Ch. 12-13

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Ch. 12-13

• Reaction Kinetics and Equilibrium

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Reaction Kinetics

• Looks at the reaction process and the factors
  that help us predict reactions

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Stability

• 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)

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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

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• Ex. A B -gt C (step 1)
• C D -gt E (step 2)
• E F -gt G (step 3)
• Total Rxn A B D F ? G

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• 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

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• Ex. N2O ? N2 O (step 1)
• N2O O ? N2 O2 (step 2)
• Total rxn. 2N2O ? 2N2 O2
• (O was an intermediate)

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• Clock Reactions
• Reaction Mechanism- teaching it

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Homogeneous Reaction

• All reactant(s) and product(s) are in the same
  phase

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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)

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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

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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

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• Crushing
• Smaller pieces increase the surface area so there
  are more possible sides for collisions
  Lycopodium Small 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

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• 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

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• Temperature
• 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

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Commercial Break What is this?
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• A Cattle List
• (get it, a catalyst)
• Ha, ha, ha, ha, ha, ha, ha,ha, ha, ha, ha (I
  crack myself up)

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• Catalyst
• A chemical that increases the speed of the
  reaction but remains chemically unchanged
• Doesnt 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

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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)

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• 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)

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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
• enzymes
• Colbalt chloride

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• Entropy
• 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)

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• Entropy
• 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
  systems particles to move and the number of ways
  they can be arranged (energy dispersal)

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• 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.

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• Entropy of gas is greater than liquid or solid
• Entropy increase when a substance is divided into
  parts
• Entropy increase w/ increase in temperature

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• Inhibitors
• 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

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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

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Energy Diagram Exothermic Rxn

• -releases energy, lower energy after rxn.

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Energy diagram Endothermic Rxn.

• -absorbed energy, higher energy after rxn.

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• Endo thermic/exothemic song

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Energy Diagram Catalyst

• -w/ catalyst product formed faster
• -lowers the activation energy requirement

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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

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• Reaction rate song (second time)

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Rate Law

• Rate is dependent on the concentration of the
  reactants
• Expression relating the rates of reaction to the
  concentration of reactants
• concentration

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A B ? AB

• Rate k A B
• k specific rate constant (proportionality
  constant relating to concentration value
  changes depending on rxn.)

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• Ex. H2 I2 gt 2HI
• rate k H2 I2
• Exp. 1- H2 1.0M
• I2 1.0M
• rate .20 M/s
• k?
• .20 k 1.0M 1.0M
• k .20

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• Exp. 2 - H2 .5 M I2 .5 M
• rate ?
• k .20
• rate k H2 I2
• rate .20 .5 .5
• rate .05 M/s

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• 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.

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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

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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

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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 Aa Bb
• ( a and b are the coefficients)

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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 )

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• Graphically
• Plot concentration vs. time identify which
  graph gives you a linear graph
• Zero Order Linear Graph A vs time
• 1st order Linear graph lnA vs time
• 2nd order Linear graph 1/A vs. time

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• 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)

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Finding overall Order

• Ex. Rate k A B2
• 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

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Practice Problems

• Rate Law
• Reaction
• 2NO(g) Cl2(g) ?2 NOCl (g)

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• Using the following data, calculate the rate law
  and constant.

NO Cl2 Rate (? /?t)
0.10 0.10 0.18
0.10 0.20 0.36
0.20 0.20 1.45
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• What is the rate law?
• Rate k NO2 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

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• 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

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Equilibrium

• -use for reversible reactions

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Equilibrium

• 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)

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• 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
• ?,

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2SO2(g) O2(g) ? 2SO3 (g)

• Steps
• 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)

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2SO2(g) O2(g) ? 2SO3 (g)

1. As reactants are used up their concentration
   decreases (forward reaction slows down)
2. As products build up, their concentration
   increases (reverse reaction speeds up)
3. Eventually the products are going to reactants at
   the same rate as the reactants are going to
   products the rxn. has reached equilibrium

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Chemical Equilibrium

• Forward reverse reactions are taking place at
  same rate (no net change in actual amts. of
  products or reactants in the system)

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2SO2(g) O2(g) ? 2SO3 (g)
CONCENTRATION

• 100
• 75
• 50
• 25
• 0

SO3
SO2
O2
TIME ?
(Have twice as much SO2 as O2 initially, then a
mixture of 3 gases is obtained at equilibrium)
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Equilibrium Position

• Given by relative concentration of reactants
  products at equilibrium
• Doesnt mean exactly 50/50 concentration at
  equilibrium
• The position indicates what is favored at
  equilibrium

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• 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)

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• Catalyst speed up the forward and reverse
  reactions equally (the activation energy is
  reduced by the same amount)
• Catalysts dont affect the amts. of reactants or
  products present at equil. ( just decrease the
  time to get to equil.)

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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.

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• Keq show the ratio of products to reactants
• If Keq gt 1 products favored at equil.
  (spontaneous rxn)
• If Keq lt 1 reactants favored at equil.
  (non-spontaneous rxn)

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• 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

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• 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

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aA bB ? cC dD

• Keq Cc x Dd
• Aa x Bb

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H2 I2 ? 2HI

• Write the Keq equation for this.
• Keq HI2
• H2 I2

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N2O4 (g) ? 2 NO2 (g)

• This is a homogeneous equil.- all substances are
  in same phase
• Write Keq equation
• Keq NO22
• N2O4

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N2O4 (g) ? 2 NO2 (g)

• Calculate the Keq if
• NO2 .0045 mol/L
• N2O4 .030 mol/L
• Keq NO22
• N2O4
• Keq .00452
• .030
• Keq 6.8 x 10 -4

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Keq 6.8 x 10 -4

• What is favored at equilibrium?
• Reactants (Keqlt1)

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N2 (g) O2(g) ? 2NO(g)

• Write the Keq
• Keq NO2
• N2 O2

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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 NO2
• .72 .72
• 2.38 x 10 31 NO2
• 4.9 x 10 16 M NO

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Book problems

• P. 614 (new book)
• 17, 21,22,23,37
• P 587 24-26 (old book)

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Le Chateliers Principle demo

• 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

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• LeChateliers 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

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• Changes in Concentration
• Change amounts of reactants or products
• System changes to minimize the original change

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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

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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)

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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

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• 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)

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• 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

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• In a reaction
• -addition of heat ? favors endo side
• -removal of heat ? favors exo side

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2SO2 O2lt-gt 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)

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• 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

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• 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

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Pressure Increase

• PCl5 (g) lt-gt PCl3 (g) Cl2(g)
• 1 mole 2 moles
• favors side w/ least moles
• Shifts toward left (reactants)
• ? PCl5 ? PCl3 ?? Cl2

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Pressure decrease

• 2PbS(s) 3O2(g)lt-gt2PbO(s) 2SO2(g)
• -look only at gases
• 3 moles 2 moles
• -favors side w/ most moles
• -shift toward left (reactants)
• ?O2 ?SO2

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test

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