Chemical Reactions: Reversible Reactions and Equilibrium

Reversible Reactions

  • A reversible reaction can proceed in both directions.
  • The products of the reaction can be turned back into the original reactants.
  • Symbol for a reversible reaction: ⇌
  • Example: A + B ⇌ C + D, where A and B react to produce C and D, and C and D can react to produce A and B.

Investigating Reversible Reactions

  • Copper(II) Sulfate:
    • Hydrated copper sulfate (blue crystals) can be converted to anhydrous copper sulfate (white) and water, and vice versa.
    • CuSO₄·5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(g)
    • Forward reaction: dehydration (endothermic).
    • Backward reaction: hydration (exothermic).
    • Heating blue copper sulfate drives off water, forming white anhydrous copper sulfate.
    • Adding water to white anhydrous copper sulfate reforms blue hydrated copper sulfate.
  • Ammonium Chloride:
    • Ammonium chloride decomposes into ammonia and hydrogen chloride upon heating, and the reverse reaction occurs upon cooling.
    • NH₄Cl (s) ⇌ NH₃(g) + HCl (g)
    • Forward reaction: decomposition (endothermic).
    • Backward reaction: addition (exothermic).
    • Heating solid ammonium chloride splits it into ammonia gas and hydrogen chloride gas.
    • Cooling the gases recombines them to form solid ammonium chloride.
  • Testing for Water:
    • Cobalt(II) chloride paper turns from blue to pink upon the addition of water.
    • Anhydrous cobalt(II) chloride + water ⇌ hydrated cobalt(II) chloride

Equilibrium

  • Reversible reaction: Reactants form products, and products react to reform reactants.
    • A + B ⇌ C + D
    • Forward and backward reactions can occur separately or simultaneously.
  • Equilibrium is a state of balance between opposing forces or processes.
    • Mechanical equilibrium: a seesaw balanced.
    • Thermal equilibrium: two objects at different temperatures eventually reaching the same temperature.
    • Chemical equilibrium: Forward reaction rate equals the backward reaction rate.

Introduction to Equilibrium

  • At equilibrium, the amounts of reactants and products remain constant.
  • Dynamic equilibrium: The forward and backward reactions continue to occur at the molecular level, but there is no overall change in the amounts of reactants and products, maintaining a state of balance.
  • Model for dynamic equilibrium: A busy road where cars enter and leave at the same rate, keeping the number of cars constant.

Closed Systems

  • A closed system is one where no matter can enter or leave, but energy can be transferred.
  • Dynamic equilibrium can only be achieved in a closed system.
  • Continuous addition or removal of substances disrupts the balance between the forward and backward reaction rates, preventing the system from reaching equilibrium.

Le Chatelier's Principle and Concentration

  • When a reaction has reached dynamic equilibrium, the concentrations of reactants and products remain constant, but the reactions have not stopped.
  • French chemist Henry Louis Le Chatelier's principle states that if a dynamic equilibrium is disturbed by changing the conditions, the equilibrium position shifts to counteract the change.
  • The equilibrium position shift refers to the concentrations of reactants and products at the new equilibrium relative to the original concentrations.
  • Increasing Reactant Concentration: Equilibrium shifts to the right to produce more products and reduce the concentration of added reactants.
  • Decreasing Reactant Concentration: Equilibrium shifts to the left to produce more reactants and reduce the concentration of products.
  • Example: Haber process for ammonia production
    • N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
    • To produce more ammonia, increase N₂ or H₂ concentrations or decrease NH₃ concentration.
  • Increasing the concentration of reactants increases the reaction rate of the forward reaction due to more frequent collisions.

Pressure, Temperature, and Catalysts

  • One mole of any gas occupies the same volume under the same temperature and pressure conditions.
    • 1 \, mole = 6.02 × 10^{23} things
  • Changing the number of moles of gas in a fixed volume will change the pressure.
  • Adding another mole of gas doubles the pressure.
  • Changing pressure affects equilibria where each side has an unequal number of moles of gas.
  • CO(g) + 2H₂(g) ⇌ CH₃OH(g)
    • Reactants: three moles of gaseous reactants.
    • Products: one mole of gaseous products.
    • Reactants have a pressure 3x greater than the products.
  • Le Chatelier's Principle: Equilibrium position shifts to minimize the effect of the pressure change.
    • Increasing pressure favors the side with fewer gas molecules.
    • Decreasing pressure favors the side with more gas molecules.
  • Example: Haber process
    • High pressure favors the production of ammonia (fewer gas molecules on the right side).

Temperature

  • Changing temperature affects equilibrium position because one direction of a reversible reaction is exothermic, and the other is endothermic.
    • N₂O₄(g) ⇌ 2NO₂(g)
      • Exothermic
      • Endothermic
    • O₂ + 2H₂ ⇌ 2H₂O
      • Exothermic
      • Endothermic
  • Le Chatelier's Principle: Equilibrium position shifts to counteract the temperature change.
    • Increasing temperature favors the endothermic reaction (absorbs heat, decreasing the temperature of the surroundings).
    • Decreasing temperature favors the exothermic reaction (releases heat, increasing the temperature of the surroundings).
  • Example: Contact process for sulfur trioxide production.
    • 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)
    • The forward reaction is exothermic, so lower temperatures favor the production of the product (SO₃).

Catalysts

  • Catalysts speed up the rate at which equilibrium is achieved; both forward and backward reaction rates speed up equally.
  • Catalysts do not change the equilibrium position but are crucial for increasing production efficiency.
    • N₂(g) + 3H₂(g) ⇌ 2NH₃(g), catalyst: iron
    • 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), catalyst: vanadium(V) oxide
  • Changing pressure and temperature also affect the reaction rate.
    • Increasing temperature increases the energy of particles, leading to more frequent, successful collisions.
    • Increasing pressure of gases forces gas particles closer together, leading to more frequent, successful collisions.

Summary of Factors Affecting Equilibrium Position

  • Temperature:
    • Increasing: Moves in the endothermic direction.
    • Decreasing: Moves in the exothermic direction.
  • Pressure:
    • Increasing: Shifts equilibrium to produce fewer gas molecules (only affects gaseous reactions).
    • Decreasing: Shifts equilibrium to produce more gas molecules (only affects gaseous reactions).
  • Concentration:
    • Increasing a substance: Equilibrium moves to produce less of that substance.
    • Decreasing a substance: Equilibrium moves to produce more of that substance.
  • Catalyst:
    • Does not affect equilibrium position but reaction reaches equilibrium faster.
  • Catalysts, pressure, temperature, and concentration can affect the equilibrium position of a reversible reaction.
  • If the conditions of a reversible reaction are changed, the equilibrium will move to counteract that change.

Haber Process

  • Symbol equation for ammonia production: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
  • Sources:
    • Nitrogen: air
    • Hydrogen: methane
  • Typical conditions: 450 °C, 20,000 kPa / 200 atm, iron catalyst
  • Developed to make ammonia from atmospheric nitrogen during World War I due to blockades preventing access to natural fertilizers.
  • Fritz Haber developed the process.
  • Ammonia produced is used for fertilizer, sustaining one-third of the Earth’s population.

Introduction to the Haber process

  • Feedstock: Methane (CH₄) is the most common source of hydrogen.
  • Natural gas (mostly methane) is used because it is abundant and cost-effective.
  • Hydrogen production:
    • CH₄(g) + H₂O(g) ⇌ CO(g) + 3H₂(g)
    • CO(g) + H₂O(g) ⇌ CO₂(g) + H₂(g)
  • Energy-intensive process produces ~2% of global carbon dioxide emissions.

The Haber Process Reaction

Reaction vessel setup:

  • Hydrogen and nitrogen mixed in 3:1 ratio at 450°C and 200 atm
  • NH3 formed as gas cools in the condenser, liquefies and is removed.
  • Unused N2 and H2 are recycled
  • Cooling, condensing, and regular collection pushes the equilibrium towards more ammonia production, maximizing yield.

Practical Applications and Environmental Impact

  • Ammonia is crucial for fertilizer production, supporting global agriculture.
  • Also used in the manufacture of: explosives, pharmaceuticals, textiles, cleaning products.
  • Fertilizers 75%
  • Nitric acid 10%
  • Others 10%
  • Nylon 5%

Contact Process

  • Symbol equation: 2𝑆𝑂2(𝑔) + 𝑂2(𝑔) ⇌ 2𝑆𝑂3(𝑔)
  • Sources:
    • Sulfur dioxide: burning sulfur or roasting sulfide ores
    • Oxygen: air
  • Typical conditions: 450 °C, 200 kPa / 2 atm, vanadium(V) oxide catalyst
  • An industrial method for producing sulfuric acid

Introduction to the Contact process

  • Raw materials:
    • Sulfur: sourced from Earth's crust or as a by-product of refining fossil fuels.
    • Oxygen: sourced from the air.
    • Water: essential for the final production stages.
  • Series of reactions:
    • Step 1: Making sulfur dioxide
      • S (s) + O₂(g) → SO₂(g)
      • Cu₂S (s) + O₂(g) → 2Cu (l) + SO₂(g)
  • Step 2:
    • Reversible reaction between sulfur dioxide and oxygen to make sulfur trioxide.
    • 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)
    • Conditions: 450°C, 2 atm, V₂O₅ catalyst
  • Step 3:
    • Converting sulfur trioxide into sulfuric acid.
    • SO₃(g) ​+ H₂​SO₄(aq) → H₂​S₂O₇​ (l)
    • H₂S₂O₇(l) + H₂O(l) → 2H₂SO₄(aq)

Practical Applications and Hazards

  • Sulfuric acid is essential in the production of fertilizers, such as ammonium sulfate.
  • Other uses include: car battery acid, and feedstock.
  • Paints, Fertilizers, Synthetic Fibres, Drugs, Car batteries, Dyes, Detergents, and Plastics.

Balancing Safety, Cost, and Efficiency

  • Catalysts (V₂O₅) are used to speed up the reaction rate without affecting the equilibrium position.
  • The removal of sulfur trioxide shifts the equilibrium position, increasing the product yield.

Factors Affecting Industrial Equilibria

  • Industrial equilibria play a crucial role in modern manufacturing processes.
    • Ammonia is a multi-million-pound industry, produced using the Haber process.
    • Sulfuric acid, produced through the Contact process, is one of the highest-volume industrial chemicals produced worldwide.
  • Equilibrium position: relative concentrations of reactants and products in a reversible reaction at equilibrium.
  • Factors affecting equilibrium position: temperature, pressure, concentration.
  • Conditions for the Haber process: 450°C, 200 atm, iron catalyst, ammonia is condensed and removed.
  • Conditions for the Contact process: 450°C, 1 – 2 atm, vanadium oxide (V₂O₅) catalyst, removal of SO₃.
  • Industrial equilibria involve a trade-off between production rate and equilibrium position.

Factors Affecting Industrial Equilibria

  • The Haber process involves a reversible reaction at dynamic equilibrium. N₂ (g) + 3H₂ (g) ⇌ 2NH₃ (g)
  • The forward reaction is exothermic, and the reverse (backward) reaction is endothermic.
  • increasing temperature will speed up reaction rate and decrease yield of product by shifting the equilibrium position.

Balancing Industry Conditions for Equilibria

  • To minimise safety risks, reduce costs To avoid excessive reaction rates, to shift the equilibrium position favourably, To prevent catalysts from being deactivated.
  • Industrial reactions are also influenced by: availability of raw materials, cost of raw materials, energy supplies, and cost.
  • Haber process: nitrogen from air and hydrogen from sources like natural gas. Energy-intensive due to high pressure and temperature.

Equilibrium Balance

  • Industrial equilibria often involve a trade-off between the rate of production and the position of equilibrium.
  • Temperature, pressure, surface area, concentration, and catalysts all impact reaction rates and equilibrium positions, sometimes in opposing ways.
  • Maintaining very high pressures is costly and increases the risk of equipment failure.
  • High pressures and temperatures improve reaction rates but must be balanced against safety and economic considerations.
  • Low temperatures tend to slow down reactions, impacting production rates, but can shift equilibrium to favor the desired product in some exothermic reactions.