CHEMICAL EQUILIBRIUM

• Chemical Equilibrium


  • Reversible reactions
  • Reaction rates of the forward and reverse reaction
• Writing expressions for the reaction quotient/equilibrium constants

  1. Write expression

  2. Explain significance of equilibrium constant value

  3. Calculate equilibrium constant

  4. Calculate pressure or concentration of reactants or products in an equilibrium mixture
• Le Chatelier's principle


  Apply it qualitatively to describe the effect of changes in pressure, concentration and temperature on a system at equilibrium reaction
• The Concept of Equilibrium

  1. Describe decomposition of N2O4 to NO2

  2. Explain when chemical equilibrium is reached

  3. Explain dynamic equilibrium
• At the beginning of the reaction, there is no NO2 so the reverse reaction (2NO2(g) ⇌ N2O4(g)) does not occur
• At equilibrium, as much N2O4 reacts to form NO2 as NO2 reacts to re-form N2O4
• At equilibrium, kf[A] = kr[B]
• Equilibrium constant


  Ratio of concentrations of products to reactants at equilibrium
• Calculating equilibrium constant


  1. Write equilibrium constant expression

  2. Consider coefficients of reactants and products

  3. Ignore concentrations of pure solids and liquids

  4. Kp has no units
• Equilibrium constant


  • Larger K means more products at equilibrium
  • Smaller K means more reactants at equilibrium
• Direction of chemical equation

  Determines whether equilibrium lies to the left or right
• Equilibrium constant for reverse direction is inverse of forward direction
• When reaction is multiplied by a number, equilibrium constant is raised to that power
• Equilibrium constant for a reaction which is the sum of other reactions is the product of the equilibrium constants for the individual reactions
• Homogeneous equilibrium


  All reactants and products in one phase
• Heterogeneous equilibrium


  One or more reactants or products in a different phase
• Concentrations of pure solids and liquids are constant and ignored in equilibrium constant expressions
• Amount of CO2 formed does not depend greatly on amounts of CaO and CaCO3 present
• Calculating equilibrium concentrations


  1. Tabulate initial and equilibrium concentrations

  2. Calculate change in concentration

  3. Use stoichiometry to calculate changes in all species

  4. Deduce equilibrium concentrations
• Predicting direction of reaction


  Depends on equilibrium constant value
• Heterogeneous Equilibria


  The concentration of a solid or pure liquid is its density divided by molar mass. Neither density nor molar mass is a variable, the concentrations of solids and pure liquids are constant.
• Decomposition of CaCO3


  1. We ignore the concentrations of pure liquids and pure solids in equilibrium constant expressions

  2. The amount of CO2 formed will not depend greatly on the amounts of CaO and CaCO3 present
• Calculating Equilibrium Constants


  1. Tabulate initial and equilibrium concentrations (or partial pressures) given

  2. If an initial and equilibrium concentration is given for a species, calculate the change in concentration

  3. Use stoichiometry on the change in concentration line only to calculate the changes in concentration of all species

  4. Deduce the equilibrium concentrations of all species

  5. Usually, the initial concentration of products is zero
• Reaction Quotient (Q)


  We define Q, the reaction quotient, for a general reaction as Q = K only at equilibrium
• If Q > K


  The reverse reaction must occur to reach equilibrium (i.e., products are consumed, reactants are formed, the numerator in the equilibrium constant expression decreases and Q decreases until it equals K)
• If Q < K


  The forward reaction must occur to reach equilibrium
• Calculating Equilibrium Constants


  1. We do not have a number for the change in concentration line

  2. Therefore, we need to assume that x mol/L of a species is produced (or used)

  3. The equilibrium concentrations are given as algebraic expressions
• Le Châtelier's Principle


  If a system at equilibrium is disturbed, the system will move in such a way as to counteract the disturbance
• If H2 is added while the system is at equilibrium

  The system must consume the H2 and produce products until a new equilibrium is established. So, [H2] and [N2] will decrease and [NH3] increases.
• Adding a reactant or product


  Shifts the equilibrium away from the increase
• Removing a reactant or product

  Shifts the equilibrium towards the decrease
• Industrial preparation of ammonia


  1. N2 and H2 are pumped into a chamber

  2. The pre-heated gases are passed through a heating coil to the catalyst bed

  3. The catalyst bed is kept at 460 - 550 °C under high pressure

  4. The product gas stream (containing N2, H2 and NH3) is passed over a cooler to a refrigeration unit

  5. In the refrigeration unit, ammonia liquefies not N2 or H2

  6. The unreacted nitrogen and hydrogen are recycled with the new N2 and H2 feed gas
• As volume is decreased


  Pressure increases
• If pressure is increased


  The system will shift to counteract the increase by removing gases and decreasing pressure
• An increase in pressure


  Favors the direction that has fewer moles of gas
• In a reaction with the same number of product and reactant moles of gas

  Pressure has no effect
• Equilibrium constant


  Is temperature dependent
• For an endothermic reaction, ΔH > 0


  Heat can be considered as a reactant
• For an exothermic reaction, ΔH < 0


  Heat can be considered as a product
• Adding heat (i.e. heating the vessel)


  Favors away from the side where there is absorption of heat