Have you ever wondered how ice cubes melt? How does water change into vapors? How does the bread get baked?
All these things happen when an endothermic chemical reaction takes place. Now we are wondering about that, what are endothermic reactions and how it takes place in chemical reactions.
What are Endothermic Reactions
The chemical reactions in which the reaction absorbs energy from the surrounding (usually in the form of heat) for breaking the reactant bonds and converting them into new products are called endothermic reactions.
In the endothermic process, energy is required to break the bond present among atoms of reactants. This energy is the activation energy which handles initiating a chemical reaction.
Once this energy adds to the system, it breaks the reactants' bonds and forms new bonds in the products.
In the end, the more energy released and makes the reaction endothermic. Energy is required to break the bonds of reactants.
If the energy is higher than the energy needed to make the bonds of the product, it will be exothermic. This will cause more energy to release.
The endothermic reaction is the reverse of the exothermic reaction. It occurs when the energy required to break the reactant bonds is lower than the energy needed to form product bonds.
In this way, the excess energy would be utilized in bond formation. This will happen at the R.H.S of reaction, and no energy would be released.
In simple words, the endothermic reactions involve the process of energy intake. Also the exothermic reactions involve the release of energy.
Since energy is required in the endothermic reactions, it tends to reduce the temperature in the surroundings. Because of it, the cooling effect is generated.
Examples of endothermic reactions
Heat location is the quickest way to find if a reaction is endothermic process or exothermic process. It makes it very easy for a user to balance their chemical equation afterwards.
There are a lot of endothermic reaction examples takes place in our surrouding. Some examples of endothermic chemical reactions and processes are given below.
Endothermic Reaction Examples
- Mixing potassium chloride with water
- Dissolving NH4 in water
- Dissolving NH4Cl in water
- The reaction between dry NH4Cl with Ba(OH)2
- The reaction between sodium carbonate and ethanoic acid
- Photosynthesis in which carbon dioxide and water reacts with the chlorophyll in the presence of sunlight
Endothermic physical processes
- Dissolving salt into the water
- Melting ice cubes
- Formation of anhydrous salt
- Ion separation
- Splitting gas molecules
- Baking bread
- Frying an egg
Difference between endothermic and exothermic reactions
Both the terms "Exo and Endo" come from the Greek language, meaning "out and within".
We've described the fundamental difference between these two reactions at the start of the article.
Let's have a look at some other points of difference between endothermic and exothermic reactions to understand the concept fully
Endothermic vs Exothermic
|Exothermic reactions||Endothermic reactions|
|Heat is taken up from the surroundings to start the reaction. Thus, it causes cooling.||Heat is released from the system into the surroundings. Thus, it increases the temperature of the surroundings.|
|Positive enthalpy change (+ΔH)||Negative enthalpy change (-ΔH)|
|Entropy decreases (ΔS<0)||Entropy increases (ΔS>0)|
How to write chemical equations for an endothermic reaction?
- There's not any specialized way of writing the chemical equations for endothermic reactions. The rules for writing endothermic reactions are the same as they are for general reactions.
- Write the reactants always at the L.H.S of the equation and products at the R.H.S. It helps in the equation balancing process.
- Use an arrow sign to differentiate between these two reaction entities. This also shows the progress and direction of the chemical reaction.
- Heat is the driving force for the endothermic reactions. Endothermic reactions cannot occur in its absence. The heat is always shown on the reactant side.
- According to the rules mentioned above, the general equation for an endothermic reaction would be as follows
Reactants + Heat → Products
CaO + H2O + Heat → Ca(OH)2
Besides using the term heat in the equation, we can also show the nature of the reaction by enthalpy change.
Since the enthalpy change in the endothermic reaction is negative, the equation would be such as
Reactants → Products (ΔH= -x) Where x could be any numeral value.
C + O2 → CO2 (ΔH = -394 KJ)
How to determine whether the given reaction is endothermic or exothermic by equation?
Generally, the term "heat" is being used to determine the nature of the reaction.
If the term heat is present at the R.H.S of the equation, it will show the exothermic reactions.
It is because these reactions involve heat release. This becomes important when we are balancing redox reactions in a basic solution.
Since the endothermic reaction requires heat to occur, the heat would be shown at the reactant side of the equation.
Let's elaborate further with the example
2Na + H2O → NaOH + Heat (exothermic reaction)
2H2 + O + Heat → 2H2O (endothermic reaction)
Besides the term heat, the heat change/enthalpy change can also be shown by a triangle (ΔH)symbol.
In this case, the chemical equations for endothermic and exothermic reactions would be as follows
SO → S + O (ΔH = 297 KJ) exothermic reaction
N + 3H → 2NH (ΔH = -92 KJ) endothermic reaction
The question which arises here is, what if the heat or enthalpy change is not mentioned besides the equation?
In such conditions, you have to calculate the enthalpies of reactants and products.
On calculating the enthalpies of each reactant, sum up them.
If the sum of reactants enthalpy would be lesser than that of the product, the reaction will be endothermic in nature and vice versa.
Similarly, the potential energy diagram can also help you to predict the nature of the reaction.
We hope this article helped you learn about endothermic reactions and exothermic reactions.
You can also find step by step guides for balancing chemical equations here. Or learn how to find percent yield of a chemical reaction.