Chemical Equilibrium

Chemical Equilibrium: Chapter 13
Definitions
Chemical equilibrium- when a reaction takes place in a closed system, it reaches a condition where the concentrations of the reactants and products remain constant over time and the rate of the forward and reverse reactions occur simultaneously and at the same rate
Law of mass action-there is a relationship between the between the concentration of reactants and products at equilibrium
Equilibrium constant- for the reaction of aA + bB <-> cC + dD the equilibrium expression is:
equilibrium expression.jpeg
Equilibrium constant- an expression providing a constant ratio of the reactants and products of the system at a constant ration: K, Keq, Kc
Homogeneous Equilibria- where all the reactants and products are in the same phase (ex. gases)
Heterogenous Equilibria- equilibria involving more than one phase
Reaction Quotient- Q, it is obtained by applying the law of mass action using initial concentrations instead of equilibrium concentrations. If Q is equal to K, the system is at equilibrium. If Q is greater than K, the system shifts to the left forming reactants. If Q is less than K, the system shifts to the right forming products.
Le Chatliers Principle- if a change is imposed on a system at equilibrium, the position of the equilibrium will shift in the direction that tends to reduce the change
Combined OLD PPT equilibrium.pptx
Learning Objectives
Learning objective 6.1- The student is able to, given a set of experimental observations regarding physical, chemical, biological, or environmental processes that are reversible, construct an explanation that connects the observations to the reversibility of the underlying chemical reactions or processes.
Learning objective 6.2- The student can, given a manipulation of a chemical reaction or set of reactions (ex. reversal or addition of two reactions), determine the effects of the manipulation on Q or K.
Learning Objective 6.3- The student can connect kinetics to equilibrium by using reasoning about equilibrium, such as Le Chatlier's principle to infer the relative rates of the forward and reverse reactions.
Learning Objective 6.4- The student can, given a set of initial conditions (concentrations or partial pressures) and the equilibrium constant, K, use the tendency of Q to approach K to predict and justify the prediction as to whether the reaction will proceed toward reactants or products as equilibrium is approached.
Learning Objective 6.5- The student can, given data (tabular, graphical, etc.) from which the state of equilibrium can be obtained, calculate the equilibrium constant K.
Learning Objective 6.6- The student can, given a set of initial conditions (concentrations or partial pressures) and the equilibrium constant, K, use the stoichiometric relationships and the law of mass action (Q equals K at equilibrium) to determine qualitatively and/or quantitatively thee conditions at equilibrium for a system involving a single reversible reaction.
Learning Objective 6.7- The student is bale, for a reversible reaction that has a large or small K, to determine which chemical species will have very large versus very small concentrations at equilibrium.
Learning Objective 6.8- The student is able to use Le Chatlier's principle to predict the direction of the shift resulting from various possible stresses on a system at chemical equilibrium.
Learning Objective 6.9- The student is able to use Le Chatlier's principle to design a set of conditions that will optimize a desired outcome, such as a product yield.
Learning Objective 6.10- The student is able to connect Le Chatlier's principle to the comparison of Q to K by explaining the effects of the stress on Q and K.
Section 13.1- The Equilibrium Condition
Although in equilibrium the concentrations of reactants and products remain the same, equilibrium is a dynamic state where the reactants and products are interconverted continually.
The concentrations of neither the reactants nor the products reach zero during equilibrium.
Equilibrium is a state in which the rate of the forward reaction equals the rate of the reverse reaction.
Equilibrium occurs in a closed system.
https://www.youtube.com/watch?v=-l5zWz_TMbM
CNX_Chem_13_01_equilibrium.jpg
Section 13.2- The Equilibrium Constant
The law of mass action: for the reaction aA + bB <-> cC + dD the equilibrium constant is
equilibrium expression.jpeg
The equilibrium constant can also be K, Kc, or Keq and is temperature dependent.
Example equilibrium expression.docx
Example of calcualting the values of K.docx
The equilibrium expression for a reaction is the reciprocal of that for the reaction in reverse. Also, when the balanced equation for a reaction is multiplied by a factor of n, the equilibrium expression for the new reaction is the original expression raised to the nth power. Finally, n is written without units.
For any given temperature there are many equilibrium positions ( a set of equilibrium concentrations) but there is only 1 value of K (which depends on the ratio of the concentrations).
https://www.youtube.com/watch?v=xfGlEXWDRZE&index=65&list=PLllVwaZQkS2op2kDuFifhStNsS49LAxkZ
Section 13.3- Equilibrium Expressions Involving Pressures
Equilibria involving gases can also be described in terms of pressure. For the equation: aA + bB <-> cC + dD
equilibrium constant partial pressure.JPG
The relationship between K and Kp can be described as Kp=K(RT)^delta (n). Delta n is the sum of the coefficients of the gaseous products minus the sum of the coefficients of the gaseous products. Thus in the equation: aA + bB <-> cC + dD, delta n is equal to ( d + c)- (a + b).
relationship between k and kp.png
equilibrium expression with pressure.docx
https://www.youtube.com/watch?v=2mIrIlVsF0I
https://www.khanacademy.org/science/chemistry/chemical-equilibrium/equilibrium-constant/a/calculating-equilibrium-constant-kp-using-partial-pressures
Section 13.4- Heterogeneous Equilibria
Homogenous equilibria is where all the reactants and products are in the same phase, like gases. Heterogeneous equilibria is where equilibria involve more than one phase.
Heterogeneous equilibria doesn't depend on the amounts of solids or liquids present since the concentration of pure liquids or pure solids cannot change.
Thus, if pure solids or pure liquids are involved in a chemical reaction, the concentrations are not included in the equilibrium expression. Only solutions and gases are included.
heterogenous equilibria.jpg
heterogenous equilibria 2.jpg
https://www.youtube.com/watch?v=xfGlEXWDRZE&index=65&list=PLllVwaZQkS2op2kDuFifhStNsS49LAxkZ
Section 13.5- Applications of the Equilibrium Constant
A large K value means that the equilibrium lies far to the right (consists mostly of products). A small K means the equilibrium lies far to the left ( consists mostly of reactants).
The equilibrium constant allows us to predict the tendency of a reaction to occur, whether a given set of concentrations represents an equilibrium condition, and the equilibrium position that will be achieved from a given set of initial concentrations.
The size of K and the time required to reach equilibrium are not directly related.
If the concentration of one of the reactants or products is zero, the reaction will shift in the direction that produces component. If all the initial concentrations are nonzero, to determine the shift the reaction quotient, Q, can be used.
The reaction quotient is obtained by applying the law of mass action using initial concentrations instead of the equilibrium concentrations.
Three possible cases: 1. Q= K ( the system is at equilibrium and no shift will occur) 2. Q is greater than K ( the system shifts to the left consuming products and forming reactants) 3. Q is less than K ( the equilibrium will shift to the right consuming reactants and forming products)
If some of the equilibrium concentrations or initial values are not given, stoichiometry needs to be used to express the concentrations at equilibrium.
reaction quotient examples.docx
https://www.youtube.com/watch?v=s_iDJkaZPdw
https://www.khanacademy.org/science/chemistry/chemical-equilibrium/factors-that-affect-chemical-equilibrium/a/the-reaction-quotient
Section 13.6- Solving Equilibrium Problems
1. Write the balanced equation for each reaction.
2. Write the equilibrium expression using the law of mass action.
3. List the initial concentrations.
4. Calculate Q, and determine the direction of the shift to equilibrium.
5. Define the change needed to reach equilibrium, and define the equilibrium concentrations by applying the change to the initial concentrations.
6. Substitute the equilibrium concentrations into the equilibrium expression, and solve for the unknown.
7. Check your calculated equilibrium concentrations by making sure they give the correct value of K.
After finding the change in concentration using a calculator or quadratic formula, use the x value that will NOT produce any negative concentrations.
If the initial concentrations divided by K is equal to a number greater than 500, the approximation method can be used.
solving equilibrium problems.docx
https://www.youtube.com/watch?v=RlfctnfZrQo
https://www.youtube.com/watch?v=mHZX9wbfO4A
https://www.youtube.com/watch?v=67t7Q2ftKVU
Section 13.7- Le Chatlier's Principle
Le Chatlier's principle- if a change is imposed on a system at equilibrium, the position of the equilibrium will shift in a direction that tends to reduce that change.
If a component (reactant or product) is added to a reaction system at equilibrium (at constant T and P or constant T and V), the equilibrium position will shift in the direction that lowers the concentration of the component. If a component is removed, the opposite effect occurs.
The addition of an inert gas increases the total pressure but has no effect on the concentrations or partial pressures of the reactants or products. (equilibrium position unaffected)
When the volume of a container is decreased, the equilibrium shifts to the side of the reaction involving a smaller number of gas molecules. The opposite is true if the volume is increased.
The value of K changes with temperature.
If the temperature increases for an exothermic reaction, the equilibrium shifts to the right. If the temperature increases for an exothermic reaction the equilibrium shifts to the left.
Le Chatliers Principle Practice.docx
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/lechv17.swf
https://www.youtube.com/watch?v=4-fEvpVNTlE
le chatliers principle.png
le chatliers image.jpg
Multiple Choice Examples
Sample Multiple Choice for Chemical Equilibrium.docx
Key for Sample Multiple Choice for Chemical Equilibria.docx
Free Response Examples
Free Response Questions for Chemical Equilibria.docx
Key for Free Response Questions for Chemical Equilibria.docx
Enables qualitative prediction of the effects of changes in concentration, pressure, and temperature on a system at equilibrium.
When a reaction takes place in a closed system, it reaches a condition where the concentration of the reactants and products remain constant over time.
The law of mass action describes the equilibrium constant. Also, there is only 1 value of K for a give system at a given temperature but there are an infinite number of equilibrium positions depending on the initial concentrations.
For a gas-phase reaction, the reactants and products can be described in terms of their partial pressures and the equilibrium constnat is called Kp
Pure solids or liquids are not included in the equilibrium expression.
A small value of K means the equilibrium lies to the left and a large value means the equilibrium lies to the right. Q, the reaction quotient, applies the law of mass action to initial concentrations. If Q>K, the system will shift to the left and if Q<K, the system will shift to the right.
To find the equilibrium concentrations, start with the initial concentrations, define the change, then apply the change to the initial concentrations and solve for the equilibrium concentrations.
14