Overview
This lecture covers the fundamentals of chemical energetics, focusing on exothermic and endothermic reactions, enthalpy changes, standard conditions, bond energies, calorimetry calculations, and Hess's Law.
Exothermic and Endothermic Reactions
- Exothermic reactions release heat to the surroundings, raising the temperature (e.g., combustion).
- Endothermic reactions absorb heat from the surroundings, lowering the temperature.
Reaction Profile Diagrams
- Reactants start with a certain energy; products in exothermic reactions have less energy than reactants (ΔH negative).
- Activation energy is the minimum energy needed to start a reaction, shown as the energy difference from reactants to the top of the graph.
- For endothermic reactions, products have more energy than reactants (ΔH positive).
Standard Enthalpy Changes and Conditions
- Standard conditions: 1 atm pressure (100 kPa), 298 K (25°C), 1 mol/dm³ concentration.
- Standard state: physical state of a substance at standard conditions.
- ΔH is measured under these conditions and denoted ΔH°.
Types of Standard Enthalpy Changes
- Standard enthalpy change of reaction (ΔH°_rxn): Heat change when reactants convert to products in their standard states under standard conditions.
- Standard enthalpy change of formation (ΔH°_f): Heat change when one mole of a compound forms from its elements in standard states.
- ΔH°_f of any element in its standard state is zero.
- Standard enthalpy change of combustion (ΔH°_c): Heat released when one mole of a substance is burned in excess oxygen under standard conditions.
- ΔH°_c for elements like O₂ or H₂O is zero.
- Standard enthalpy change of neutralization: Heat change when one mole of water forms from acid and base under standard conditions.
- Standard enthalpy change of atomization: Energy needed to form one mole of gaseous atoms from the element in its standard state.
Bond Energies
- Bond energy: Energy required to break one mole of bonds in gaseous molecules.
- Breaking bonds requires energy (endothermic); forming bonds releases energy (exothermic).
- ΔH calculation: ΔH = (energy to break bonds) - (energy released forming bonds).
- Stronger (shorter) bonds require more energy to break.
- Covalent bonds are stronger than hydrogen bonds, which are stronger than van der Waals forces.
Calorimetry and Enthalpy Calculations
- Q (energy change) = mass (g) × specific heat capacity (4.18 J/g·K) × temperature change (°C).
- ΔH = Q / n (moles of limiting reactant); sign is negative for exothermic, positive for endothermic.
- Mass of solution ≈ volume in cm³ for aqueous solutions.
Experimental Errors
- Calculated ΔH may differ from actual due to heat loss, incomplete combustion, evaporation, or apparatus absorption.
Hess’s Law
- The total enthalpy change is independent of the reaction pathway if initial and final states are the same.
- Use enthalpy of combustion or formation data to construct cycles and solve for unknown ΔH values.
Key Terms & Definitions
- Exothermic reaction — reaction that releases heat.
- Endothermic reaction — reaction that absorbs heat.
- Enthalpy change (ΔH) — heat change at constant pressure.
- Standard state — physical state of a substance under standard conditions.
- Activation energy — minimum energy required to start a reaction.
- Bond energy — energy needed to break one mole of specific bonds in gaseous state.
- Hess's Law — total enthalpy change is path-independent.
Action Items / Next Steps
- Memorize definitions for all types of standard enthalpy changes.
- Practice drawing reaction profile and energy level diagrams.
- Solve problems on ΔH using bond energies and calorimetry.
- Complete additional questions on Hess’s Law as recommended.