Unit 14 Equilibrium Notes

Unit 14 Equilibrium

Reversible Reactions

• Chemical reaction where products can regenerate original reactants.

• Expressed with a double arrow (\leftrightarrow).

• Example: 2NO2(g) \leftrightarrow N2O_4(g)

• Forward Reaction: 2NO2 \rightarrow N2O_4

• Reverse Reaction: N2O4 \rightarrow 2NO_2

• Reversibility:

  • Some reverse on their own.

  • Some reverse under specific conditions (temperature, pressure, catalyst).

  • Some do not reverse.

• Example of a non-reversible reaction: Single Replacement Reactions: 3CuCl2(aq) + 2Al(s) \rightarrow 3Cu(s) + 2AlCl3(aq)

Chemical Equilibrium

• State where product and reactant concentrations remain constant (not necessarily equal).

• Rate of forward and reverse reactions are equal.

• [ ] denotes concentration.

• Reaction rates are affected by concentration.

  • Increased reactant concentration increases the rate of the forward reaction.

• As the reaction proceeds:

  • Reactant concentration decreases.

  • Product concentration increases.

  • The forward reaction rate decreases.

• Reaching equilibrium does not mean the reaction has stopped.

• It indicates the rate of the forward and reverse reactions are equal.

• Dynamic equilibrium: Stays constant over time.

Changes As a Reaction Approaches Equilibrium

• Concentrations of reactants and products change until they reach a constant state.

• Concentrations are constant, but not necessarily equal, at equilibrium.

Reaction Rates & Equilibrium: Collision Theory

• Theoretical Model.

• Reactions occur when particles collide.

• Successful collision requires:

  • Sufficient energy.

  • Correct orientation.

• Effective collisions lead to product formation.

• Ineffective collisions do not lead to product formation.

• Effective collisions:

  • Enough energy.

  • Correct orientation to form products.

• Ineffective collisions revert to original reactants.

• Successful collision:

  • Bond breaking (endothermic).

  • Bond forming (exothermic).

• Activation Energy (E_a):

  • q WAr e a áMinimum energy needed for an effective collision.

• Transitional structure from successful collision:

  • Activated complex.

  • Unstable and short-lived.

  • Neither reactant nor product.

• Catalyst:

  • Increases reaction rate.

  • Provides mechanism with lower activation energy.

• Homogeneous system: One phase (e.g., g \rightarrow g \rightarrow g).

• Heterogeneous system: More than one phase (e.g., s \rightarrow l \rightarrow g).

Energy Diagrams

• Shows energy changes during a reaction.

• Endothermic:

  • Reactants to Products

• Exothermic:

  • Reactants to Products

• Forward Reaction

• Reverse Reaction

• Energy of Reactants

• Activated Complex

• Energy of Products

• Reaction Pathway

• E_a forward reaction (80 kJ)

• E_a reverse reaction (140 kJ)

• \Delta H = -60 kJ (for.)

• \Delta H = +60 kJ (rev.)

• Added Catalyst

Factors Affecting Reaction Rates

• Chemical Kinetics: Concerned with reaction rate.

• Reaction rate:

  • Measure of the speed of a chemical reaction.

  • Determined by measuring the change in concentration of reactants and products over time.

  • Determined experimentally.

• Rate-influencing factors:

  • Affect reaction rate by altering:

    • Frequency.

    • Orientation.

    • Energy of particle collisions.

Factors Affecting Reaction Rates - Collision Theory

• Nature of Reactants:

  • Structure: Complexity of bonds broken and formed, and orientation.

  • State: Homogeneous systems (reactants and products in the same state) vs. Heterogeneous systems (reactants and products in different states).

    • Homogeneous systems usually react faster than heterogeneous systems.

• Temperature:

  • (Average kinetic energy!)

  • Rule of thumb: Every 10°C increase in temperature doubles the reaction rate.

• Concentration:

  • Increased concentration increases the number of collisions.

  • Aqueous solutions: Can change concentration or molarity.

  • Gases: Can change pressure.

  • Solids and pure liquids (like water): Cannot change concentration.

• Surface Area:

  • Increased surface area increases the frequency of collisions.

  • Especially true for heterogeneous systems (s \rightarrow l \rightarrow g).

• Catalyst:

  • Increases reaction rate without being consumed.

  • Speeds up reaction by lowering the activation energy (E_a).

• Inhibitors:

  • Decrease reaction rate.

  • Take the place of a reactant and stop the reaction.

  • (Opposite of a catalyst.)

Energy Diagram Involving a Catalyst

• Uncatalyzed pathway

• Catalyzed pathway

• Activation energy of catalyzed pathway

• Reactants

• Products

• Progress of Reaction

Equilibrium Over Time

• Initial state: Only reactants (H2 + N2).

• Forward reaction: 3H2 + N2 \rightarrow 2NH_3 (synthesis, slows down).

• Reverse reaction: 2NH3 \rightarrow 3H2 + N_2 (decomposition, speeds up).

• Concentration changes over time:

  • Reactants (H2 + N2) decrease.

  • Product (NH_3) increases.

• Equilibrium is reached when the concentrations of reactants and products are constant.

• At equilibrium, the concentrations of reactants (H2 + N2) are greater than the concentration of product (NH_3).

• The reverse reaction is favored.

Equilibrium Systems and Stress – Le Chatlier’s Principle

• Le Chatelier's Principle: If a stress is applied to a system at equilibrium, the equilibrium will shift to relieve the stress.

• Stresses include changes in concentration, pressure, and temperature.

• Pure liquids and solids are not affected by changes in equilibrium.

• Le Chatelier's Principle is used to manipulate reversible reactions to maximize product formation.

The Haber Process

• Application of Le Chatelier's Principle.

• Used to produce ammonia (NH_3).

• Important for:

  • Nitrogen Containing Explosives

  • Household cleaners.

  • Fertilizers: Increased food production.

Factors Determining Reaction Direction

• Determines whether a reaction favors reactants (lies to the left) or products (lies to the right).

• Factors:

  • Changes in Concentration: Adding/Removing Reactants and Products.

  • Changes in Pressure.

  • Changes in Temperature.

Changes in Concentration - (Reactants and Products)

• Adding a substance: Shifts the system AWAY from that substance, making more of the opposite side.

• Removing a substance: Shifts the system toward REPLACING that substance.

• Example: N2O4(g) \leftrightarrow 2NO_2(g)

  • \uparrow N2O4(g) shift right making more NO_2

  • \uparrow NO2(g) shift left making more N2O_4

  • \downarrow N2O4(g) shift left making more N2O4

  • \downarrow NO2(g) shift right making more NO2

Changes in Pressure

• Avogadro's Law: Equal volumes of gases at the same temperature and pressure have the same number of molecules. (1 mole of any gas at STP = 22.4L = 6.02 x10^{23} particles)

• Boyle's Law: An increase in pressure means a decrease in volume. So if the pressure is increased on a system at equilibrium, the side which occupies the lower volume will be favored. (inverse relationship)

• Increased pressure: Shifts the reaction to the side with fewer moles of gas.

• Examples:

  • 2NO2(g) \leftrightarrow N2O4(g) ; \uparrow pressure shift to right making more N2O_4

  • 2H2O2(g) \leftrightarrow 2H2O(g) + O2(g); \uparrow pressure shift to left making more H2O2

  • H2(g) + Cl2(g) \leftrightarrow 2HCl(g); \uparrow pressure shift to neither! making more same # of moles each side

• Pressure changes only affect gases!

Changes in Temperature

• Exothermic Reactions (release heat):

  • Increased temperature: Favors the reaction that absorbs heat.

  • Decreased temperature: Favors the reaction that releases heat.

  • Example: H2(g) + I2(g) \leftrightarrow 2HI(g) + Heat

    • Increasing the temperature will produce a higher yield of H2 + I2.

    • Lowering the temperature will produce a higher yield of HI.

• Endothermic Reactions (absorb heat):

  • Increased temperature: Favors the reaction that absorbs heat.

  • Decreased temperature: Favors the reaction that absorbs heat.

  • Example: HEAT + NH4Cl(s) \leftrightarrow NH3(g) + HCl(g)

    • Increasing the temperature will produce a higher yield of NH_3 + HCl.

    • Lowering the temperature will produce a higher yield of NH_4Cl.

Conclusion

• Le Chatelier's Principle: A system at equilibrium, when subjected to a stress, will temporarily adjust itself to relieve the stress.

• Shift to the right or left (increased forward or reverse reaction) is temporary, and a new equilibrium will be reestablished.

Practice

1. 2SO2(g) + O2(g) \leftrightarrow 2SO_3(g) + heat

   What conditions of temperature and pressure favor high equilibrium concentrations of SO_3?

   (high) pressure; (low) temperature

2. 3H2(g) + N2(g) \leftrightarrow 2NH_3(g) + heat

   The commercial production of ammonia uses the Haber Process which is expressed by the above equation.

   What condition of temperature and pressure will provide a maximum yield of NH_3?

   (high) pressure: (low) temperature

3. 4HCl(g) + O2(g) \leftrightarrow 2H2O(g) + 2Cl_2(g) + heat

   Increasing the temperature of the reaction will (decrease) the forward reaction.

   Decreasing the pressure on the system will (decrease) the forward reaction.