Presentation on theme: "Chemistry 30 Equilibrium. Equilibrium in Chemical Systems Reactions do not always go forward to completion (reactants do not all react) Instead reach."— Presentation transcript:
Chemistry 30 Equilibrium
Equilibrium in Chemical Systems Reactions do not always go forward to completion (reactants do not all react) Instead reach equilibrium: concentration of reactants and products becomes constant (but not equal) Use an arrow in the equation to show equilibrium
Ex) ColorlessViolet Colorless
Dynamic equilibrium – a balance between the forward and reverse reactions occurring at the same rate with constant macroscopic properties (temp, ph, color etc).
Although rates are the same, concentrations of reactants and products are different at equilibrium equilibrium is reached when concentrations stop changing
Demonstration of simulated chemical equilibrium
Two Types of Equilibrium 1. Phase change equilibrium Solid in a saturated solution ex) sugar and sugar water Gas above a liquid ex) pop Vapour above a solid ex) mothballs in a dresser
2. Chemical equilibrium Homogeneous: reactants and products are in the same phase or state Ex) C 2 H 5 OH(l) + CH 3 COOH(l) CH 3 COOC 2 H 5 (l) + H 2 O(l) Heterogeneous: reactants and products are in different phases or states Ex) CaCO(s) CaO(s) + CO 2 (g)
4 Conditions for Equilibrium 1. Forward rate and reverse rate of reaction are the same 2. Macroscopic properties are constant 3. Closed system 4. Equilibrium can be approached from either direction
Classes of Chemical Reaction Criteria Percent Reaction Description at Equilibrium Position at equilibrium <50% Reactants favored >50% Products favored >99% No eq’b
Equilibrium Constant Ratio of the product and reactant concentrations is a constant value. This value is called the equilibrium constant K c
Finding K c For the reaction: aA + bB cC + dD Kc= [C] c [D] d = products [A] a [B] b reactants *Only substances whose concentrations or pressures change can be included in the equilibrium expression. Liquids and solids are not included.
If K>1 the reaction favors products If K<1 the reaction favors reactants If K=1 there are equal concentrations of reactants and products
Calculations with Equilibrium Constants Case 1 – Given all the amounts of species at equilibrium, calculate the equilibrium constant. Case 2 – Given the initial and final conditions, calculate the equilibrium constant.
In each case, you will need: The balanced chemical equation for the system at equilibrium The equilibrium constant expression (remember to include only those that are solutions or gases, not liquids or solids)
Case 1 Examples 1. Calculate the Kc for the following system at a constant temperature: 2NO(g) + O 2 (g) 2NO 2 (g) [NO]: 1.0 mol/L [O 2 ]: 6.0 mol/L [NO 2 ]: 2.0 mol/L
2. For the reaction CO(g) + Cl 2 (g) COCl 2 (g) at a particular temperature, Kc = 5.1 x At equilibrium, there are 0.30 mol of Cl 2 (g) and mol of COCl 2 (g) in a 2.0L container. What is the equilibrium concentration of CO(g)?
Case 2 Examples Use the ICE tables I= initial C=change E= equilibrium Need to use molar ratio as well
1. For the reaction given below, 3.00 mol of A and 4.00 mol of B are placed in a 5.00 L container. 2A(g) + B(g) C(g) At equilibrium, the concentration of A is 0.40 mol/L. Determine the value of K c.
2. Initially 2.0 mol of N 2 and 4.0 mol of H 2 were added to a 1.0 L container and the following reaction occurred. 3H 2 (g) + N 2 (g) 2NH 3 (g) The equilibrium concentration of NH 3 is 0.68 mol/L. What is the Kc ?
Le Chatelier’s Principle when a system at equilibrium is disturbed by a change in a property, the system adjusts in a way that opposes this change. Change can occur to: Concentration Pressure or volume Temperature (this will also change Kc) Add a Catalyst
The application of Le Chatelier’s Principle involves three stages: 1. The initial state of equilibrium in the system 2. Stress exerted on the system by some change in property (temp, conc, pressure) that creates dis- equilibrium 3. Shift towards reactants or products, to re-establish a new state of equilibrium Le Chatelier’s Principle can provide a method of predicting the reaction’s response to the change. y/flash/lechv17.swf y/flash/lechv17.swf
1.Concentration Changes The addition of more reactant or the removal of a product will increase the yield of the product, shifting the equilibrium towards the right (products) A + B C If the opposite occurs, the shift will be towards the left (reactants).
Concentration change can occur by adding something that makes a precipitate. Ex) AgNO 3(s) Ag + (aq) + NO 3 - (aq) If you add NaCl, the Cl - (aq) reacts with Ag + (aq) to form AgNO 3(s) Response is to shift to the right AgNO 3(s) Ag + (aq) + NO 3 - (aq)
2. Pressure / Volume Change Gases only If pressure decreases, response is to increase pressure by shifting to the side with more moles of gas Ex) Decrease pressure and : A (g) + 2B (g) 1C (g)
In a gaseous system, decreasing the volume of the container increases the pressure. Adding more gas increases the pressure Pressure change has no effect on equilibrium if moles are the same on both side.
3. Temperature Change The energy in a chemical reaction is treated as if it were a reactant or product. Endothermic: reactants + energy products Exothermic: reactants products + energy Energy can be added or removed by heating or cooling the system. Equilibrium shifts to minimize the change. Will change K c
Ex) 2 SO 3(g) + 97 kJ 2 SO 2(g) + O 2(g) If you heat it up (the stress), response is to cool down and equilibrium shifts right. Kc increases If you cool it down (the stress), response is to heat up and equilibrium shifts left. Kc decreases.
4. Catalyst Decreases time to reach equilibrium Lowers activation energy by the same amount in both the forward and reverse directions Does not effect final equilibrium position
Summary VariablesChange Response of the system ConcentrationIncrease Shift to consume added reactant or product ConcentrationDecrease Shift to replenish some of the removed reactant or product TemperatureIncrease Shifts to consume some of the added heat TemperatureDecrease Shifts to replenish some of the removed heat Volume (overall pressure) Increase volume / decrease pressure Shifts toward the side with the larger number of gaseous molecules Volume (overall pressure) Decrease volume / increase pressure Shifts toward the side with the smaller number of gaseous molecules