1. Reaction Rates and Equilibrium
M.Elizabeth
2011

2. Collision Theory Used to Explain Reaction Rates
Atoms, ions, and molecules can form a chemical bond when they collide, provided the particles have enough kinetic energy.
Particles lacking the necessary kinetic energy to react still collide, but simply bounce apart.
Activation energy - the minimum energy colliding particles must have in order to react.

3. Chemical Reactions
ordinarily occur as a result of collisions between reacting particles. Consider the reaction:
CO(g) + NO2(g) ----> CO2(g) + NO(g) rate = k(conc CO)(conc NO2(g))
doubling the [CO], holding [NO2] constant, the number of collisions in a given time doubles.
doubling the [NO2] , holding CO constant, has the same effect.
the number of collisions per unit time is directly proportional to the concentration of CO or NO2.
The fact that the rate is directly proportional to these concentrations indicates that reaction occurs as a direct result of collisions between CO and NO2 molecules.

4. NOT EVERY COLLISION LEADS TO REACTION!!!!!
It is possible to calculate the rate at which molecules collide with each other by using the kinetic theory. Consider a mixture of CO and NO2 at 700 K and a concentration of 0.10 mol/L
every molecule would collide with about a billion other molecules in one second!
if every collision resulted in a reaction, then the whole mixture would be reacted in a fraction of a second.
the actual reaction takes about 20 seconds.

5. Effective Collisions
In order for collisions to be effective, there must be considerable force in the collisions. The slower moving molecules do not have enough kinetic energy to react when they collide...they bounce off one another and retain their identity.
Only those molecules moving at high speed have enough energy for collisions to result in a reaction.
Every reaction requires a certain minimum energy for the reaction to occur--it is called activation energy, Ea, and is expressed in kJ.

6. Activated Complex-an unstable, high energy species
forward reaction exothermic (∆H<0), Ea is smaller than Ea1
forward reaction endothermic (∆ H>0), Ea is larger than Ea1
if ∆ H = +200 kJ, then Ea = Ea1 + DH = Ea kJ

7. Factors that Affect Reaction Rates Collision Theory
1. Temperature
2. Concentration
3. Particle size
4. Catalyst

8. Factors that Affect Reaction Rates Collision Theory
Temperature
Increasing temperature increases the number of particles that have enough kinetic energy to react when they collide.
Concentration changes (amt per vol)
Cramming more particles into a fixed volume increases the collision frequency.

9. Factors that Affect Reaction Rates Collision Theory
Particle Size
the smaller the particle size, the larger the surface area for a given mass of particles. Decreasing particle size will increase the rate of reaction.
Catalyst
A catalyst is a substance that increases the rate of a reaction without being used up itself in the reaction.

10. Effects of Catalyst on Activation Energy
Enzymes are biological catalyst usually made of proteins.
Speed reactions by lowering the activation energy of the reaction.

11. Chemical Equilibrium Dynamic (in constant motion) Reversible
Chemical equilibrium occurs when the forward and reverse reaction are taking place at the same rate.
There is no net change in the actual amounts of the components of the system.

12. Equilibrium Example

13. Rate vs Equilibrium At Equilibrium: RATES ARE EQUAL
the concentrations of reactants and products are constant.
D [ ]’s = 0
The forward and reverse reactions continue after equilibrium is attained.

14. Kinetics Reaction Rate Orders
the order of reaction with respect to a certain reactant, is defined as the power to which its concentration term in the rate equation is raised.
For example, 2A + B → C
r = k[A]2[B]1 the reaction order with respect to A would be 2 and with respect to B would be 1, the total reaction order would be 2 + 1 = 3.
Reaction orders can be determined only by experiment. The reaction order is not necessarily related to the stoichiometry of the reaction, unless the reaction is elementary. Complex reactions may or may not have reaction orders equal to their stoichiometric coefficients

15. The Reaction Quotient, Q
In general, all reacting chemical systems are characterized by their REACTION QUOTIENT, Q.
When the system is at equilibrium, Q = K

16. Reactions Review
“Systems”: two reactions that differ only in direction
Any reversible reaction H2 + I2 ↔ 2HI
noted by the double arrow; ↔
Two reactions: only difference is the Direction
H2 + I2 ↔ 2HI
Reactant products
2HI ↔ H2 + I2
Left Right

17. Reversible Reactions H2 + I2 ↔ 2HI
the products may react back to original reactants.
“closed system”: ONLY if all reactant are present
If one piece is completely gone it has ”gone to competition” and no longer reversible
Examples: Reversible Reactions.
Unopened Soda
Breathing
Rechargeable batteries
Color changing shirt

18. Blue to pink Co(H2O)4Cl2 + 2 H2O  Co(H2O)6Cl2
The product is hexa-coordinated with water and the reactant is tetra-coordinated with water. Notice that the net charge on the left and right sides of the first equation is zero.
Inclusion of 4[H2O] assumes additional information that the reactant is hexa-coordinated with water and the product is tetra-coordinated with water. Notice that the net charge on the left and right sides of the first equation is zero.
Blue to pink Co(H2O)4Cl2 + 2 H2O  Co(H2O)6Cl2

19. Properties of an Equilibrium
Pink to blue
Co(H2O)6Cl2 → Co(H2O)4Cl H2O
Blue to pink
Co(H2O)4Cl H2O → Co(H2O)6Cl2
Equilibrium systems are
DYNAMIC (in constant motion)
REVERSIBLE
can be approached from either direction

20. Reaction at Equilibrium
A and B are _________. C and D are _______.

21. Factors Affecting Equilibrium
Changes in temperature, pressure, and concentration affect equilibrium.
The outcome is governed by LE CHÂTELIER’S PRINCIPLE
“...if a system at equilibrium is disturbed, the system tends to shift its equilibrium position to counter the effect of the disturbance.”

22. Writing and Manipulating Keq
Solids NEVER appear in equilibrium expressions.
S(s) + O2(g) ---> SO2(g)
Liquids NEVER appear in equilibrium expressions.
NH3(aq) + H2O(liq) ---> NH4+(aq) + OH-(aq)

23. Henri Le Châtelier Henri Le Châtelier 1850-1936
Studied mining engineering.
Interested in glass and ceramics.

24. A + B  C + D Change in Concentration
What happens when there is an increase in reactant or product?
What happens when there is a decrease in reactant or product?

25. Change in Concentration
A + B  C + D
Change in Concentration
Stress Shift
Increase in A or B forward
Increase in C or D reverse
Decrease in A or B reverse
Decrease in C or D forward

26. Changes in Pressure
Only affects equilibrium system with an unequal number of moles of gaseous reactants and products
Decrease in pressure shifts the reaction in the direction that produces the larger number of moles of gas.
An increase in pressure shifts the reaction in the direction that produces the smaller number of moles of gas.

27. Product or Reactant Favored?
Keq greater than 1 Products favored
Keq less than 1 Reactants favored

28. Endothermic vs Exothermic
Exothermic - Heat is a product. (-ΔH)
Endothermic - Heat is a reactant. (+ΔH)

29. Changes in Temperature
An increase in temperature favors endothermic reactions.
A decrease in temperature favors exothermic reactions.

30. Changes in Temperature
+ΔH = Endothermic
↑ temp favors the forward rxn.
-ΔH = Exothermic
↑ temp favors the reverse rxn.

31. Temperature Effects on Equilibrium
N2O4 (colorless) + heat NO2 (brown)

32. Equilibrium and Catalysts
Add catalyst = no change equilibrium concentration
A catalyst only affects the RATE of approach to equilibrium
Catalytic exhaust system.

33. Applying a Stress to a System at Equilibrium
N2 + 3H2  2NH3 + Heat
ΔH = -92 kJ/mol rxn
Increase temperature
Increase pressure
Add a catalyst No Change
Adding H2
Removing NH3

34. Work on Equilibrium Practice Problems

35. The Equilibrium Constant
For any type of chemical equilibrium of the type: a A + b B ↔ c C + d D Keq is a CONSTANT (at a given Temp)

36. Equilibrium Constant
N2 (g) + O2 (g) ↔ 2NO (g)

37. Equilibrium Constant
CH4 (g) + Cl2 (g) ↔ CH3Cl (g) + HCl (g)

38. Magnitude of Keq Varies only with temperature
Is constant at a given temperature
Is independent of the initial concentrations

39. Magnitude of Keq A + B ↔ C + D Keq > 1 mostly products
Keq < 1 mostly reactants
Keq ~ 1 equal amounts of products and reactants

40. The Meaning of Keq For N2(g) + 3 H2(g)  2 NH3(g)
Keq is used to tell if a reaction will favor products or reactants.
For N2(g) + 3 H2(g)  2 NH3(g)
Concentration of products is much greater than that of reactants at equilibrium.
The reaction strongly favors products

41. The Meaning of Keq Concentration of products is much less
AgCl(s) ↔ Ag+(aq) + Cl (aq)
Keq = 1.8 x 10-5
Concentration of
products is much less
than that of reactants at
equilibrium.
The reaction strongly
favors reactants

42. The Meaning of Keq
AgCl(s) ↔ Ag+(aq) + Cl (aq) Keq = 1.8 x 10-5 Reactant Favored The Reverse reaction Ag+(aq) + Cl-(aq) ↔ AgCl(s) Keq = 1.8 x 105 is product-favored.

43. Calculation of Keq – Learning Check
For the reaction
2HI(g) ↔ H2(g) +I2(g), at 448•C,
The equilibrium concentrations are
HI = M,
H2 = M, and
I2 = M.
Calculate the equilibrium constant at this temperature.
Write the equilibrium expression (Prod/React)
Fill in the values and solve

44. Calculation of Keq 2HI(g) ↔ H2(g) + I2(g),
Equil

46. Calculation of Keq Learning Check
N2 + 3H2  2NH3
When equilibrium is established in a L container the equilibrium concentrations are: N2, 3.01 mol; H2, mol; NH3, mol.
Calculate Keq
Write the equation for Keq
Calculate molar concentration (M)
Fill in and solve.

47. Calculation of Keq [N2] = 3.01 mol/ 5.00L [H2] = 2.10 mol/ 5.00L
N2 + 3H2  2NH3
[N2] = 3.01 mol/ 5.00L
[H2] = 2.10 mol/ 5.00L
[NH3] = .565 mol/ 5.00L
[N2] = .602 M
[H2] = .420 M
[NH3] = .113 M
We must convert
concentrations to
Molarity = mol/L

48. Calculation of Keq
N H2 ↔ 2NH3
Equil

49. Uses of Keq
We can us the Keq to find equilibrium concentrations at that temperature.

50. Factors that Affect Equilibria
Once a reaction has reached equilibrium, it remains at equilibrium until it is disturbed by some change in conditions.