Endothermic and exothermic processes Knowledge-Based Study

Study notes

1. Energy in Chemical and Physical Changes

   In thermochemistry, every process is studied by dividing the universe into two parts:

   • System: The specific part of the universe being studied (such as the reactants and products in a beaker).
   • Surroundings: Everything outside the system (such as the solvent, the container, the air, and any measuring apparatus like a thermometer).

   Energy can be transferred between the system and its surroundings as heat ($q$). By convention, heat transfer is defined from the perspective of the system: heat entering the system has a positive sign ($q > 0$), whereas heat leaving the system has a negative sign ($q < 0$).

   Releasing Thermal Energy to the Surroundings

   An exothermic process is a physical change or chemical reaction that releases heat to its surroundings.

   Key Characteristics:

   • Heat Flow: Energy flows from the system to the surroundings ($q < 0$).
   • Enthalpy Change: The change in enthalpy is negative ($\Delta H < 0$), meaning the products have lower potential energy than the reactants.
   • Temperature Effect: Because thermal energy is transferred to the surroundings, the temperature of the surroundings increases.

   Common Examples:

   • Combustion: Burning of methane:
     CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) + Energy
   • Neutralization: Reaction between an acid and a base.
   • Physical Phase Changes: Freezing (liquid to solid), condensation (gas to liquid), and deposition (gas to solid).

   Absorbing Thermal Energy from the Surroundings

   An endothermic process is a physical change or chemical reaction that absorbs heat from its surroundings.

   Key Characteristics:

   • Heat Flow: Energy flows from the surroundings into the system ($q > 0$).
   • Enthalpy Change: The change in enthalpy is positive ($\Delta H > 0$), indicating that the products have higher potential energy than the reactants.
   • Temperature Effect: Because thermal energy is removed from the surroundings, the temperature of the surroundings decreases.

   Common Examples:

   • Thermal Decomposition: Decomposition of calcium carbonate:
     CaCO3(s) + Energy → CaO(s) + CO2(g)
   • Dissolution: Dissolving ammonium nitrate (NH4NO3) in water, commonly used in instant cold packs.
   • Physical Phase Changes: Melting (solid to liquid), vaporization (liquid to gas), and sublimation (solid to gas).

   Bond Breaking versus Bond Forming

   Chemical reactions involve rearranging atoms, which requires breaking existing bonds and forming new ones:

   • Breaking chemical bonds: Always requires an input of energy (an endothermic step, $\Delta H > 0$).
   • Forming chemical bonds: Always releases energy (an exothermic step, $\Delta H < 0$).

   The net enthalpy change of a reaction depends on the balance between these two contributions:

   $$\Delta H_{\text{reaction}} = \sum \text{Energy of bonds broken} - \sum \text{Energy of bonds formed}$$

   • If more energy is released during bond formation than is consumed during bond breaking, the overall reaction is exothermic ($\Delta H < 0$).
   • If more energy is absorbed to break bonds than is released when new bonds form, the overall reaction is endothermic ($\Delta H > 0$).

   Comparing Endothermic and Exothermic Changes

   To identify whether an observed process is endothermic or exothermic, consider the following diagnostic points:

   • Sign of $\Delta H$: Negative for exothermic; positive for endothermic.
   • Thermometer Reading: If the temperature of the reaction vessel or surrounding liquid rises, heat is leaving the reaction system, confirming an exothermic change. If the temperature drops, heat is entering the reaction system, confirming an endothermic change.
   • Position of Heat in an Equation: In thermochemical equations, heat written on the product side indicates an exothermic reaction, while heat written on the reactant side indicates an endothermic reaction.