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Ch. 12-13. Reaction Kinetics and Equilibrium. Reaction Kinetics. Looks at the reaction process and the factors that help us predict reactions. Stability. Thermodynamically Stable: Reaction does not spontaneously occur Kinetically Stable: Spontaneous

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ch 12 13

Ch. 12-13

Reaction Kinetics and Equilibrium

reaction kinetics
Reaction Kinetics
  • Looks at the reaction process and the factors that help us predict reactions
  • Thermodynamically Stable:
    • Reaction does not spontaneously occur
  • Kinetically Stable:
    • Spontaneous
    • Reaction is occurring so slow it is undetected ( but things are still reacting)
    • Ex. Decomposition of H2O2 (needs brown bottle)
reaction mechanism
Reaction Mechanism
  • Rxn occurs in a series of steps
  • Reaction Mechanism:
    • Series of reaction steps that must occur for a reaction to go to completion
    • Each step has 2 particles colliding
Ex. A + B -> C (step 1)

C + D -> E (step 2)

E + F -> G (step 3)

Total Rxn : A+ B + D + F  G

There were intermediates in between that you never saw (C, E)
  • you only see the original reactants and the final products
  • Intermediates:
    • Something that appears in the series but not in the final product
Ex. N2O  N2 + O (step 1)

N2O + O  N2 + O2 (step 2)

Total rxn. : 2N2O  2N2 + O2

(O was an intermediate)

Clock Reactions
  • Reaction Mechanism- teaching it
homogeneous reaction
Homogeneous Reaction
  • All reactant(s) and product(s) are in the same phase
heterogeneous reaction
Heterogeneous Reaction
  • Reaction that takes place at the interface between 2 phases
  • Zn (s) + HCl (aq)  H2 (g) + ZnCl2 (aq)

( HCl bubbles on the surface of the Zn)

collision theory
Collision Theory

In order for a rxn. to occur their particles must collide & those collisions must result in interactions

  • Collisions must:
    • Collide w/ enough energy
    • Have particles positioned in a way that enables them to react
    • Rate of reaction song
factors that affect reaction rates
Factors that affect reaction rates:
  • Nature of reaction:
    • Dependent on the type of bond involved
    • Ionic reaction rates, faster than covalent
  • Stirring:
    • Molecules in faster motion increase probability that the particles will hit & collide w/ enough energy
    • Smaller pieces increase the surface area so there are more possible sides for collisions LycopodiumSmall scale- creamer Mythbusters-creamer
  • Concentration: (video)
    • Quantity of matter that exists in a unit of volume
    • Increasing concentration increase # of collisions therefore increasing rate
    • Ex. Double the concentration  4x the collisions
Pressure (works for gases)
    • Increase pressure, decrease volume
    • So you have the same # of molecules in a smaller space, more molecules per unit of volume (i.e. higher concentration)  more collisions that could occur  increase rate
    • Measure of average kinetic energy (frequency of collisions)
    • Increase that frequency , the collisions increase
    • Increase temperature does 2 things:
      • Heating up molecules, moves them faster, more chances for collisions
      • More kinetic energy in molecules increase the motion of particles, easier to get over that activation energy, rate of reaction will increase
A “Cattle List”

(get it, a catalyst)

Ha, ha, ha, ha, ha, ha, ha,ha, ha, ha, ha (I crack myself up)

    • A chemical that increases the speed of the reaction but remains chemically unchanged
    • Doesn’t change the normal position of the equilibrium
    • Same amounts of product will be formed w/ or w/out the catalyst –just takes longer
    • Types: homogeneous & heterogeneous
    • Sugar/sulfuric
heterogeneous catalyst
Heterogeneous Catalyst
  • Surface catalyst
  • Ex. metal oxides, platinum
  • Works by adsorption – the adherence of one substance to the surface of another
  • Catalyst has specific lumps that hold the chemicals in the right position to react (increase the chance of them coming together)
Catalytic converter:
    • Platinum honeycomb structure (more surface area)
    • Pollution  SO2, CO2, NO
      • Converter lets H2O react w/ gases to convert them to weak acids (more complete combustion)
homogeneous catalyst
Homogeneous Catalyst
  • In same phase as reactants
  • Forms an intermediate or activated complex
  • Reactant reacts better w/ the catalyst than the other reactant
  • Ex. Sulfuric acid in ester reaction


Colbalt chloride

    • Chemical systems tend to achieve the lowest possible energy state (more stable)
    • Law of Disorder – states that things move spontaneously in the direction of maximum chaos
    • Entropy
      • Can be thought of as– measure of the disorder of the system or the randomness (more stable)


    • More exact definition- measure of the number of possible ways that the energy of a system can be distributed; related to the freedom of the system’s particles to move and the number of ways they can be arranged (energy dispersal)

Misconceptions about Entropy

  • This view of the second law of thermodynamics is very popular, and it has been misused. Some argue that the second law of thermodynamics means that a system can never become more orderly. Not true. It just means that in order to become more orderly (for entropy to decrease), you must transfer energy from somewhere outside the system, such as when a pregnant woman draws energy from food to cause the fertilized egg to become a complete baby, completely in line with the second line\'s provisions.
Entropy of gas is greater than liquid or solid
  • Entropy increase when a substance is divided into parts
  • Entropy increase w/ increase in temperature
    • Prevents reaction from happening for a certain length of time (delays reaction)
    • Not opposite of catalyst
    • Ex. Lemon juice on apples
    • A + B  AB
      • w/ inhibitor: A-inh + B  no rxn.
      • Once inhibitor is used up then: A + B  AB
energy diagrams
Energy Diagrams
  • Activation Energy:
    • Energy required to start a chemical reaction
    • High activation energy  few collisions have enough energy for a reaction  get slow undetected reaction
  • Activated Complex:
    • Product formed when reactants have collided w/ sufficient energy to meet activation energy requirement
energy diagram exothermic rxn
Energy Diagram: Exothermic Rxn

-releases energy, lower energy after rxn.

energy diagram endothermic rxn
Energy diagram: Endothermic Rxn.

-absorbed energy, higher energy after rxn.

energy diagram catalyst
Energy Diagram: Catalyst

-w/ catalyst product formed faster

-lowers the activation energy requirement

reaction rate
Reaction Rate:
  • Rate of disappearance of one of the reactants or rate of appearance of one of the products
  • Unit: (mole/L)/s
    • Change in molarity per second
rate law
Rate Law:
  • Rate is dependent on the concentration of the reactants
  • Expression relating the rates of reaction to the concentration of reactants
  • [ ] = concentration
a b ab
A + B  AB
  • Rate = k [A] [B]
  • k= specific rate constant (proportionality constant relating to concentration – value changes depending on rxn.)
Ex. H2 + I2 => 2HI

rate= k [H2 ] [I2 ]

Exp. 1- [H2 ] = 1.0M

[I2 ] = 1.0M

rate= .20 M/s


.20= k [1.0M] [1.0M]

k = .20

Exp. 2 - [H2] = .5 M [I2] = .5 M

rate = ?

k= .20

rate= k [H2 ] [I2 ]

rate = .20 [.5] [.5]

rate= .05 M/s


The rate law for elementary reactions is just the product of the reactants, reactions that have more than one step you would need to figure out the order of reaction.

rate determining step
Rate Determining Step
  • The step or reaction in the series that is slower than all the others  the reaction rate is dependent on this

Ex. Person going 45 in the left lane on I-94

reaction order or order of reaction
Reaction Order or Order of Reaction
  • Changing the concentration of substances taking part in a reaction usually changes the rate of reaction
  • A rate equation shows this effect mathematically
  • Orders of reaction are a part of this rate equation (helps us describe the reaction )
  • Orders of Reaction are always found by doing experiments
elementary reactions
Elementary Reactions
  • A reaction with no intermediate steps (very rare) – not a reliable way to determine order
  • One can determine the order with the coefficients
  • Rate is proportional to the concentration of the reactants raised to the power of the coefficients

Rate expressed as:

aA + bB cC + dD

Rate = k [A]a [B]b

( a and b are the coefficients)

reaction order
Reaction Order
  • Can determine reaction order experimentally or graphically
  • Experimentally:
  • Gather data and see what happen to rates if you change the concentration (1st order- double [ ] doubles rate, 2nd order – double[ ] quadruples rate, zero order- rate constant with any [ ] )
  • Plot concentration vs. time – identify which graph gives you a linear graph
    • Zero Order: Linear Graph [A] vs time
    • 1st order: Linear graph ln[A] vs time
    • 2nd order: Linear graph 1/[A] vs. time
Sum of the power to which all the reactant concentrations are raised (always defined in terms of reactant concentrations (no products))
  • Overall order = a + b

(exponents added together)

finding overall order
Finding overall Order

Ex. Rate = k [A] [B]2

Rate is 1st order for reactant A

Rate is 2nd order for reactant B

Overall order =(a + b) = 3rd order

-if you double [A] = doubles rate

-if you double [B] = quad. rate

practice problems
Practice Problems
  • Rate Law:
  • Reaction:
    • 2NO(g) + Cl2(g) 2 NOCl (g)

What is the rate law?

  • Rate = k [NO]2 [Cl2]
  • What is the order of the reaction with respect to NO?
  • 2nd order
  • What is the order of the reaction with respect to Cl2
  • 1st order
Using the data and rate law, calculate the rate constant.
  • k = 180
  • Assign: p 567 #21, #27 a,b #30 a, 39, 65 a,b


-use for reversible reactions

  • Means a state of balance
  • We will look at dynamic equilibrium – where changes are taking place but the overall balance is maintained (happening at the molecular level)
  • Not Static equilibrium- where nothing is moving any more (see-saw)
Use equilibrium w/ reversible reactions, where reactants convert to products and products convert to reactants simultaneously
  • Ex. Equilibrium reaction

2SO2(g) + O2(g) ↔ 2SO3 (g)

Can also use these symbols


2so 2 g o 2 g 2so 3 g
2SO2(g) + O2(g) ↔ 2SO3 (g)


  • In a reversible rxn., the rate of the reverse process is zero at the beginning. At that point no products are going back to reactants.
  • As the concentration of products build up, some products start converting to reactants (reverse rxn. starts)
2so 2 g o 2 g 2so 3 g1
2SO2(g) + O2(g) ↔ 2SO3 (g)
  • As reactants are used up their concentration decreases (forward reaction slows down)
  • As products build up, their concentration increases (reverse reaction speeds up)
  • Eventually the products are going to reactants at the same rate as the reactants are going to products; the rxn. has reached equilibrium
chemical equilibrium
Chemical Equilibrium
  • Forward & reverse reactions are taking place at same rate (no net change in actual amts. of products or reactants in the system)
2so 2 g o 2 g 2so 3 g2
2SO2(g) + O2(g) ↔ 2SO3 (g)











(Have twice as much SO2 as O2 initially, then a mixture of 3 gases is obtained at equilibrium)

equilibrium position
Equilibrium Position
  • Given by relative concentration of reactants & products at equilibrium
  • Doesn’t mean exactly 50%/50% concentration at equilibrium
  • The position indicates what is favored at equilibrium
A B= product bond is weak and you have mostly reactants at equil.

A B = product bond is strong and you have mostly products at equil.

(larger arrow indicates the favored direction)

Catalyst speed up the forward and reverse reactions equally (the activation energy is reduced by the same amount)
  • Catalysts don’t affect the amts. of reactants or products present at equil. ( just decrease the time to get to equil.)
equilibrium constant keq
Equilibrium Constant (Keq)
  • Use of constant is a concise way of stating whether reactants or products are favored in a rxn.
  • The constant #’s relate the amt. of reactants to products at equil.
Keq show the ratio of products to reactants
  • If Keq > 1 products favored at equil. (spontaneous rxn)
  • If Keq < 1 reactants favored at equil. (non-spontaneous rxn)
Ex. aA + bB  cC + dD

(a) moles of reactant A react w/ (b) moles of reactant B and give (c) moles of product C and (d) moles of product D

Keq (equil. constant)=

ratio of product concentrations to reactant concentrations w/ each concentration raised to a power given by the # of moles of that substance in the balanced chem. rxn.

-Keq is dependent on temp., as temp changes Keq changes

aa bb cc dd
aA + bB  cC + dD

Keq = [C]c x [D]d

[A]a x [B]b

h 2 i 2 2hi
H2 + I2  2HI

Write the Keq equation for this.

Keq = [HI]2

[H2] [I2]

n 2 o 4 g 2 no 2 g
N2O4 (g)  2 NO2 (g)
  • This is a homogeneous equil.- all substances are in same phase
  • Write Keq equation:

Keq= [NO2]2


n 2 o 4 g 2 no 2 g1
N2O4 (g)  2 NO2 (g)
  • Calculate the Keq if:

[NO2] = .0045 mol/L

[N2O4] = .030 mol/L

Keq= [NO2]2


Keq = [.0045]2


Keq = 6.8 x 10 -4

keq 6 8 x 10 4
Keq = 6.8 x 10 -4
  • What is favored at equilibrium?

Reactants (Keq<1)

n 2 g o 2 g 2no g
N2 (g) + O2(g)  2NO(g)
  • Write the Keq

Keq = [NO]2

[N2] [O2]

n 2 g o 2 g 2no g1
N2 (g) + O2(g)  2NO(g)
  • If [N2] and [O2] = .72 M and Keq = 4.6 x 10 –31. What is [NO]?

4.6 x 10-31 = [NO]2

[.72] [ .72]

2.38 x 10 –31 = [NO]2

4.9 x 10 –16 M = [NO]

book problems
Book problems
  • P. 614 (new book)
  • # 17, 21,22,23,37
  • P 587 #24-26 (old book)
le chatelier s principle demo
Le Chatelier’s Principledemo
  • Delicate balance exists between reactants and products in a system at equilibrium
  • If equil. conditions are changed – system shifts to restore equilibrium
  • Any application of stress to the system disrupts the system
  • Henri LeChatelier studied the changes in systems w/ stress application
LeChatelier’s Principle =

-if a stress is applied to a system in a dynamic equilibrium, the system changes to relieve the stress

-stress types: concentration, temperature, pressure

Changes in Concentration
    • Change amounts of reactants or products
    • System changes to minimize the original change
ex co 2 h 2 0 h 2 co 3
Ex. CO2 + H20 H2CO3
  • Adding a reactant always pushes a reversible reaction in the direction of the products
    • Shifts reaction to right ()
    • Forms more product, uses up excess reactants
ex co 2 h 2 0 h 2 co 31
Ex. CO2 + H20 H2CO3
  • Removing a reactant always pulls a reversible rxn in the direction of the reactants
    • Shifts reaction to the left ()
    • Forms more reactants (less products)
ex co 2 h 2 0 h 2 co 32
Ex. CO2 + H20 H2CO3
  • If product is added at equil. the reaction shifts to the formation of more reactants
    • Shifts to the left ()
    • Forms less product, more reactant
If product is removed at equil. reaction shifts to direction of products
    • Shifts to the right ()
    • Forms more product, less reactant
    • Removal is a trick used by chemists to increase the yield of a desired product (ex. Hens & eggs)
Change in temperature
    • Increasing the temperature causes the equilibrium position of a reaction to shift in a direction that absorbs the heat energy
    • Keq changes if temp. changes
    • Reversible rxns are endothermic in one direction, exothermic in the other
    • Effect of a change depends on which is endo
In a reaction

-addition of heat  favors endo side

-removal of heat  favors exo side

2so 2 o 2 2so 3 heat kcal
2SO2 + O2<-> 2SO3 + heat (kcal)
  • Think of heat as a product
  •  exo direction(points towards heat)
  •  endo direction (points away from heat)
  • If I add heat– shifts away from heat to cool system (shifts left-toward reactants)
  • If I cool – shifts towards heat (shifts right-toward products)
Change In Pressure
    • Affects gas phase only
    • Affects the number of moles
    • Similar effect as increasing concentration of any gas
    • Le Chat. States that if the pressure on an equilibrium system is changed the rxn. is driven in a direction that relieves that stress
An increase in pressure (decreases volume) favors the side w/ the least moles
  • A decrease in pressure (increases volume) favors the side w/ more moles
  • If moles are equal on both sides- no pressure effect
pressure increase
Pressure Increase

PCl5 (g) <-> PCl3 (g) + Cl2(g)

1 mole 2 moles

  • favors side w/ least moles
  • Shifts toward left (reactants)
  • PCl5 PCl3  Cl2
pressure decrease
Pressure decrease

2PbS(s) +3O2(g)<->2PbO(s) +2SO2(g)

-look only at gases

3 moles 2 moles

-favors side w/ most moles

-shift toward left (reactants)

O2 SO2

  • 18 multiple choice
  • 4 short answer
  • 2 calculations
  • 3 essay