This guide covers how energy, in the form of heat, is either released or absorbed during chemical reactions. Understanding the two main types of reactions and the simple calculation involved is essential.

1. Key Concepts: Exothermic vs. Endothermic Reactions

All chemical reactions involve a change in energy. We classify them based on whether they release heat into the surroundings or take heat in from the surroundings.

Exothermic Reactions

• Definition: A reaction that releases heat into the surroundings.

• Observation: The temperature of the surroundings increases. The reaction mixture feels hot.

• Energy Change: The products have less energy than the reactants. The extra energy is released as heat.

Energy Level Diagram: The reactants are shown at a higher energy level than the products, with an arrow pointing downwards to show energy being released.Energy

  ^

  |  Reactants

  |      |

  |      ↓ (Heat Released)

  |  Products

  +------------>

    Reaction Progress

• 
  

• Common Examples:

  • Combustion: Burning fuels like wood, petrol, or a candle.

  • Neutralization: The reaction between an acid and an alkali.

  • Cellular Respiration: The release of energy in living cells.

  • Adding water to quicklime (calcium oxide).

  • Dissolving solid Sodium Hydroxide (NaOH) in water.

Endothermic Reactions

• Definition: A reaction that absorbs heat from the surroundings.

• Observation: The temperature of the surroundings decreases. The reaction mixture feels cold.

• Energy Change: The products have more energy than the reactants. Energy must be taken in from the surroundings to make the reaction happen.

Energy Level Diagram: The reactants are shown at a lower energy level than the products, with an arrow pointing upwards to show energy being absorbed.Energy

  ^

  |  Products

  |      ↑ (Heat Absorbed)

  |      |

  |  Reactants

  +------------>

    Reaction Progress

• 
  

• Common Examples:

  • Photosynthesis: Plants absorb light energy to make food.

  • Thermal Decomposition: Breaking down a compound with heat, like heating calcium carbonate to produce quicklime.

  • Dissolving ammonium chloride or urea in water (this is why some instant cold packs work).

2. Calculating Heat Changes

You can calculate the amount of heat energy absorbed or released in a reaction (usually in a solution) using a specific formula.

• The Formula: Q = mcθ

  1. Q = Heat energy change (in Joules, J).

  2. m = mass of the substance being heated or cooled (usually the solution, in kilograms, kg).

  3. c = specific heat capacity of the substance (for water/dilute solutions, this is 4200 J kg⁻¹ °C⁻¹). This value will usually be given to you.

  4. θ (theta) = change in temperature (in degrees Celsius, °C). This is (Final Temperature - Initial Temperature).

• Important Assumptions for Experiments:

  1. We assume the density of the solution is the same as water (1 g/cm³ or 1000 kg/m³). This lets us use the volume to find the mass (e.g., 100 cm³ of solution has a mass of 100 g or 0.1 kg).

  2. We assume the specific heat capacity of the solution is the same as water (4200 J kg⁻¹ °C⁻¹).

  3. We assume no heat is lost to the surroundings or the apparatus (like the beaker).

Exam Tips & Tricks

1. Exo = Exit, Endo = Enter: This is the easiest way to remember. Exothermic reactions let heat exit. Endothermic reactions need heat to enter.

2. Temperature is the Clue: If the temperature goes UP, it's EXOthermic. If the temperature goes DOWN, it's ENDOthermic. This is a very common way questions are phrased.

3. Watch Your Units in Q = mcθ: The most common mistake is forgetting to convert the mass from grams (g) to kilograms (kg) before calculating. Divide the mass in grams by 1000.

4. Know Your Examples: Be ready to classify everyday reactions. Combustion and neutralization are the classic examples of exothermic reactions. Photosynthesis is the ultimate endothermic reaction.

5. Energy Level Diagrams: Be prepared to sketch or interpret these simple diagrams. Just remember: for exothermic, reactants are high and products are low; for endothermic, reactants are low and products are high.

Important Points to Remember

• Energy is conserved in chemical reactions; it just changes form or is transferred.

• Exothermic: Reactants → Products + Heat (Products are more stable).

Endothermic: Reactants + Heat → Products (Products are less stable).