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Chemical Reactions and Energy

• Every chemical reaction involves energy transfer, either absorbed or released.

• Reactions Types:

• Endothermic: Absorbs energy (e.g., ( ext{A} + ext{B} + ext{Heat} \rightarrow ext{C} ))

• Exothermic: Releases energy (e.g., ( ext{A} + ext{B} \rightarrow ext{C} + ext{Heat} ))

Enthalpy of Reaction (9;ΔH9;)

• Endothermic Reactions:

• ( ext{ΔH is positive (+)} )

• Example: ( 2 ext{HCl(g)} \rightarrow ext{H}_2(g) + ext{Cl}_2(g), ext{ΔH} = 184.6 ext{ kJ/mol} )

• Exothermic Reactions:

  • ( ext{ΔH is negative (-)} )

  • Example: ( 2 ext{H}_2(g) + ext{O}_2(g) \rightarrow 2 ext{H}_2 ext{O}(g), ext{ΔH} = -483.6 ext{ kJ/mol} )

Energy Changes in Chemical Reactions

• Bonds have different energy values.

• • Bond Breaking: Requires energy (endothermic)

  • Bond Formation: Releases energy (exothermic)

Bond Energy Overview

• Bond energies are the energies required to break specific bonds.

• More stable (stronger) bonds have higher bond energies:

• Example Values:

  • Cl-Cl = 243 kJ/mol

  • H-H = 436 kJ/mol

  • O-H = 463 kJ/mol

  • C-H = 413 kJ/mol

  • O=O = 498 kJ/mol

  • C=O = 803 kJ/mol

• Calculate overall energy transfer during the reaction by comparing energies needed to break old bonds versus energies released from new bonds formed.

Estimating Energy Changes in Reactions

Example 1

• Reaction: ( 2 ext{H}_2 ext{O(g)} \rightarrow 2 ext{H}_2(g) + ext{O}_2(g) )

• Consider energy to break/reactants vs. form/products to determine ( ext{ΔH} ).

Example 2

• Reaction: ( ext{CH}_4(g) + 2 ext{O}_2(g) \rightarrow ext{CO}_2(g) + 2 ext{H}_2 ext{O}(g) )

Summary: Exothermic vs. Endothermic Reactions

• Exothermic Reaction: ( ext{CH}_4(g) + 2 ext{O}_2(g) \rightarrow ext{CO}_2(g) + 2 ext{H}_2 ext{O}(g), ext{ΔH} = -810 ext{ kJ/mol} )

• Endothermic Reaction: ( 2 ext{H}_2 ext{O}(g) \rightarrow 2 ext{H}_2(g) + ext{O}_2(g), ext{ΔH} = +482 ext{ kJ/mol} )

Reaction Rate: Speed of Reactions

• Measurement: Change in concentration over time (e.g., moles per liter per minute).

Factors Influencing Reaction Rates

• High rate: Reaction forms products quickly.

• Activation Energy: Minimum energy needed for reactants to collide effectively. Must overcome the energy barrier.

  • Exothermic and endothermic reactions defined through energy relationships.

Catalysts

• Catalysts increase the reaction rate without being consumed.

• They work by lowering activation energy, enabling more molecules to react.

Chemical Equilibrium

• Definition: A state where the forward and reverse reactions occur at the same rate, maintaining concentrations of reactants and products.

• Disturbances to equilibrium can shift the reaction towards products or reactants based on Le Chatelier's principle.

Le Chatelier's Principle

• When a system at equilibrium is disturbed, it responds to offset the disturbance.

• Reactions may shift towards products or reactants; can be affected by changes in concentration, temperature, and pressure.

Disturbing Equilibrium: Effects of Concentration Changes

• Examples of how equilibrium shifts with different stresses on systems.

Disturbing Equilibrium: Effects of Temperature Changes

• Understanding shifts for endothermic and exothermic reactions when temperature changes.

Additional Factors Influencing Rate and Equilibrium

• Effects of adding a catalyst on the reaction rate versus equilibrium.