Exothermic and endothermic reactions

1. Overview

All chemical reactions involve a transfer of energy between the reaction system and its surroundings. This energy is usually transferred as thermal energy (heat), which determines whether a reaction feels hot or cold. Understanding these energy changes is vital for controlling industrial processes and predicting how substances will react.

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Key Definitions

• Exothermic Reaction: A reaction that transfers thermal energy to the surroundings, leading to an increase in the temperature of the surroundings.
• Endothermic Reaction: A reaction that takes in thermal energy from the surroundings, leading to a decrease in the temperature of the surroundings.
• Enthalpy Change ($\Delta H$): The transfer of thermal energy during a reaction (at constant pressure).
• Activation Energy ($E_a$): The minimum energy that colliding particles must have to react.
• Bond Energy: The amount of energy required to break one mole of a specific covalent bond.

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Core Content

Exothermic Reactions

• Energy is released to the surroundings.
• The temperature of the surroundings increases.
• Examples: Combustion of fuels, respiration, and the reaction between magnesium and hydrochloric acid.
  • Word Equation: Magnesium + hydrochloric acid → magnesium chloride + hydrogen
  • Symbol Equation: $Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)$

Endothermic Reactions

• Energy is absorbed from the surroundings.
• The temperature of the surroundings decreases.
• Examples: Photosynthesis, thermal decomposition of calcium carbonate, and the reaction between citric acid and sodium hydrogencarbonate.
  • Word Equation: Calcium carbonate → calcium oxide + carbon dioxide
  • Symbol Equation: $CaCO_3(s) \rightarrow CaO(s) + CO_2(g)$

Reaction Pathway Diagrams (Core)

These diagrams show the relative energy levels of reactants and products.

• Exothermic: The products are at a lower energy level than the reactants because energy has been lost to the surroundings.
• Endothermic: The products are at a higher energy level than the reactants because energy has been absorbed.

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Extended Content (Extended Curriculum Only)

Enthalpy Change ($\Delta H$)

The energy change is expressed as $\Delta H$ (measured in kJ/mol):

• Exothermic: $\Delta H$ is negative ($-$) because the system loses energy.
• Endothermic: $\Delta H$ is positive ($+$) because the system gains energy.

Activation Energy ($E_a$)

Even exothermic reactions require an initial "push" to start. This is the activation energy ($E_a$), represented by the height of the "hump" from the reactant level to the peak of the curve on a diagram.

Detailed Reaction Pathway Diagrams

When drawing these for exams, you must label:

1. Reactants and Products on their respective energy levels.
2. Enthalpy Change ($\Delta H$): The vertical difference between reactants and products.
3. Activation Energy ($E_a$): The arrow from the reactant level to the peak of the curve.

Bond Breaking and Bond Making

• Bond breaking is endothermic: Energy must be taken in to break chemical bonds.
• Bond making is exothermic: Energy is released when new bonds form.
• The overall enthalpy change depends on the balance:
  • If more energy is released making bonds than is taken in breaking them, the reaction is exothermic.
  • If more energy is taken in breaking bonds than is released making them, the reaction is endothermic.

Calculating Enthalpy Change using Bond Energies

Formula: $\Delta H = \text{Energy taken in to break bonds} - \text{Energy released making bonds}$

Worked Example: Calculate the enthalpy change for the reaction: $H_2(g) + Cl_2(g) \rightarrow 2HCl(g)$ Bond energies: H-H = 436 kJ/mol; Cl-Cl = 242 kJ/mol; H-Cl = 431 kJ/mol

1. Energy in (breaking bonds):
   • 1 × (H-H) = 436
   • 1 × (Cl-Cl) = 242
   • Total In = $436 + 242 = 678 \text{ kJ/mol}$
2. Energy out (making bonds):
   • 2 × (H-Cl) = $2 \times 431 = 862 \text{ kJ/mol}$
3. Enthalpy Change ($\Delta H$):
   • $\Delta H = 678 - 862 = -184 \text{ kJ/mol}$
   • Since the value is negative, the reaction is exothermic.

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Key Equations

• Enthalpy Change: $\Delta H = \text{Energy}{broken} - \text{Energy}{formed}$
  • $\Delta H$: Enthalpy Change (kJ/mol)
  • Positive value ($+$): Endothermic
  • Negative value ($-$): Exothermic

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Common Mistakes to Avoid

• ❌ Wrong: Thinking that exothermic reactions "take in" heat because they feel hot.
• ✅ Right: Exothermic reactions release heat, which is why the surroundings (and your thermometer) get hotter.
• ❌ Wrong: Drawing the $\Delta H$ arrow from the peak of the graph.
• ✅ Right: $\Delta H$ is only the difference between the reactants and products.
• ❌ Wrong: Forgetting to multiply bond energies by the coefficients in the balanced equation (e.g., $2HCl$ means $2 \times H-Cl$ bonds).

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Exam Tips

• Command Words: If asked to "Interpret" a diagram, look at whether the product line is higher or lower than the reactant line to identify the reaction type.
• Calculations: Always show your working in three clear steps: (1) Bonds broken, (2) Bonds made, (3) The subtraction. This ensures partial marks if you make a calculator error.
• Signs: Always include the $+$ or $-$ sign in your final answer for $\Delta H$.
• Real-world context: Be prepared to identify reactions in context. Portable hand-warmers use exothermic reactions, while instant ice packs use endothermic reactions.