5.1 Chemical Energetics – Exothermic and Endothermic Reactions

Table of Contents

5.1 Exothermic and Endothermic Reactions

Chemical Energetics – IGCSE Chemistry Study Notes

5.1 Exothermic and Endothermic Reactions #

IGCSE Chemistry Topic 5 – Chemical Energetics

Chemical reactions always involve energy changes.

Think of it like this – to take molecules apart (break bonds), you need to put energy in, like pulling apart two magnets.

When molecules come together (make bonds), they release energy, like when magnets snap together.

The overall energy change tells us if a reaction is exothermic or endothermic.

Enthalpy Change and Activation Energy #

Before we look at exothermic and endothermic reactions, we need to understand two important energy concepts that apply to all chemical reactions: enthalpy change and activation energy. These concepts help us measure and understand the energy changes that happen during reactions.

What is Enthalpy Change (ΔH)? #

Enthalpy change (ΔH) is the scientific way to measure how much heat energy is transferred during a chemical reaction. Think of it as the “energy difference” between the reactants (what you start with) and the products (what you end up with). We measure enthalpy change in kilojoules per mole (kJ/mol).

Simple Way to Remember: Enthalpy change tells us if a reaction gives out heat or takes in heat, and exactly how much energy is involved.

The Sign of ΔH Tells Us the Type of Reaction:
ΔH is negative (-) → Exothermic reaction (gives out heat)
ΔH is positive (+) → Endothermic reaction (takes in heat)

Why the Sign Matters: The negative sign shows that energy is leaving the chemical system and going to the surroundings (making them warmer). The positive sign shows that energy is entering the chemical system from the surroundings (making them cooler).

Examples:
• Burning methane: ΔH = -890 kJ/mol (exothermic – releases 890 kJ of heat per mole)
• Photosynthesis: ΔH = +2802 kJ/mol (endothermic – absorbs 2802 kJ of heat per mole)

What is Activation Energy (Ea)? #

Activation energy (Ea) is the amount of energy that particles need when they collide for a chemical reaction to start. Even if a reaction will release energy overall, the particles still need this initial “push” of energy to get the reaction going.

Think of it Like This: Imagine rolling a ball over a hill. Even if the ball will roll down the other side by itself (releasing energy), you still need to give it enough energy to get it over the top of the hill first. That initial energy needed is like activation energy.

Important Facts About Activation Energy:

ALL reactions need activation energy to start (even exothermic ones)
Activation energy is always positive (energy must be added)
Higher activation energy = reaction is harder to start
Lower activation energy = reaction starts more easily
Catalysts work by lowering the activation energy

Real-Life Example: A match needs activation energy from striking to start burning. Once it starts, the burning reaction releases much more energy than was needed to light it. This is why matches don’t just burst into flames by themselves – they need that initial energy input.

Exothermic Reactions #

An exothermic reaction is a chemical reaction that transfers thermal energy (heat) to the surroundings. This means the reaction gives out heat, making the surroundings warmer. You can feel this happening – if you touch the container where an exothermic reaction is taking place, it will feel hot.

Key Features of Exothermic Reactions:

Temperature of surroundings increases
Heat energy is released to the surroundings
Products have less energy than reactants
More energy is released when making new bonds than is needed to break the old bonds
The enthalpy change (ΔH) is -negative

Common Examples: Combustion (burning), respiration in cells, neutralization reactions between acids and alkalis, and the reaction between water and calcium oxide (quicklime).

Endothermic Reactions #

An endothermic reaction is a chemical reaction that takes in thermal energy (heat) from the surroundings. This means the reaction absorbs heat, making the surroundings cooler. If you touch the container where an endothermic reaction is happening, it will feel cold.

Key Features of Endothermic Reactions:

Temperature of surroundings decreases
Heat energy is absorbed from the surroundings
Products have more energy than reactants
More energy is needed to break the old bonds than is released when making new bonds
The enthalpy change (ΔH) is +positive

Common Examples: Photosynthesis in plants, thermal decomposition reactions (like heating copper carbonate), dissolving ammonium chloride in water, and the reaction between citric acid and sodium hydrogencarbonate.

Reaction Pathway Diagrams #

Reaction pathway diagrams (also called energy profile diagrams) show how energy changes during a chemical reaction. These diagrams help us visualize what happens to energy as reactants turn into products. The vertical axis shows energy, and the horizontal axis shows the progress of the reaction from start to finish.

How to Read Reaction Pathway Diagrams: Start at the left with reactants. Follow the line as it goes up (energy needed to start the reaction) then down to products on the right. If products are lower than reactants, it’s exothermic. If products are higher than reactants, it’s endothermic.

Supplement: Enthalpy Change (ΔH) #

The enthalpy change (ΔH) is the scientific way to measure the heat energy transferred during a chemical reaction. It tells us exactly how much energy is released or absorbed. We measure it in kilojoules per mole (kJ/mol).

Remember:

Exothermic reactions: ΔH is negative (e.g., ΔH = -286 kJ/mol)

Endothermic reactions: ΔH is positive (e.g., ΔH = +178 kJ/mol)

Why the Sign Matters: The negative sign for exothermic reactions shows energy is leaving the system (going to surroundings). The positive sign for endothermic reactions shows energy is entering the system (coming from surroundings).

Activation Energy (Ea) #

Activation energy (Ea) is the minimum energy that particles must have when they collide for a reaction to happen. Think of it like the energy needed to push a ball over a hill – even if the ball will roll down the other side by itself, you still need to give it a push to get it over the top first.

Important Point: All reactions need activation energy to start, even exothermic reactions that release energy overall. This is why a match needs a strike to light, even though burning is exothermic.

Facts About Activation Energy:

It’s always positive (energy must be added)
Lower activation energy means reactions happen more easily
Catalysts work by lowering the activation energy
On energy diagrams, it’s the height from reactants to the peak

Drawing and Labeling Reaction Pathway Diagrams #

When drawing reaction pathway diagrams for exams, you must include four key parts: reactants, products, enthalpy change (ΔH), and activation energy (Ea). Each part must be clearly labeled.

How to Draw a Reaction Pathway Diagram:

Draw axes: Vertical axis = Energy, Horizontal axis = Progress of reaction
Mark reactants: Draw a horizontal line on the left showing the energy level of reactants
Draw the curve: Draw a sqaure curve going up (activation energy) across and then down to products
Mark products: Draw a horizontal line on the right showing the energy level of products
Add arrow for ΔH: Draw a vertical arrow from reactants to products level
Add arrow for Ea: Draw a vertical arrow from reactants to the peak of the curve
Label everything: Label reactants, products, ΔH, and Ea

Bond Breaking and Bond Making #

Every chemical reaction involves breaking bonds in the reactants and making new bonds in the products. Understanding the energy involved in these processes is key to understanding why reactions are exothermic or endothermic.

1. Bond breaking is ALWAYS Endothermic (needs energy input)
2. Bond making is ALWAYS Exothermic (releases energy)

Why This Makes Sense: Breaking bonds is like pulling apart magnets – you need to put energy in to separate them. Making bonds is like letting magnets snap together – they release energy when they connect.

The overall energy change of a reaction depends on the balance between energy needed to break bonds and energy released when making bonds. If more energy is released than needed, the reaction is exothermic. If more energy is needed than released, the reaction is endothermic.

The Key Formula:
$\Delta H = \text{Energy to break bonds} – \text{Energy released when bonds form}$

Calculating Enthalpy Change Using Bond Energies #

Bond energy is the amount of energy needed to break one mole of a specific type of bond. We can use bond energies to calculate the overall enthalpy change (ΔH) of a reaction. This involves adding up all the energy needed to break bonds, then subtracting all the energy released when new bonds form.

Worked Example: Formation of Ammonia #

Reaction: N₂ + 3H₂ → 2NH₃

N≡N

H—H

→

H—N—H

Given bond energies:

Bond	N≡N	H—H	N—H
Bond energy (kJ/mol)	945	435	390

Step 1: Count and calculate energy to break bonds

Bonds broken in reactants:

• 1 × N≡N bond = 1 × 945 = 945 kJ

• 3 × H—H bonds = 3 × 435 = 1305 kJ

Total energy to break bonds = 945 + 1305 = 2250 kJ

Step 2: Count and calculate energy released when bonds form

Bonds formed in products:

• 6 × N—H bonds

• 6 × 390 = 2340 kJ

Total energy released = 2340 kJ

Step 3: Calculate ΔH

$\Delta H = \text{Energy to break bonds} – \text{Energy released when bonds form}$
$\Delta H = 2250 – 2340 = -90 \text{ kJ/mol}$

Answer: ΔH = -90 kJ/mol (negative, so the reaction is exothermic)

Example 2: Decomposition of Water #

Reaction: 2H₂O → 2H₂ + O₂

H—O—H

→

H—H

O=O

Given bond energies:

Bond	O—H	H—H	O=O
Bond energy (kJ/mol)	464	435	498

Step 1: Count and calculate energy to break bonds

Bonds broken in reactants:

• 4 × O—H bonds

• 4 × 464 = 1856 kJ

Total energy to break bonds = 1856 kJ

Step 2: Count and calculate energy released when bonds form

Bonds formed in products:

• 2 × H—H bonds = 2 × 435 = 870 kJ

• 1 × O=O bond = 1 × 498 = 498 kJ

Total energy released = 870 + 498 = 1368 kJ

Step 3: Calculate ΔH

$\Delta H = \text{Energy to break bonds} – \text{Energy released when bonds form}$
$\Delta H = 1856 – 1368 = +488 \text{ kJ/mol}$

Answer: ΔH = +488 kJ/mol (positive, so the reaction is endothermic)

Tips for Bond Energy Calculations:

Always draw out the structural formulas with all bonds shown – this makes counting much easier
Count bonds carefully – remember that N₂ has a triple bond (N≡N) and O₂ has a double bond (O=O)
Don’t forget to multiply by the number of molecules (use the coefficients in the equation)
Check your sign: negative ΔH = exothermic, positive ΔH = endothermic

Summary of Key Points #

Essential Facts to Remember:

Exothermic reactions: Release heat to surroundings, temperature increases, ΔH is negative
Endothermic reactions: Absorb heat from surroundings, temperature decreases, ΔH is positive
Activation energy (Ea): Minimum energy needed for particles to react when they collide
Bond breaking: Always endothermic (needs energy)
Bond making: Always exothermic (releases energy)
Calculating ΔH: Energy to break bonds minus energy released when bonds form
Reaction pathway diagrams: Must show:
- reactants
- products
- ΔH
- Activation Energy – Ea

Exam Tip: When drawing bonds for calculations, always show every single bond clearly. For example, show H₂ as H—H, not just H₂. Show NH₃ with all three N—H bonds visible. This helps you count bonds accurately and avoid mistakes in calculations.