Chemistry -Entrance Exam ( Universidad de Navarra) - unit 5 (1)

Unit 4: Kinetics and Equilibrium

1. Chemical Dynamics

• Definition: Study of rates and mechanisms of chemical reactions.

• Focus: Energy transfer among molecules during collisions in different phases.

• Reaction Dynamics: Emphasizes individual reactive collisions to understand fundamental chemical reactions.

• Key Equation:

  • A + B → AB

  • Concentrations of A and B decrease over time, while AB increases.

• Rate Calculation:

  • Rate = Moles of product formed / Time elapsed

  • Rate = Moles of reactants transformed / Time elapsed

2. Collisions in Reactions

• Requirement: Collisions between reactant molecules are necessary for reactions.

• Reactive Collisions: Successful collisions that lead to a reaction.

• Reaction Rate: Change in concentration (reactants/products) over time, measured in mol dm⁻³ s⁻¹.

• Factors Influencing Reaction Rate:

  • Concentration of reactants

  • Surface area of solid reactants

  • Temperature

  • Presence of a catalyst

• Graphical Representation:

  • Concentration of A decreases; concentration of B increases over time.

  • Steeper gradient indicates faster reaction rate.

3. Effect of Concentration and Surface Area

• Concentration: Higher concentration leads to more collisions and increased reaction rates.

• Surface Area: Greater surface area of solids increases exposure and reaction frequency.

4. Effect of Temperature

• Kinetic Energy: Average kinetic energy of particles rises with temperature, leading to faster movement and more collisions.

• Boltzmann Distribution: Temperature increase shifts the distribution curve to the right with a lower peak.

5. Collision Theory

• Concept: Explains how collisions affect reaction rates.

• Conditions for Reaction:

  1. Sufficient energy

  2. Proper orientation

• Temperature's Role: Increased temperature boosts kinetic energy, leading to more frequent effective collisions.

• Catalysts: Provide alternative pathways with lower activation energy, enhancing reaction rates.

• Rate-Determining Step: Slowest step in a sequence of reactions that dominates the overall reaction rate.

6. Effective Collisions

• Definition: Collisions with sufficient energy and proper orientation.

• Factors Increasing Effective Collisions:

  • Higher temperature (increases energy)

  • Higher concentration (increases collision frequency)

  • Use of catalysts (lowers activation energy)

7. Catalysts

• Definition: Substances that increase reaction rates without being consumed.

• Mechanism: Provide alternative pathways for reactions, changing the reaction mechanism.

• Effect on Equilibrium: Catalysts speed up the approach to equilibrium without affecting composition.

8. Dynamic Equilibrium Characteristics

• Conditions: Achieved in closed systems with reversible reactions.

• Properties:

  • Both reactants and products present.

  • Forward and backward reactions occur at equal rates.

  • Concentrations remain constant unless disturbed.

9. Static vs Dynamic Equilibrium

• Static: No changes; forward and backward reactions halt.

• Dynamic: Continuous change while maintaining overall composition.

10. Equilibrium Constant (Kc)

• Expression: Links Kc to the concentrations of reactants/products in a reversible reaction.

  • Kc = [C]^c [D]^d / [A]^a [B]^b

• Solids: Not included in Kc expressions.

11. Mole Fraction and Partial Pressure

• Partial Pressure: The pressure a gas would exert if alone in a container.

• Mole Fraction: Ratio of moles of one gas to total moles present.

• Total Pressure: Sum of partial pressures from each gas.

12. Precipitation Reactions

• Definition: Involves forming an insoluble solid from two aqueous solutions.

• Role of Solubility Product Constant (Ksp): Determines if precipitation occurs based on ion product (Q) vs Ksp.

• Double Displacement Reactions: Cations/anions switch to form products, one being insoluble precipitate.

13. Chemical Equilibrium Factors

• Le-Chatelier’s Principle: Predicted shifts in equilibrium based on changes in concentration, temperature, and pressure.

• Changes:

  • Concentration: Shifts towards more products.

  • Temperature: High temperature favors endothermic reactions.

  • Pressure: Increased pressure shifts equilibrium toward fewer gas moles.

  • Catalysts: Do not affect equilibrium position but speed up reactions.

14. Haber-Bosch Process

• Purpose: Converts nitrogen and hydrogen into ammonia (NH₃).

• Historical Significance: Developed by Fritz Haber and Carl Bosch, pivotal for fertilizers.

• Reaction: N₂(g) + 3H₂(g) ⇌ 2NH₃(g); ΔH = −92.4 kJ/mol.

• Conditions: High pressure (200 atm), high temperature (450°C), iron catalyst.

• Process Efficiency:

  • Continuous removal of NH₃ ensures reaction balance.

  • Raw materials: Nitrogen from atmosphere, hydrogen from methane and steam reactions.

15. Thermodynamics of Haber Process

• Spontaneity: The reaction is spontaneous with a decrease in enthalpy and entropy.

• Optimal Conditions: Use of high pressure to maximize production and steady reaction rate under controlled thermodynamics.