Endothermic Reactions: Reactions that absorb energy from the surroundings.
Example: Photosynthesis, which uses sunlight to produce glucose.
Enthalpy (H)
Definition: A measure of the total energy of a thermodynamic system.
Change in Enthalpy (ΔH): The difference between the enthalpy of products and reactants.
ΔH is negative for exothermic reactions (energy released).
ΔH is positive for endothermic reactions (energy absorbed).
Bond Energies
Breaking Bonds: Requires energy (endothermic).
Forming Bonds: Releases energy (exothermic).
Overall Reaction Energy: The total energy change is the sum of bond energies of reactants minus the sum of bond energies of products.
Stability and Activation Energy
Stability: Lower energy states are more stable.
Activation Energy: The initial energy required to start a reaction by breaking bonds in reactants.
High activation energy can prevent spontaneous reactions even if they are exothermic.
Practical Applications: Calorimetry
Calorimetry: An experimental technique to measure the heat absorbed or released during a reaction.
Involves conducting the reaction in an insulated container to prevent heat loss.
Calorimetry Calculation
Heat (Q): Calculated using the equation Q = mCΔT
m: Mass of the substance (in grams).
C: Specific heat capacity (4.18 J/g·K for water).
ΔT: Change in temperature (in Celsius or Kelvin).
Example Calculation
Scenario: Reaction in 100g of water increases temperature from 20°C to 30°C.
Calculation:
Q = 100 g * 4.18 J/g·K * 10 K = 4,180 J
Conclusion: The reaction gave off 4,180 Joules of heat.
Summary
We explored the concepts of exothermic and endothermic reactions, changes in enthalpy, the importance of bond energies, and the role of activation energy.
Calorimetry is a practical approach to measure these heat changes, providing insight into the energetics of chemical reactions.