Chemical Equilibrium

Chemical Equilibrium A dynamic process.

Chemical Equilibrium When compounds react they eventually form a mixture of products and un-reacted reactants in a dynamic equilibrium. • Dynamic Equilibrium – consists of a forward reaction, in which substances react to give products; and a reverse reaction in which products react to give the original reactants. • Chemical Equilibrium – is the state reached by a reaction mixture when the rates of the forward and reverse reactions have become equal.

Chemical Equilibrium Equilibrium is a state where there are no observable changes. Chemical Equilibrium is achieved when: • The rates of the forward and reverse reactions are equal. (Opposing Reactions) • These opposing reactions occur simultaneously and the rate at which product is formed during the forward reaction is equal to the rate at which reactants are re-formed during the reverse reaction. • The concentrations of the reactants and products remain constant. • The quantities of reactants and products do NOT have to be equal but, they must remain constant to have equilibrium.

Chemical Equilibrium • Chemical Equilibrium occurs in a closed system (nothing can get in or out), therefore the reactions are continuous. • Double arrow = equilibrium

Types of Equilibrium • Physical Equilibrium – equilibrium of physical processes such as phases of matter or dissolving. • Phase Equilibrium - equilibrium between two phases in a closed system. • Evaporation and condensation occur at equal rates The rates of forward and reverse reactions are equal. As 0.50 moles of water are evaporated into the space above the water 0.50 moles of gaseous water are condensed back to liquid. The concentration of the reactants and products remain constant. Because the rates of the forward and reverse reaction are equal there is no net change to the amount of reactants or products produced. The may certainly be more of one than the other but those amounts do not change.

1. Physical Equilibrium a. Phase equilibrium • Melting and freezing occur at the same rate. There may not be the same amounts of solid and liquid present, but the rate of melting will be equal to the rate of freezing

1. Physical Equilibrium b. Solution equilibrium • Solution Equilibrium (two types) • gases in Liquids – equilibrium exists between the gas in the liquid and the gas above the liquid. The rates of forward and reverse reactions are equal. As 0.50 moles of CO2 gas comes out of solution into the space above the soda, 0.50 moles of it are dissolved back into solution. The concentration of the reactants and products remain constant. Because the rates of the forward and reverse reaction are equal there is no net change to the amount of reactants or products produced. There is certainly more gas dissolved into the liquid than in the space above it, but at equilibrium the amounts in each area remain constant.

1. Physical Equilibrium b. Solution equilibrium • Gases in Liquids – equilbrium exists between the gas in the liquid and the gas above the liquid. Factors that affect the equilibrium of gases in liquids • Increasing Temperature – decreases the solubility of gas in the liquid (warm soda has less bubble than cold soda). • Increasing Pressure – increases the solubility of gas in the liquid (bottling soda).

1. Physical Equilibrium b. Solution equilibrium • solids in liquids – equilibrium exists between dissolving and re-crystallizing of a solute in a liquid. • Solids and liquids exist in equilibrium in a saturated solution. Factors that affect the equilibrium of solids in liquids • Increasing/Decreasing Temperature - increases/decreases the solubility of solids in liquids (more sugar will dissolve in hot tea than iced tea).

Types of Equilibrium • Chemical Equilibrium – Forward and Reverse chemical reactions occur at the same time and rate. Equilibrium is reached when the rates of the forward and reverse reactions are equal. As a certain amount of HI is produced in the forward reaction that same amount of H2 and I2 is produced in the reverse reaction. Upon Initially mixing the reactants (H2 and I2) no products are present. As time goes on and more and more products are formed and we start to see an increase in the rate or the reverse (backward) reaction.

Reactions that go to Completion • Will not Reverse • 1) Closed System - if products are removed as they are made. The equilibrium is upset and the reaction will NOT reverse. • 2) Open System – reactions not taking place in a closedsystem. • Precipitate formed; Pb(NO3)2(l)+ 2KI(l) 2KNO3(l) + PbI2(s) • Gas formed; FeS(s) + 2HCl(aq) FeCl2(aq) + H2S(g) • Un-ionized product is formed; HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Questions • Which factors must be equal when a reversible chemical process reaches equilibrium. • mass of products and mass of reactants. • rate of forward reaction and rate of reverse reaction. • concentration of product and concentration of reactants. • activation energy of the forward reaction and activation energy of the reverse reaction. • A solute is added to water and a portion of the solute remains un-dissolved. When equilibrium between the dissolved and un-dissolved solute is reached, the solution must be • dilute • saturated • unsaturated • supersaturated • Which description applies to a system in a sealed flask is half full of water? • only evaporation occurs, but it eventually stops. • only condensation occurs but it eventually stops. • neither evaporation nor condensation occurs. • both evaporation and condensation occur.

Questions • Solution equilibrium always exists in a solution that is. • unsaturated. • saturated. • dilute. • concentrated. • Given a reaction at equilibrium. The addition of a catalyst will • shift the equilibrium to the right • shift the equilibrium to the left • increase the rate of forward and reverse reactions equally • have no effect on the forward or reverse reactions • If a catalyst is added to a system at equilibrium and the temperature and pressure remain constant, there will be no effect on the • rate of the forward reaction. • rate of the reverse reaction. • activation energy of the reaction. • heat of reaction (∆H).

Questions • In a reversible reaction, chemical equilibrium is attained when the • rate of the forward reaction is greater than the rare of the reverse reaction. • rate of the reverse reaction is greater than the rate of the forward reaction. • concentration of the reactants reaches zero. • concentration of the products remains constant. • The addition of a catalyst to a system at equilibrium will increase the rate of • the forward reaction only. • the reverse reaction only. • both the forward and reverse reactions. • neither the forward nor the reverse reaction. • A system is said to be in a state of dynamic equilibrium when the • concentration of the products is greater than the concentration of the reactants. • concentration of the products is less than the concentration of the reactants. • rate at which the product are formed is greater than the rate at which reactants are formed. • rate at which the products are formed is the same as the rate at which reactants are formed.

Le Chatelier’s Principle • Any change in concentration, temperature, or pressure on an equilibrium system is called a stress. • Le Chatelier’s Principle explains how a system at equilibrium responds to a stress on the system. • The equilibrium will shift in a way that tehnds to counteract the change.

Concentration Changes • An increase in concentration of any substance shifts the reaction away from the location of the increased substance. • Increasing the concentration of a reactant will shift the reaction to the right producing more products. • Increasing the concentration of a product will shift the reaction to the left producing more reactants. • An decrease in concentration of any substance shifts the reaction toward the location of the decreased substance. • Decreasing the concentration of a reactant will shift the reaction to the left producing more reactants. • Decreasing the concentration of a product will shift the reaction to the right producing more product.

Concentration Changes - Example Stress: More NH3 Stress: Less O2 Stress: Less NO Stress: More H2O

Common Ion Effect • Adding a substance with an ion that is common to the reaction. • Causes a shift in equilibrium because it favors the formation of a precipitate. What Happens if NaCl(s) is added? What happens if KBr(s) is added?

Temperature Changes • An increase or decrease in temperature changes the amount of energy into/out of the system. • We can consider heat as a reactant or product. • An increase in temperature shifts the reaction away from heat. • A decrease in temperature shifts the reaction towards heat.

Temperature Changes - Example Stress: More Heat (increased temperature) Stress: Less Heat (decreased temperature)

Pressure Changes • Pressure does not affect the rate of reaction of solids or liquids. • Only gases are affected!!

Pressure Changes • Pressure Increases: shift to the side with the least moles of gas. • System shifts to the right forming more NH3(g) • Pressure Decreases: shift to the side with the most moles of gas. • System shifts to the left forming more N2(g) and H2(g)

Pressure Changes • Both sides have the same number of moles of gas. • In this case, pressure changes have NO EFFECT on the system.

Catalysts • Favor the forward and reverse reactions equally. • Does NOT shift equilibrium.

Equilibrium Diagrams • The graph below represents the following chemical equilibrium before and after a stress has been imposed upon the system. Stress: increase H2 Result: equilibrium shifts away from the increase get decreased N2 get increased NH3

Practice Problem • What effect do the following changes have on the position of equilibrium for this reversible reaction? How must the reaction shift to relieve the “stress caused by change”? • addition of H2 • decrease in pressure • adding heat • removal of HI as it is formed

Questions • Consider the equation for the following reaction at equilibrium; The concentration of product could be increased by • adding a catalyst • adding more heat to the system • increasing the concentration of Y • decreasing the concentration of X

Questions • Consider the following equation What stress would cause the equilibrium to shift to the left? • increasing the temperature • increasing the pressure • adding N2(g) to the system • adding H2(g) to the system

Questions • Consider the equation for the following reaction at equilibrium. When Na2SO4 is added to the system, how will equilibrium shift? • The amount of CaSO4 will decrease, and the concentration of Ca2+(aq) will decrease. • The amount of CaSO4 will decrease, and the concentration of Ca2+(aq) will increase. • The amount of CaSO4 will increase, and the concentration of Ca2+(aq) will decrease. • The amount of CaSO4 will increase, and the concentration of Ca2+(aq) will icrease.

Questions • Consider the following equation As the concentration of HCl(aq) decreases at constant temperature, the rate of the forward reaction • decreases • increases • remains the same

Questions • Consider the following equation Which change will shift the equilibrium to the right? • decreasing the concentration of SO2 • decreasing the pressure • increasing the concentration of O2 • increasing the temperature

Practice Problems • What effect do the following changes have on the position of equilibrium for this reversible reaction? How must the reaction shift to relieve stress? • addition of Cl2(g) • increase in pressure • removal of heat • removal of PCl3 as it is formed

Practice Problem • How is the equilibrium position of this reaction affected by the following changes? • lowering temperature • increasing pressure • removing hydrogen from the equilibrium mixture • adding water to the equilibrium mixture • adding a catalyst • raising the temperature and decreasing the pressure

Practice Problems • The industrial production of ammonia is described by this reversible reaction. What effect do the following changes have on the equilibrium position? • addition of heat • increase in pressure • addition of a catalyst • removal of heat • removal of NH3

Review of Le Chatelier’s Principle

Chemical Equilibrium