Unit 2: Energy Changes in Chemical Reactions

Energy In Chemical Reactions

• Energy is a crucial factor in chemical processes and reactions.
• Reactants must overcome a specific minimum energy to initiate a reaction, called activation energy.
  • This can be thought of as a barrier to be overcome (analogous to a hill).

System and Surroundings

• System: The materials involved in the chemical reaction.
• Surroundings: Everything else in the universe.
• Law of Conservation of Energy states that total energy in the universe is constant.
  • Energy cannot be created or destroyed; it only changes forms.
  • Energy entering the system must come from the surroundings, and vice versa.

Energy Transfers in Reactions

• Reactions can either absorb or release heat due to the making and breaking of bonds:
  • Making chemical bonds releases heat energy (exothermic).
  • Breaking chemical bonds requires energy input (endothermic).
  • Heat input is often necessary to start a reaction.

Exothermic Reactions

• Definition: Reactions that release heat.
• Example: Magnesium reacting with hydrochloric acid.
  • Energy diagram shows reactants have more energy than products.
  • The heat released can be used to keep the reaction ongoing (e.g., rocket fuel).
  • Summary:
  • More energy is released by reactants than absorbed by products.
  • Excess energy is emitted as heat.
  • Activation energy may be needed initially.

Endothermic Reactions

• Definition: Reactions that absorb heat.
• Example: Ammonium nitrate in an instant ice pack.
  • Reactants absorb energy from surroundings, resulting in temperature decrease.
  • Summary:
  • More energy is taken in than released.
  • Creates a cooling effect.

Energy Changes in Reactions

• Two types of energy changes:
  • Endothermic: Heat absorbed from surroundings.
  • Exothermic: Heat released to surroundings.

Bond Breaking and Formation

• Energy Required: Energy must be absorbed to break bonds in reactants.
• Energy Released: Energy is released when new bonds form in products.
• General rules:
  • Bond breaking is endothermic; it requires energy.
  • Bond forming is exothermic; it releases energy.

Measuring Energy Changes

• Energy changes can be observed through temperature changes during reactions:
  • If temperature increases, the reaction is exothermic.
  • If temperature decreases, the reaction is endothermic.
• Enthalpy (ΔH):
  • Represents energy changes in reactions.
  • ΔH = H(products) - H(reactants)
  • Refers to total energy stored within chemical bonds.

Energy-Level Diagrams

• Diagrams illustrate the energy of reactants and products.
  • Activation Energy (EA): Energy difference between reactants and the transition state.
  • Enthalpy Change (ΔH): Difference between energy levels of reactants and products.
  • Endothermic vs. Exothermic:
  • Endothermic: Overall ΔH positive, energy absorbed.
  • Exothermic: Overall ΔH negative, energy released.

Thermochemical Equations

• Balanced equations include ΔH to indicate energy changes:
  • Exothermic: Energy on product side (heat released).
  • Endothermic: Energy on reactant side (heat absorbed).