Equilibrium

•      HF (aq) = H+ (aq) + F- (aq)      K1 =6.8*10 -4

•       H2C2O4(aq) = 2 H +(aq) +C2O4-2(aq)                                                              K2=3.8*10 -6

•      What is K for

•      2 HF + C2O4-2(aq) = 2 F- (aq) + H2C2O4(aq)

 

•      [.12]

example

•      For the reaction

•            2H2(g) + O2 (g) = 2 H2O(l)

•                            Kp = 1.5 * 1083

•      Is this reaction reversible?

 

•      Will this reaction proceed?

 

 

For the Haber process

•      N2(g) + 3 H2(g) = 2 NH3 (g)

•            Kp = 1.45*10-5 at 500 C, in an equilibrium mixture, PH2 = 0.928 atm, PN2 = 0.432 atm, what is PNH3?

 

Heterogeneous Equilibriaple

•      Sometimes the substances in equilibrium are different phases –

•            CaCO3(s) = CaO(s) + CO2(g)

•            Kc = [CaO] [CO2] / [CaCO3]

•      How do we express the concentration of a solid substance?

•            Kc’ = c1 [CO2] / c2 = c1/c2 [CO2]

•      A pure liquid or solid’s concentration is not included in the equilibrium expression

•      Bell Jar experiment

 

•      Each of the following mixtures was placed in a closed container and allowed to stand – which mixtures are capable of attaining equilibrium for

•      CaCO3(s) = CaO(s) + CO2 (g)

•      1.   Pure CaCO3(s)

•      2.   CaO and a pressure of CO2(g) greater than Kp

•      3.   Some CaCO3(s) and a pressure of CO2(g) greater than Kp

•      CaCO3(s) and CaO(s)

 

Calculate K

•   1.       Tabulate initial and equilibrium concentration of all species involved in equilibrium

•   2.       Calculate the change in concentration for known species

•   3.       Use stoichiometry to calculate changes for all other species

•   4.calculate K

Exampple

•      a mixture of 5.00 10-3 mol of H2 and 1.00*10-2 mol of I2 are placed in a 5.00 L container at 448 C and allowed to come to equilibrium.  [HI]eq = 1.87 *10-3 M, calculate Kc for the reaction

•            H2 (g) + I2 (g) = 2 HI (g)

•      I                                                   

•      D

•      F                                         Kc = 51

Le Chatelier’s Principle

•      If a system at equilibrium is disrupted by a change in temperature, pressure or the concentration of one of the components – the system will shift its equilibrium positions so as to counteract the effect of the disturbance

Example

•      Consider

•            N2 (g) + 3 H2 (g) = 2 NH3 (g)

 

•      What would happen if I added H2?

•      What if I increased pressure?

•      If I wanted to make a lot of NH3 – how could I do that?  Why would N2 be hard to break up?

Q, reaction quotient

•      K is a constant at a given T, Q is the nonequilibrium analog of K - recall K only valid at equilibrium

•      Q approaches K

•            If Q > K reaction goes ί

•            If Q < K reaction goes ΰ

•            If Q = K equilibrium

 

Example

•      Consider if we have a mixture of 2.00 mol of H2 and 1.00 mol of N2 and 2.00 mol of NH3 in a 1.00 l container at 472 C.  Kc is 0.105 – What direction will the reaction proceed in?