9.1 reversible reactions

9.1 Reversible Reactions

6.3 Reversible Reactions and Equilibrium

  • Definition: Some chemical reactions are reversible.

  • Equilibrium in Closed Systems:

    • Indicated with a specific symbol.

    • Achieved when:

      • The rate of the forward reaction equals the rate of the reverse reaction.

      • Concentrations of reactants and products remain constant.

  • Changing Conditions:

    • Can influence the direction of a reversible reaction.

Factors Affecting Equilibrium Position

Changing Temperature

  • Effects of Heat on Hydrated Compounds: Examine how temperature changes affect reactions involving hydrates.

Changing Pressure

  • Pressure changes impact gas reactions; it can shift the position of equilibrium.

Changing Concentration

  • Altering concentrations of reactants or products shifts equilibrium towards the side that reduces the change.

Catalysts

  • Catalysts do not affect the position of equilibrium but speed up reaction rates, allowing equilibrium to be reached faster.

Key Chemical Reactions and Conditions

Contact Process (Sulfuric Acid Manufacture)

  • Symbol Equation: N2(g) + 3H2(g) ⇌ 2NH3(g)

  • Typical Conditions:

    • Temperature: 450°C

    • Pressure: 200 kPa (2 atm)

    • Catalyst: Iron (Fe)

Sulfur Dioxide to Sulfur Trioxide

  • Symbol Equation: 2SO2(g) + O2(g) ⇌ 2SO3(g)

  • Typical Conditions:

    • Temperature: 450°C

    • Pressure: 20000 kPa (200 atm)

    • Catalyst: Vanadium(V) oxide

Reversible Reactions: Basics

  • Definition: Reversible reactions can proceed in both forward and backward directions depending on conditions.

  • Products can react or decompose back into reactants.

Reversible Hydration of Salts

Copper(II) Sulfate

  • Hydrated Form: Copper(II) sulfate pentahydrate (CuSO4•5H2O)

    • Appears as blue crystals.

    • Reaction: Heat imparted transforms it to anhydrous copper(II) sulfate (CuSO4).

    • Reaction is endothermic.

  • Anhydrous Form: Usually white, converting back to hydrated form upon water addition, which is highly exothermic.

Cobalt(II) Chloride

  • Hydrated Form: Cobalt(II) chloride hexahydrate (CoCl2•6H2O)

    • Appears pink.

    • Heating removes water, forming anhydrous cobalt(II) chloride (CoCl2) which appears blue.

  • Reactions: Both hydration and dehydration are reversible processes.

Chemical Tests for Presence of Water

  • Copper(II) sulfate: Changes from white (anhydrous) to blue (hydrated).

  • Cobalt(II) chloride: Changes from blue (anhydrous) to pink (hydrated).

Dynamic Equilibrium

  • Defined as the state in which:

    • Rate of forward reaction = Rate of reverse reaction

    • Concentrations of reactants and products remain constant.

  • Requires a closed system for equilibrium to be established.

Le Chatelier's Principle

  • States that any change in equilibrium conditions will shift the reaction to minimize the impact of that change.

  • Effect of Conditions:

    • Temperature

    • Pressure

    • Concentration

Effects of Changing Conditions on Equilibrium

Effect of Temperature

  • Increase favoring endothermic reactions, decrease favoring exothermic reactions.

Effect of Pressure

  • Affects gaseous reactions; increased pressure favors side with fewer gas molecules.

Effect of Concentration

  • Increase in reactant concentration shifts to produce more products.

  • Increase in product concentration shifts to produce more reactants.

Summary of Equilibrium Responses

  1. Temperature:

    • Increase: Shift to endothermic direction.

    • Decrease: Shift to exothermic direction.

  2. Pressure:

    • Increase: Shift toward fewer gas molecules.

    • Decrease: Shift toward more gas molecules.

  3. Concentration:

    • Increase of a substance: Shift to reduce that substance.

    • Decrease of a substance: Shift to increase that substance.

  4. Catalysts: Do not affect position, but speed up reaching equilibrium.

Exam Tips

  • When conditions change at equilibrium, the system responds by shifting in the opposite direction to counteract the change.