeninreactants)}-\sum ∆H_{(bondsmadeinproducts)}$‍

The overall enthalpy of the reaction is negative, i.e., it’s an exothermic reaction where energy is released in the form of heat.

In a chemical reaction, some bonds are broken and some bonds are formed. During the course of the reaction, there exists an intermediate stage, where chemical bonds are partially broken and partially formed. This intermediate exists at a higher energy level than the starting reactants; it is very unstable and is referred to as the transition state. The energy required to reach this transition state is called activation energy. We can define activation energy as the minimum amount of energy required to initiate a reaction, and it is denoted by $E_{act}$‍.

An energy diagram can be defined as a diagram showing the relative potential energies of reactants, transition states, and products as a reaction progresses with time. One can calculate the $E_{act}$‍ and $\Delta H$‍ for any reaction from its energy diagram.

Let’s draw an energy diagram for the following reaction:

The activation energy is the difference in the energy between the transition state and the reactants. It’s depicted with a red arrow. The enthalpy change—$\Delta H$‍—of the reaction is depicted with a green arrow. So, now you should be able to clearly differentiate between $E_{act}$‍ and $\Delta H$‍ on an energy diagram.

In the case of an endothermic reaction, the reactants are at a lower energy level compared to the products—as shown in the energy diagram below. In other words, the products are less stable than the reactants. Since we are forcing the reaction in the forward direction towards more unstable entities, overall $\Delta H$‍ for the reaction is positive, i.e., energy is absorbed from the surroundings.

In the case of an exothermic reaction, the reactants are at a higher energy level as compared to the products, as shown below in the energy diagram. In other words, the products are more stable than the reactants. Overall $\Delta H$‍ for the reaction is negative, i.e., energy is released in the form of heat.