Equilibrium • The state where the concentrations of all reactants and products remain constant with time.
Reactions are reversible • A + B C + D ( forward) • C + D A + B (reverse) • Forward and Reverse Rxns can be shown by double arrow A + B C + D
A + B C + D • Initially there is only A and B so only the forward reaction is possible • As C and D build up, the reverse reaction speeds up while the forward reaction slows down. • Eventually the rates are equal • So concentrations of the reactants and products no longer change with time
Forward Reaction Reaction Rate Equilibrium Reverse reaction Time
Static or Dynamic? • At equilibrium, forward and reverse reaction rates are equal • May seem like no changes are occurring but there are changes • No NET changes • On the molecular level, there is frantic activity. Equilibrium is not static, but is a highly dynamic situation. • chemical reactions take place ,but concentrations of reactants and products remain unchanged
Analogies and Metaphors to think about • In a football game, the number of players on the field is constant although exchange of players (substitution) changes actual persons. • Connected fish bowl analogy . Two fish tanks are connected by a tube large enough to allow passage of fish. A number of fish are placed in one of the tanks. At equilibrium, the number of fish in each tank will eventually become unchanged. • Two jugglers analogy. • Drinking fountain line: • Ten students waiting in line to get a drink of water on a hot day. As each gets a drink, the same student reenters the line (equilibrium in a closed system). • (b) Same situation as "a," except as each student gets a drink and leaves, a new student enters the line (steady-state in an open system). • Picture a number of horses and wranglers in a corral. As each wrangler mounts a horse, the wrangler is bucked off. The equilibrium is: Horse + Wrangler Mounted wrangler
Molecular Simulation • In this simulation two gaseous reactants collide to produce a more dense solid. A + B C gaseous R dense P • http://www.absorblearning.com/media/attachment.action?quick=w8&att=2310
Homo vs Hetero • Homogeneous Equilibria all reacting species are in the same phase • gas phase • equilibrium constant can be expressed in terms of pressure or concentration, Kp or Kc • Solution (aqueous) phase • concentration term for the pure liquid does not appear in the expression for the equilibrium constant but aqueous substance concentrations do appear • Heterogeneous Equilibria all reacting species are not in the same phase • concentration term for solid or liquid does not appear in the expression for the equilibrium constant
Equilibrium Summarized • Forward and Reverse rates are equal • Concentrations are not. • Rates are determined by concentrations and activation energy. • Molecular Motion is frantic and constantly changing • Macroscopically no net change is occurring (We can’t observe any changes)
Distinguishing between Physical and Chemical Equilibrium • As with physical and chemical changes, physical and chemical equilibrium follow the same rules: • Physical no changes to the chemical properties of the substances involved • Ex. equilibrium of water vapor with liquid water in a partly filled sealed bottle • Chemical involve changes in the chemical composition of substances. Bond breaking and bond formation is involved. • Ex. dissociation of acetic acid water into acetate and hydronium ion
Activity • Model Dynamic Equilibrium with Coins • Now lets plot the data using excel
c c Law of Mass Action • For a reaction: aA + bB ⇄ cC + dD • equilibrium constant: K • Pure liquids and pure solids have concentrations of 1.
c c Playing with K • If we write the reaction in reverse. cC + dD ⇄ aA + bB • Then the new equilibrium constant is
They are simply the inverse of one another. • Forward Reaction aA + bB ⇄ cC + dD So we call this K1 And K1= 1 = K2-1 K2 • Reverse Reaction cC + dD ⇄ aA + bB So we call this K2 And K2= 1 = K1-1 K1
The units for K • Are determined by the various powers and units of concentrations. • They depend on the reaction. • will always have the same value at a certain • temperature (Why will the T effect it?) • no matter what amounts are initially added • ratio at equilibrium will always be same
K • Has no units • Is constant at any given temperature. • Is affected by temperature. • Equilibrium constants are reaction, phase, temperature and pressure dependent • There is a K for each temperature. • Equilibrium constant values are thus established for a specific reaction in a specific system and will be unchanging (constant) in that system, providing the temperatures does not change.
What does the size of my K mean? • Large K > 1 • products are "favored“ • Ex. 1 x 1034 • K = 1 • neither reactants nor products are favored • Small K < 1 • reactants are "favored“ • Ex. 4 x 10-41
Now Let’s Calculate Your K • Using your data from the simulation calculate K
Different Equilibrium Constants (All of these are known as Keq) • Kcis the most used general form with molar concentrations. • Kpcan be used with partial pressures when working with a gas phase reaction. • Kais used for the dissociation of weak acids in water. • Kbis used for the dissociation of weak bases in water. • Kwis the equilibrium expression for the dissociation of water into its ions. • Kspis used for the dissociation into ions of sparingly soluble solids in water.
Practice Writing the Equilibrium Expression 4NH3(g) + 7O2(g) 4NO2(g) + 6H2O(g) First write the equilibrium expression using no concentration values. What is the value for K if the concentrations are as follows NH3 1.0 M O2 1.0 M NO2 1.4 M H2O 1.8 M
Equilibrium with Gases • Equilibria involving only gases can be described using pressures or concentrations • If using pressures, • use pA not [A] • KP not KC • be sure all pressure are in the same units N2(g) + 3H2(g) 2NH3(g)
Calculating Kc from Kp • where • Δn is the difference in moles of gas on either side of the equation (np – nr) • R is the gas law constant: 0.08206 • T is Kelvin temperature For: N2(g) + 3H2(g) 2NH3(g) • Δ n = (2) – (1+3) = -2 • h
Practice • Setup the expression for KP in terms of KC, R and T 2NO(g) + Cl2(g) 2NOCl(g)
What the equilibrium constant tells us… • if we know the value of K, we can predict: • tendency of a reaction to occur • if a set of concentrations could be at equilibrium • equilibrium position, given initial concentrations • If you start a reaction with only reactants: • concentration of reactants will decrease by a certain amount • concentration of products will increase by a same amount
Using this we can make ICE charts The following reaction has a K of 16. You are starting reaction with 9 O3 molecules and 12 CO molecules. Find the amount of each species at equilibrium.
Practice 1 ICE Charts • Consider the following reaction at 600ºC • 2SO2(g) + O2(g) 2SO3(g) • In a certain experiment 2.00 mol of SO2, 1.50 mol of O2 and 3.00 mol of SO3 were placed in a 1.00 L flask. At equilibrium 3.50 mol were found to be present. Calculate • The equilibrium concentrations of O2 and SO2, K and KP
Practice Problem 2 ICE Charts • Consider the same reaction at 600ºC • In a different experiment .500 mol SO2 and .350 mol SO3 were placed in a 1.000 L container. When the system reaches equilibrium 0.045 mol of O2 are present. • Calculate the final concentrations of SO2 and SO3 and K
The Reaction Quotient (Q) • Tells you the directing the reaction will go to reach equilibrium • Calculated the same as the equilibrium constant, but for a system not at equilibrium • Q = [Products]coefficient [Reactants] coefficient • Compare value to equilibrium constant
What Q tells us IF THEN Q = K reaction is at equilibrium Q > K too much products, left shift Q < K too much reactants, right shift
Example 1 Reaction Quotient For the synthesis of ammonia at 500°C, the equilibrium constant is 6.0 x 10-2. Predict the direction the system will shift to reach equilibrium in the following case:
Example 2 In the gas phase, dinitrogen tetroxide decomposes to gaseous nitrogen dioxide: Consider an experiment in which gaseous N2O4 was placed in a flask and allowed to reach equilibrium at a T where KP = 0.133. At equilibrium, the pressure of N2O4 was found to be 2.71 atm. Calculate the equilibrium pressure of NO2.
Example 3 At a certain temperature a 1.00 L flask initially contained 0.298 mol PCl3(g) and 8.70x10-3 mol PCl5(g). After the system had reached equilibrium, 2.00x10-3 mol Cl2(g) was found in the flask. PCl5(g) PCl3(g) + Cl2(g) Calculate the equilibrium concentrations of all the species and the value of K.
Approximations • If K is very small, we can assume that the change (x) is going to be negligible compared to the initial concentration of the substances • can be used to cancel out when adding or subtracting from a “normal” sized number to simplify algebra
Example 4 At 35°C, K=1.6x10-5 for the reaction 2NOCl(g) ⇄ 2NO(g) + Cl2(g) Calculate the concentration of all species at equilibrium for the following mixtures 2.0 mol NOCl in 2.0 L flask 1.0 mol NOCl and 1.0 mol NO in 1.0 L flask 2.0 mol NO and 1.0 mol Cl2 in 1.0 L flask
Le Chatelier’s Principle • can predict how certain changes or stresses put on a reaction will affect the position of equilibrium • helps us determine which direction the reaction will progress in to achieve equilibrium again • system will shift away from the added component or towards a removed component
Change Concentration • equilibrium position can change but not K • system will shift away from the added component or towards a removed component Ex: N2 + 3H2 2NH3 • if more N2 is added, then equilibrium position shifts to right (creates more products) • if some NH3 is removed, then equilibrium position shifts to right (creates more products)
Adding Gas • adding or removing gaseous reactant or product is same as changing concentration • adding inert or uninvolved gas • increase the total pressure • doesn’t effect the equilibrium position
Change the Pressure by changing the Volume • only important in gaseous reactions • decrease V • requires a decrease in # gas molecules • shifts towards the side of the reaction with less gas molecules • increase V • requires an increase in # of gas molecules • shifts towards the side of the reaction with more gas molecules