Chemical Equilibrium

Unit Objectives  Define chemical equilibrium.  Explain the nature of the equilibrium constant.  Write chemical equilibrium expressions and carry out calculations involving them  Construct for any chemical equation the corresponding reaction quotient expression.  Clearly distinguish between Q and Keq.  Perform calculations required for a Q vs. Keq analysis.

Chemical Equilibrium  In a chemical equilibrium there are finite concentrations of reactants and products and these concentrations remain constant with time.  Equilibrium reactions always result in smaller amounts of products than the theoretical yield predicts  [reactants] and [products] at equilibrium provide a quantitative way of determining how successful a reaction has been.

Chemical Equilibrium

 Equilibria are dynamic (two opposing processes occur simultaneously) Product-favored (reaction tends to the right) products predominate over reactants Reactant-favored (reaction tends to the left) most of equilibrium mixture consists of reactants

Chemical Equilibrium  Though most reactions go to completion, chemical reactions are reversible (in theory) by changing the conditions of the reaction mixture.  Nano-scale rate of forward reaction exactly equals the rate of the reverse reaction CH 3 COOH (aq) + H 2 O (l) ↔ CH 3 COO - (aq) + H 3 O +

Chemical Equilibrium  Equilibrium is independent of direction of approach Whether you start with reactants or products the same equilibrium state is achieved  Catalysts do not affect equilibrium concentrations If a catalyst is present the same equilibrium state will be achieved, but more quickly

Equilibrium Constant and Reaction Quotient  At any given temperature, the quantity  is a constant, independent of the initial amounts of reactants and products, the volume of the container, or the total pressure. This constant is referred to as the equilibrium constant, K c

Equilibrium Constant and Reaction Quotient  For the balanced chemical equation aA + bB ⇌ cC + dD  Constant varies with temperature and coefficients of balanced equation

Equilibrium Constant and Reaction Quotient For the reverse reaction  Use the reciprocal rule: the equilibrium constants for the forward and reverse reactions are the reciprocals of each other

Equilibrium Constant and Reaction Quotient: K Expressions for Related Reactions  N 2 (g) + 3 H 2 (g) ⇌ 2 NH 3 (g) K c1 = 3.5 x 10 8 (at 25C)  1/2 N 2 (g) + 3/2 H 2 (g) ⇌ NH 3 (g) K c2 = ?  Is the value of K c2 the same as the value of K c1 ?

Equilibrium Constant and Reaction Quotient: K Expressions for Related Reactions  K c1 is the square of K c2 K c1 = (K c2 ) 2

Equilibrium Constant and Reaction Quotient: K Expressions for Related Reactions Whenever the stoichiometric coefficients of a balanced chemical equation are multiplied by some factor, the equilibrium constant for the new equation is the old equilibrium constant raised to the power of the multiplication factor

Equilibrium Constant and Reaction Quotient: K for Reactions that Combine Two or More Reactions N 2 (g) + O 2 (g) ⇌ 2NO (g)K c1 2NO (g) + O 2 (g) ⇌ 2NO 2 (g)K c2 N 2 (g) + 2O 2 (g) ⇌ 2NO 2 (g) K c1 x K c2  If two chemical equations can be summed to give a third, the equilibrium constant for the overall equation equals the product of the two equilibrium constants for the equations that were summed.

Equilibrium Constant and Reaction Quotient: Equilibrium Constants in Terms of Pressure Product pressures raised to powers of coefficients Reactant pressures raised to powers of coefficients

Equilibrium Constant and Reaction Quotient: Equilibrium Constants in Terms of Pressure  K p indicates equilibrium constant has been expressed in terms of partial pressures. Sometimes K c = K p, for most equilibria it does not.  So, it is useful to relate the two:

Equilibrium Constant and Reaction Quotient:Determining Equilibrium Constants 1.Write the balanced equation for the equilibrium reaction. From it derive the equilibrium constant expression 2.Set up an ICE table: a table containing initial concentration, change in concentration, and equilibrium concentration for each substance included in the equilibrium constant expression. Enter all known information into this table 3.Use x to represent the change in concentration of one substance. Use the stoichiometric coefficients in the balanced equilibrium equation to calculate the other changes in terms of x

Equilibrium Constant and Reaction Quotient:Determining Equilibrium Constants 4.From initial concentrations and the changes in concentrations, calculate the equilibrium concentrations in terms of x and enter them in the table 5.Use the simplest possible equation to solve for x. Then use x to calculate the unknown you were asked to find.

Equilibrium Constant and Reaction Quotient:Determining Reaction Quotient Applications of the Equilibrium Constant  K c, K p, K a, etc. provides qualitative and quantitative information with regard to the extent to which a reaction will occur  Knowing the equilibrium constant allows one to predict whether or not a reaction will proceed.

Equilibrium Constant and Reaction Quotient:Determining Reaction Quotient In general (qualitative)  If the equilibrium constant is very large, the forward reaction will proceed far to the right (system contains mostly products)  If the equilibrium constant is very small, no reaction proceeds in the forward direction (system consists almost entirely of reactants)  If the equilibrium constant is neither very large or very small product and reactant concentrations are about the same.

Equilibrium Constant and Reaction Quotient:Determining Reaction Quotient Direction of the Reaction (quantitative)  Use the reaction (concentration) quotient, Q to predict reaction direction  For the general equilibrium expression K c

Equilibrium Constant and Reaction Quotient:Determining Reaction Quotient  The original concentration quotient Q, may be used to determine the direction of the reaction

Equilibrium Constant and Reaction Quotient:Determining Reaction Quotient  Q rarely equals K c (reactions usually do not start at equilibrium) so the reaction will proceed in one direction or the other until the equilibrium concentrations are reached.  If Q<K c reaction proceeds to the right ([products] increases, [reactants] decreases), Q increases until it becomes equal to K c.  If Q>K c reaction proceeds in the reverse direction ([products] are too ‘high’ and [reactants] are too ‘low’ to meet equilibrium conditions; [products] decreasing, [reactants] increasing)  Use Q to predict the reaction direction then establish the equilibrium concentrations using the previously presented procedure (equilibrium expression, ICE table, etc)

Le Chatelier’s Principle: Effects of Change on an Equilibrium System  3 ways to disturb an equilibrium system  Add or remove gaseous reactants or products  Change the volume of the system (change the # of moles of gas present)  Change the temperature If a system at equilibrium is disturbed by some change, the system will shift as to partially counteract the effect of the change, if possible.

Le Chatelier’s Principle: Effects of Change on an Equilibrium System Add or Remove Gaseous Reactants or Products If we disturb a chemical system at equilibrium by adding a gaseous species (product or reactant), the reaction will proceed in such a direction as to consume part of the added species. Conversely, if we remove a gaseous species, the system will shift so as to restore part of that species.

Le Chatelier’s Principle: Effects of Change on an Equilibrium System N 2 O 4 (g) ⇌ 2NO 2 (g) ChangeEffect Adding N 2 O 4 Reaction will move to the right consuming some of the added N 2 O 4 Adding NO 2 Reaction will move to the left (reverse) using up part of the NO 2 added Removing N 2 O 4 Reaction occurs in the reverse direction to restore part of N 2 O 4 Removing NO 2 Forward reaction occurs restoring part of the NO 2 removed.

Le Chatelier’s Principle: Effects of Change on an Equilibrium System Changes in Volume When the volume of an equilibrium system is decreased, reaction takes place in the direction that decreases the total number of moles of gas. When the volume is increased, the reaction that increases the total number of moles of gas takes place.

Le Chatelier’s Principle: Effects of Change on an Equilibrium System SystemV IncreasesV Decreases N 2 O 4 (g) ⇌ 2NO 2 (g) →← N 2 (g) + 3H 2 (g) ⇌ 2NH 3 (g) ←→ N 2 (g) + O 2 (g) ⇌ 2NO(g) 00

Le Chatelier’s Principle: Effects of Change on an Equilibrium System Since changes in volume result from changes in applied pressure we could also say that: an increase in pressure shifts the position of the equilibrium in such a way as to decrease the number of moles of gas (the reverse is also true)

Le Chatelier’s Principle: Effects of Change on an Equilibrium System Changes in Temperature  An increase in temperature causes the endothermic reaction to occur (H = +)  A decrease in temperature causes the exothermic reaction to occur (H = -) If the forward reaction is endothermic, K c becomes larger as the temperature increases. If the forward reaction is exothermic, K c becomes smaller as the temperature increases.

Le Chatelier’s Principle: Effects of Change on an Equilibrium System A few more things…  Treat Heat of Reaction as a reactant (endothermic) or product (exothermic)  2SO 3 (g) ⇌ 2SO 2 (g) + O 2 (g) H = 197 kJ  Heat as a reactant T ↑, shift right, K ↑. Opposite for Heat as a product.

Le Chatelier’s Principle: Effects of Change on an Equilibrium System Increase in Pressure (another look at volume changes)  If pressure increases in an equilibrium system, the reaction will shift toward the side with fewer moles of gas.  Note: decreasing volume has the same effect (the only way to change P without changing T or n is to decrease V) For 2SO 3 (g) ⇌ 2SO 2 (g) + O 2 (g)H = 197 kJ  Shift left toward SO 3  if moles of gas are the same on both sides – no shift

Le Chatelier’s Principle: Effects of Change on an Equilibrium System  Adding solids or an inert gas (no change in volume) No shift in equilibrium. Neither are part of the equilibrium expression. [ ] of all components of equilibrium expression remain unchanged.  Inert gas Added at constant P (volume changes) Shift to side of equation w/more moles of gas. Partial pressure of all gases decreases as volume increases.