08.00 Chapter Summary

1. Rates of Reaction

1.1 Factors Affecting Reaction Rates

Concentration of Reactants:

• Effect: Higher concentration increases the number of collisions, speeding up the reaction.
• Example: Doubling the concentration of hydrogen peroxide in its decomposition reaction doubles the rate.

Pressure (for Gaseous Reactions):

• Effect: Increased pressure raises the concentration of gas molecules, leading to more frequent collisions.
• Example: Increasing the pressure in the Haber process (N₂ + 3H₂ ⇌ 2NH₃) shifts equilibrium to produce more ammonia.

Surface Area of Solid Reactants:

• Effect: Greater surface area (e.g., powdered vs. lump solids) exposes more particles to reactants, increasing the reaction rate.
• Example: Powdered magnesium reacts faster with hydrochloric acid than magnesium ribbon.

Temperature:

• Effect: Higher temperatures increase particle kinetic energy, resulting in more frequent and energetic collisions, thus speeding up the reaction.
• Example: Heating hydrochloric acid increases the rate at which magnesium reacts with it.

Presence of a Catalyst:

• Effect: Catalysts provide an alternative reaction pathway with lower activation energy, increasing the rate without being consumed.
• Example: Iron catalyst in the Haber process accelerates ammonia production.

1.2 Collision Theory

• Principles:
  • Necessary Collisions: Reactant particles must collide to react.
  • Sufficient Energy: Collisions must have enough energy (≥ activation energy) to break bonds.
  • Proper Orientation: Particles must collide in the correct orientation for a reaction to occur.
• Types of Collisions:
  • Successful Collisions: Lead to product formation.
  • Unsuccessful Collisions: Particles bounce off without reacting.

1.3 Measuring Reaction Rates

Methods:

• Mass Loss: Monitoring mass loss if a gas is produced (e.g., magnesium reacting with hydrochloric acid).

• Gas Volume: Collecting and measuring the volume of gas produced over time (e.g., decomposition of hydrogen peroxide).

• Color Change: Timing how long a color change takes to occur (e.g., the disappearing cross experiment with sodium thiosulfate).

• Graphing Data:
  • Rate vs. Time: Initial rate is steepest, decreasing as reactants are consumed.
  • Tangents: Used to determine instantaneous rates at specific points.

1.4 Evaluating Methods of Measuring Reaction Rates

• Gas Syringe:
  • Advantages: Accurate, suitable for all gas-producing reactions.
  • Disadvantages: Fragile, expensive, can stick, limited gas volume capacity.

• Inverted Measuring Cylinder:
  • Advantages: Simple setup, uses common lab equipment.
  • Disadvantages: Gas can be lost if not properly sealed, difficult to read upside-down.

• Mass Measurement:
  • Advantages: Easy to set up, uses common equipment.
  • Disadvantages: Not suitable for low-mass gases like hydrogen.

• Disappearing Cross Experiment:
  • Advantages: Simple, no specialized equipment needed.
  • Disadvantages: Subjective timing, potential for equipment contamination.

2. Physical vs. Chemical Changes

• Indicators of Chemical Changes:
  • Formation of a precipitate.
  • Evolution of gas (effervescence).
  • Color change.
  • Temperature change (exothermic/endothermic).
• Indicators of Physical Changes:
  • Change in state or form without new substances.

2.2 Factors Affecting Reaction Rates

• Remember the Five Factors:
  • Concentration
  • Pressure (gases)
  • Surface Area
  • Temperature
  • Catalysts
• Use the Collision Theory Framework: Focus on collision frequency and energy.

3. Equilibrium Shifts

• Use Le Chatelier’s Principle: Identify how changes affect the position of equilibrium.
• Predict Direction: Determine whether shifts favor reactants or products based on changes in concentration, pressure, or temperature.

4. Redox Reactions

• Use Oxidation Numbers: Assign and track changes to identify oxidized and reduced species.
• Mnemonic:OIL-RIG
  • Oxidation Is Loss (of electrons).
  • Reduction Is Gain (of electrons).

5. Industrial Processes

• Haber Process:
  • Key Conditions: 450°C, 200 atm, Iron catalyst.
  • Purpose: Synthesis of ammonia for fertilizers.

• Contact Process:
  • Key Conditions: 450°C, 2 atm, V₂O₅ catalyst.
  • Purpose: Production of sulfuric acid.

6. Worked Examples

6.1 Collision Theory Example

• Question: Explain how increasing temperature affects the rate of reaction.
• Answer:
  • Explanation: Higher temperature increases the kinetic energy of particles, resulting in more frequent and energetic collisions. This leads to a higher proportion of collisions exceeding the activation energy, thus increasing the reaction rate.

6.2 Redox Reaction Identification

• Question: In the reaction Fe + Cu²⁺ → Fe²⁺ + Cu, identify the oxidizing and reducing agents.
• Answer:
  • Fe: Loses electrons (oxidized) → Reducing Agent.
  • Cu²⁺: Gains electrons (reduced) → Oxidizing Agent.

6.3 Le Chatelier’s Principle Application

• Question: For the reaction 2NO₂ ⇌ N₂O₄, what happens when pressure is increased?
• Answer:
  • Effect: Increasing pressure favors the side with fewer gas molecules (N₂O₄).
  • Observation: Mixture becomes more colorless as more N₂O₄ is formed.

6.4 Redox Reaction Example

• Question: Identify which species is acting as the reducing agent in the reaction Fe + Br₂ → FeBr₂.
• Answer:
  • Fe: Loses electrons (oxidized) → Reducing Agent.
  • Br₂: Gains electrons (reduced) → Oxidizing Agent.

7. Key Concepts and Formulas

7.1 Chemical Equations for Reversible Reactions

• Symbol: ⇌ (double arrow with harpoons).
• Example: N₂ + 3H₂ ⇌ 2NH₃.

7.2 Oxidation Number Rules

• Element in pure form: 0.
• Monatomic ions: Equal to charge.
• Oxygen: Usually -2.
• Hydrogen: +1 with non-metals, -1 with metals.
• Fluorine: Always -1.
• Sum in compound: Equals overall charge.

7.3 Le Chatelier’s Principle Responses

• Increase in Reactant Concentration: Shift to the right (more products).
• Increase in Product Concentration: Shift to the left (more reactants).
• Increase in Pressure (gases): Shift toward fewer gas molecules.
• Increase in Temperature: Shift toward endothermic direction.