Dynamic Equilibrium

Equilibrium Condition:

·  A balance between ______and ______processes

·  Constancy of macroscopic, but not microscopic properties.

·  MUST be a ______system

·  A system where some of all ______and ______are present

Reaching/Establishing Equilibrium:

·  As the reaction proceeds in the forward direction ( ) reactants are converted into products and the rate of the forward reaction (Ratef) ____ .

·  The amount of product ___ and the rate of the reverse reaction ____ . Products are converted back to reactants which reduces the Rater.

·  Eventually, the amount of new product forming is the same as reactant reforming and

Testing for Equilibrium:

·  Alter conditions such as concentration, temperature or pressure

·  if there is a visible macroscopic change, then the system ____ at equilibrium and is reacting to reach a new equilibrium.

·  if there is ___ change then the reaction is ______, or there is no ______at all.

Types of Equilibrium systems:

I)  Solubility

II) Phase

III)  Chemical Reactions

Percentage Reaction

• all reactions are considered reversible

• the amount of products or reactants can be determined by the % rxn:

If,

% Rxn / Conclusion
< 1 %
< 50 %
> 50 %
> 99 %
Percentage Reaction / A2 + B2 D 2AB / Equilibrium Conditions / % rxn
[AB]eq / 1.35 / 1.59 / 2.03
[B2]eq / 0.325 / 0.380 / 0.485
[A2]eq / 0.325 / 0.380 / 0.485
Initial Conditions / [AB]max
[AB]i / 0.00 / 1.35 / 3.00
[B2]i / 1.00 / 0.500 / 0.000
[A2]i / 1.00 / 0.500 / 0.000

ICE Tables

• more commonly used to determine equilibrium concentrations, [ ]eq than % rxn as [ ]max is not required.

eg. When 0.900 mol of N2 and 2.70 mol of Cl2 are added to a 1.00 L container and allowed to achieve equilibrium, the [Cl2]eq is found to be 0.440 M. Determine the [N2]eq and [NCl3]eq, according to:

N2 + 3 Cl2 D 2 NCl3


Equilibrium Law

1) Equilibrium Constant, Keq:

If, aA + bB D cC + dD

eg. For A2 + B2 D 2AB

2) Significance of Keq magnitude:

·  If Keq is very large then is large and the rxn is complete.

·  If Keq is very small then no rxn occurs.

3) Keq and the Balanced Chemical Eq’n:

·  any action performed on the chemical rxn, the Keq expression is raised to that action.

4) Keq and Reaction Kinetics:

If A + B D C + D

and both the forward and reverse rxns are elementary steps, then

Ratef = kf [A]1[B]1 and Rater = kr [C]1[D]1

At equilibrium:

5) Keq & the Effect of Temperature:

·  usually Eaf ≠ Ear.

·  an ­ in T won’t affect the forward and reverse reactions equally, so,

Ratef ≠ Rater

and the equilibrium will change

If exothermic:

at higher T an increased Rater creates more reactants than at lower T, and

If endothermic then the opposite occurs:


6) Heterogeneous Equilibria:

·  reactants and products in different physical states (s, l ,g, aq)

·  Pure liquids (not aq) and solids have constant densities

·  as a result, pure solids and liquids are not written in the Keq expression.

eg. 2H2O(l) D 2H2 (g) + O2 (g)

eg. Zn(s) + Cu2+(aq) D Cu (s) + Zn2+(aq)


Equilibrium and Spontaneity

∆Gº = -RT ln Keq where R = 8.314 J/mol•K

Calculate ∆Gº from Keq

eg. PCl3 + Cl2 « PCl5 ; Keq = 0.18

∆Gº = -RT ln Keq

Calculate Keq from ∆Gº

∆Gº = - RT ln Keq

ln Keq = - (∆Gº) / (RT)

Keq = e-(∆Gº)/(RT) * Note : ∆Gº in J/mol

eq. From before

2CO(g) + O2(g) à 2CO2(g)

and ∆Gº = -514.5 kJ


Quantitative Changes in Equilibrium Systems

Reaction Quotient, Q

· is a test calculation of the equilibrium system using initial, NOT equilibrium concentrations.

· by comparing the value of Q to Keq, the direction that the equilibrium will change is determined.

If: / Reaction System Response
Q < Keq
Q = Keq
Q > Keq

eg. Consider the following system:

2 SO2 + O2 D 2 SO3 ; Keq =

a) 2.0 1.0 0 Q = 0 < Keq \®

b) 0.50 0.75 1.0 Q = 5.3>Keq \¬

Calculating [ ]eq from [ ]i and Keq

· Write the balanced chemical reaction & Keq.

· Calculate Q to determine the sign of “C” (Change).

· Create an I.C.E. table, input I ( [ ]i ) and C (± nx) and then form the E, ( [ ]eq ) expressions, [ ]i ± nx.

· Substitute [ ]eq expressions into Keq.

· Solve for the unknown, “x”, by using known relationships, (perfect squares), assumption, (100 Rule, ) or use the quadratic equation, ().

· Determine [ ]eq values by substituting the calculated “x” value.

· Check your calculated [ ]eq values by substituting into Q and comparing to Keq.


eg. The Keq for the dissociation of I2 into I atoms is 3.80 x 10-2. If the [I2]i is 0.200 M, calculate the [I]eq.


eg. If 0.100 mol of HCl is placed in a 1.00 L container and allowed to reach equilibrium, find [ ]eq for all species given that Keq = 1.84 x 10-2


eg. If 1.0 mol of NOCl is placed in a 2.0L flask, calculate [ ]eq if Keq = 1.6 x 10-5.


Le Chatelier’s Principle

· An equilibrium under stress will ______in whichever direction ______the stress

· If temperature is maintained, Keq is the same (constant)

· One Rate rxn will increase to offset the stress until

Rrxn f = Rrxn r

and equilibrium is re-established

CoCl42-(al) + 6H2O D Co(H2O)62+ + 4Cl-1 + NRG

2CrO42- + 2H1+ D Cr2O72- + H2O

N2O4(g) + NRG D 2NO2(g)

CO2 + H2O D 2H1+ + HCO31-

(with bromothymol blue indicator)

Tb2- + H3O+ D HTb 1- + H2O

HTb 1- + H3O+ D H2Tb + H2O


Solubility Product Constant, Ksp

Solubility:

• is the ______amount of solute in a solvent at a given temperature

• saturated solution, [ ]max

• equilibrium between: (see Dissolving salts OH)

______« ______

(see dissolving salts handout)

eg. AgNO3(s) « Ag1+(aq) + NO31-(aq)

• solubility is a contest between:

• ionic bond strengths vs ______

• lattice energies vs ______

• usually exothermic with ______disorder

• if undissolved solid is present, then it is a ______with a saturated sol’n phase

Solubility Product Constant, Ksp:

• equilibrium expression for slightly soluble salts

If: AxBy(s) « xAm+(aq) + yBn-(aq)

Saturation can be tested by Q:

If: Q = Ksp then:

Q > Ksp

Q < Ksp

Terminology:

• refers to the ______of solid that can dissolve not how much is in ______

• solubility in g/100g or g/100 mL or g/L

[ ]max = solubility x Msolute and volume adjustment

• molar solubility (mol/L) ______

= solubility / Msolute

Calculations:

1) Ksp from solubility

eg. Calculate the Ksp of Ag2CO3 given its solubility of 0.0014 g/100g.

2) Solubility from Ksp

types: solubility, molar solubility, [ion]eq, amount of solid that will dissolve

eg. How much PbI2 at SATP will dissolve in 1.00 L of water? Give the solute ion concentration, [Pb2+]eq, solubility (g/100mL) and molar solubility.

(Ksp of PbI2 from textbook is 8.5 x 10-9)

Predicting Precipitation:

• Used to determine precipitation when mixing 2 sol’ns as with double displacement rxns

eg. Will precipitation occur when 5.0 ml of 0.030 M AgNO3 is mixed with 1.0 ml of 0.0050 M Na2CrO4? (Ksp = 1.1 x 10-12 for Ag2CrO4).


Common Ion Effect

· Application of Le Chatelier’s Principle to solubility

For: AB2(s) « A2+(aq) + 2 B1-(aq)

the equilibrium can shift in the:

i) ______direction by redissolving solid

ii) ______direction by forming precipitate

Re-dissolving Precipitates

1.  Add more solvent

then

For AB ßà A+ + B-

o  A precipitate forms or exists if

o  Adding solvent it ______the ion concentrations and equilibrium shifts to the ______

2.  Addition of an Acid

a) CaCO3(s) ßà Ca2+ + CO32-

CO32- + 2H+ à CO2 + H2O

As CO32- falls below ksp

b) Cu(OH)2(s) ßà Cu2+(aq) + 2OH1-(aq)

H+ + OH- à H2O

3.  Addition of a Complexing Agent

Cu(OH)2 ßà Cu2+ (down) + 2OH-

Added:

Cu2+(aq) + NH3 à Cu(NH3)42+

eg. Given that Ksp(AgI) = 4.5 x 10-17, calculate the molar solubility of AgI in:

a)  pure water

b) 0.10M NaI