Chemistry (0620) | Exothermic and endothermic reactions

Hello IGCSE Chemists! Understanding Heat Changes in Reactions

Welcome to Chemical Energetics! This is one of the most exciting topics because it explains why some reactions feel hot and others feel cold. Every chemical reaction involves a change in energy, usually thermal energy (heat).

In these notes, we will look at the two main ways reactions handle heat: releasing it (making things hot!) or absorbing it (making things cold!). Understanding this concept is essential not only for your exams but for understanding processes like combustion and even instant cold packs. Let’p>

1. Core Concepts: Exothermic and Endothermic Reactions

1.1 Exothermic Reactions: Heat Out!

Think of a bonfire or turning on your heating at home. These processes release heat energy into the room.

• Definition: An exothermic reaction is a reaction that transfers thermal energy to the surroundings.
• Observation: This transfer of energy leads to an increase in the temperature of the surroundings. The container or test tube holding the reaction will feel hot.
• Real-World Examples:
  • Combustion (Burning): Methane gas reacting with oxygen to produce heat and light.
  • Neutralisation: Adding acid to alkali.
  • Respiration: The process our bodies use to release energy from food.

💡 Memory Aid for Exothermic: "EXO" sounds like "EXIT". Heat EXITS the reaction and enters the surroundings.

1.2 Endothermic Reactions: Heat In!

Think of using an instant cold pack for a sports injury. The pack feels cold because it is sucking heat energy away from your skin.

• Definition: An endothermic reaction is a reaction that takes in thermal energy from the surroundings.
• Observation: Since the reaction is absorbing heat from its surroundings, this leads to a decrease in the temperature of the surroundings. The container will feel cold.
• Real-World Examples:
  • Thermal Decomposition: Heating limestone (calcium carbonate).
  • Melting Ice: (Although this is a physical change, it requires absorbing energy).
  • Instant Cold Packs: Often involve the dissolving of a salt, like ammonium nitrate, which requires a large input of energy.

💡 Memory Aid for Endothermic: "ENDO" sounds like "ENTER". Heat ENTERS the reaction from the surroundings.

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2. Extended Concept: Enthalpy Change (\(\Delta H\))

For Extended candidates, we use a formal chemistry term to describe the heat transferred during a reaction: Enthalpy Change.

2.1 Defining Enthalpy Change ($\Delta H$)

The enthalpy change ($\Delta H$) is the transfer of thermal energy during a reaction (measured in $\text{kJ/mol}$). It represents the difference in energy between the products and the reactants.

• For Exothermic Reactions: The products have less energy stored inside them than the reactants did, so energy is released.
  $\Delta H$ is negative.
  (The system lost energy, so the change is negative.)
• For Endothermic Reactions: The products have more energy stored inside them than the reactants did, because energy had to be absorbed from the surroundings.
  $\Delta H$ is positive.
  (The system gained energy, so the change is positive.)

⚠️ Common Mistake Alert: Students often confuse $\Delta H$ signs! Remember: A cold endothermic reaction has a positive $\Delta H$. The surroundings feel cold, but the system itself gained potential energy.

3. Reaction Pathway Diagrams (Energy Profile Diagrams)

We can visualise the energy changes in a reaction using a diagram. This diagram plots the potential energy of the substances as the reaction progresses.

3.1 Activation Energy ($E_a$) (Supplement)

Before a reaction can start, particles must collide with enough energy to break the existing bonds.

• Definition: The activation energy ($E_a$) is the minimum energy that colliding particles must have in order to react.
• Analogy: Imagine rolling a ball over a hill. Even if the valley on the other side is lower (exothermic), you still need to give the ball a strong push ($E_a$) to get it over the top!

3.2 Drawing and Interpreting the Diagrams (Core & Supplement)

All diagrams must have the Y-axis labeled Energy and the X-axis labeled Reaction Pathway or Progress of Reaction.

Exothermic Reaction Pathway Diagram

The energy of the products is lower than the energy of the reactants. The overall energy released ($\Delta H$) is negative.

• (A) Reactants: Starting energy level.
• (B) Products: Final energy level (lower than reactants).
• (C) Activation Energy ($E_a$): The height of the peak, measured from the Reactants level up to the transition state (the top of the hill). This energy must be supplied to start the reaction.
• (D) Enthalpy Change ($\Delta H$): The vertical difference between the Reactants and Products level. Since products are lower, energy was released, so $\Delta H$ is negative.

(Note: In a true HTML format, drawing the diagram is impossible, but the interpretation points above cover the syllabus requirements for labeling).

Endothermic Reaction Pathway Diagram

The energy of the products is higher than the energy of the reactants. The overall energy absorbed ($\Delta H$) is positive.

• (A) Reactants: Starting energy level.
• (B) Products: Final energy level (higher than reactants).
• (C) Activation Energy ($E_a$): The height of the peak, measured from the Reactants level up to the transition state. This is often larger than in an exothermic reaction.
• (D) Enthalpy Change ($\Delta H$): The vertical difference between the Reactants and Products level. Since products are higher, energy was absorbed, so $\Delta H$ is positive.

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4. Energy and Chemical Bonds (Supplement)

Chemical reactions involve rearranging atoms. To rearrange them, we must first break the bonds that hold the reactants together, and then form new bonds to create the products.

4.1 Bond Breaking vs. Bond Making

Energy changes in a reaction are determined by the balance between the energy needed to break bonds and the energy released when new bonds are formed.

• Bond Breaking: This process always requires energy input. You must put energy in to pull atoms apart. Therefore, bond breaking is an endothermic process.
• Bond Making: This process always releases energy. Energy is released when atoms join together to form stable bonds. Therefore, bond making is an exothermic process.

The overall enthalpy change ($\Delta H$) of the reaction is the sum of these two steps.

$$ \Delta H = (\text{Energy required to break bonds}) + (\text{Energy released from making bonds}) $$

Why a reaction is Exothermic or Endothermic:

1. If the energy released (bond making) is greater than the energy required (bond breaking), the reaction is exothermic ($\Delta H$ is negative).
2. If the energy required (bond breaking) is greater than the energy released (bond making), the reaction is endothermic ($\Delta H$ is positive).

4.2 Calculating Enthalpy Change using Bond Energies (Supplement)

The syllabus requires you to calculate $\Delta H$ using bond energies. Bond energy is the amount of energy needed to break one mole of a specific type of bond.

Here is the step-by-step process:

Step 1: Calculate Total Energy Absorbed (Endothermic Part)

Sum the energies of all the bonds in the reactants. This is the energy input (positive value).

Step 2: Calculate Total Energy Released (Exothermic Part)

Sum the energies of all the bonds in the products. This is the energy output (negative value).

Step 3: Calculate the Overall Enthalpy Change ($\Delta H$)

$$ \Delta H = \sum (\text{Bond energies of reactants}) - \sum (\text{Bond energies of products}) $$

(Alternatively written as: $\Delta H = \text{Energy In} - \text{Energy Out}$)

Example Walkthrough (Combustion of Methane):

Imagine you are calculating the $\Delta H$ for: $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$

1. Bonds Broken (Reactants): You need to break 4 C–H bonds and 2 O=O bonds. Add up the energy needed for these 6 bonds. (Positive energy, $\text{E}_{\text{in}}$).
2. Bonds Formed (Products): You form 2 C=O bonds (in $\text{CO}_2$) and 4 O–H bonds (in two $\text{H}_2\text{O}$ molecules). Add up the energy released when these 6 bonds form. (Negative energy, $\text{E}_{\text{out}}$).
3. Result: Since combustion is highly exothermic, the energy released ($\text{E}_{\text{out}}$) will be much larger than the energy absorbed ($\text{E}_{\text{in}}$), resulting in a large negative $\Delta H$ value.

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Quick Chapter Summary

• Exothermic: Releases heat, temperature increases, $\Delta H$ is negative.
• Endothermic: Absorbs heat, temperature decreases, $\Delta H$ is positive.
• Activation Energy ($E_a$): Minimum energy needed to start the reaction (the energy "hill").
• Bond Rule: Breaking bonds requires energy (Endo). Making bonds releases energy (Exo).
• Calculation: $\Delta H = \text{Energy needed to break bonds} - \text{Energy released by forming bonds}$.

Don't worry if the calculation seems tricky at first. Practice drawing the pathway diagrams and remembering the $\Delta H$ sign convention—that's half the battle! Good luck!