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3/9  Objective: SWBAT determine the driving force of a

chemical reaction. Do Now: 1) Hand in work completed yesterday. 2) Using the models at your pod, construct the

reactants of the following equation: H2 + O2 à

HW – Unit 9 pg 23 Q 1  

Making  things  Happen  

Unit 9: Kinetics, Thermodynamics & Equilibrium Lesson 1: Kinetics

2  

What  causes  a  chemical  reac9on?  Kinetics:

The study of chemical reactions Mechanism:

How a reaction happens Think of reactant molecules…why do they react? 1.  They must collide with each other 2.  The collisions must have enough energy 3.  The molecules must have the proper orientation Effective Collisions:

Collisions of reactant that cause them to react. 3  

Reac9on  Mechanisms  Can be very complex. Here’s an easy one: Theoretical Reaction:

2A + B à A2B Reaction Mechanism:

A + B à AB + A àA2B

4  

Here’s  a  Real  One:  Reaction:

H2 + I2 à 2HI Step 1:

H2 à 2H –Absorbs energy Step 2:

I2 à 2I –Absorbs energy Step 3:

2H + 2I à 2HI—Releases energy Net Equation:

H2 + I2 à 2HI 5  

All  Road  Lead  To  Rome  Reactions can have different pathways depending on local conditions. Experimentation is needed to determine the reaction pathway. Rate-determining Step:

The slowest step in the reaction pathway. It determines the overall rate of the reaction.

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Factors  that  Affect  Reac9on  Rate  Reaction Rate:

How many reactions happen/Unit of time

NOT HOW FAST/SLOW THE REACTION IS!!! Reaction Rate can be affected in 2 ways: 1.  Changing the reaction mechanism. 2.  Changing the number of collisions between

reactants

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Changing  the  Reac9on  Rate  

8  

1.    Use  of  a  catalyst  Catalyst:

1. Changes one or more steps in a reaction (usually lowers the activation energy). 2. Shortens the reaction mechanism. 3. Is not consumed during the reaction.

Examples: – Biological enzymes – Catalytic converter in your car – Dissolving substances in water.

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2.    Use  of  an  Inhibitor  Inhibitor:

1. Changes one or more steps in a reaction (usually increases the activation energy).

2. Lengthens the reaction mechanism. 3. Is not consumed in the reaction

Examples: – Patina on the surface of a metal

preventing oxidation of the rest of the metal

– Painting your car 10  

Changing  the  Number  of  Collisions  

11  

1.    Nature  of  Reactants  Different substances react at different rates. Has to do with how much energy those substances need to react. Ionic reactions happen much faster than covalent reactions. Why?

In solution, ionic bonds are broken, covalent bonds aren’t.

It takes more energy to break the covalent bonds.

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2.    Temperature  Why? Hotter the substance = higher the average KE. Higher the average KE = higher the rate of effective collisions. Examples: – Putting batteries in the fridge. – Heating up your food to cook it.

Which has a higher rate of collisions? A car going 30mph or a car going 100mph?

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3.    Concentra9on  Why? Higher concentration = more of a substance per unit of area. More substance present = higher rate of effective collisions Increasing pressure on gases is the same thing as increasing concentration. Examples: – Dilute vs. Concentrated Acids.

Which has a higher rate of collisions? A parking lot with 10 cars or 100 cars?

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4.    Surface  area  Why? Higher surface area = more surface to collide with other reactants = higher rate of effective collisions. Only effects solids. Examples: – Dissolving a sugar cube vs. dissolving powdered

sugar. – Using a shotgun vs. using a bb gun. – Crushed garlic flavors more than a whole clove.

Which has a higher rate of collisions? A parking lot with 10 cars or 100 cars?

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To  Review:  Reactions depend upon effective collisions. Effective collisions require the proper orientation and energy of reactants. Reactions happen according to reaction pathways. Anything that affects the number of effective collisions or the pathway of the reaction will effect the rate of the reaction.

16  

Take out a sheet of paper, calc.., and ref tables.

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CO + 2H2 à CH3OH

1.2 grams of H2 are made to react with 7.45 grams of CO. What is the limiting reactant? How much CH3OH should be produced? If we actually recover 7.52g of CH3OH, what is our % yield?

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%yield  =  moles  actual/moles  theore9cal    If we actually recover 7.52g of CH3OH, what is our

% yield?

19   20  

What  now?  

Any  Ques)ons?  

3/10  Objective: SWBAT calculate the effect of

temperature of reaction rate. Do Now: Describe ways to speed up the following

reaction: N2 + H2 à NH3

HW: Study for quiz on %yield and Collision Theory

Quiz

21   22  

23   24  

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3/23/11  

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25  

Alka  Seltzer  Lab  

26  

3/11  Objective: SWBAT construct a potential energy diagram given a chemical reaction. Do Now: Quiz. HW: Chapter 9 pgs. 26-27.

27  

Fire  and  Ice  

Unit 9: Kinetics, Thermodynamics & Equilibrium Lesson 2: Potential Energy Diagrams

28  

What  Happens  During  a  Reac9on?  We know it involves reactants à products. But what about energy? A reaction can do one of two things with energy: 1.  Absorb it (endothermic) 2.  Release it (exothermic)

29  

All  About  H  H = enthalpy:

A measurement of the amount of energy stored in a substance (units kJ/mole). For endothermic reactions ΔH is positive. For exothermic reactions ΔH is negative. ΔH values of many reactions are listed on Reference Table I.

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Endothermic  Reac9ons  General Formula:

A + B + Energy à C + D a.  The reactants absorb energy (heat) b.  This causes the temperature of the

surroundings to decrease. c.  The products have more energy than the

reactants (stored in their bonds). d.  The products are less stable than the reactants. This is how explosives are made!

31  

A  General  Example  For the reaction:

A + B à C HA = 40kJ HB = 20kJ Hc = 110kJ

How much energy is absorbed during this reaction? What is ΔH? Rewrite the equation to show the Conservation of Energy:

32  

A  Table  I  Example  N2 (g) + O2 (g) à 2NO (g) ΔH = +182.6 kJ/mole

(+) ΔH = Endothermic Reaction. This means that Nitrogen and Oxygen need to absorb 182.6 kJ of energy per mole of NO that will be formed. NO is an unstable compound.

33  

Exothermic  Reac9ons  General Formula:

A + B à C + D + energy a.  The reactants release energy (heat) b.  This causes the temperature of the surroundings

to increase. c.  The products have less energy than the reactants

(stored in their bonds). d.  The products are more stable than the reactants. This is what is left behind after an explosion!

34  

A  General  Example  For the reaction:

A + B à C HA = 60kJ HB = 40kJ Hc = 30kJ

How much energy is released during this reaction? What is ΔH? Rewrite the equation to show the Conservation of Energy:

35  

A  Table  I  Example  C (g) + O2 (g) à CO2 (g) ΔH = -393.5 kJ/mole

(-) ΔH = Exdothermic Reaction. This means that Carbon and Oxygen release 393.5 kJ of energy per mole of CO2 that will be formed. CO2 is a stable compound.

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3/23/11  

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Poten9al  Energy  Diagrams  A graph that shows what happens to potential energy as a reaction occurs. What has to happen for a reaction to occur? 1.  The reactants have some amount of energy

stored in their bonds. 2.  To get them to react, some energy is put in to

the reactants (the “activation energy”) to get them to collide effectively. PE goes up until the reactants become the “activated complex”.

3.  Once the activated complex reacts, the PE decreases to the energy of the products

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What  goes  in  a  PE  diagram  1.  Y-axis: PE (kJ) 2.  X-axis: Reaction coordinate (unmeasured time). 3.  H of reactants 4.  H of activated complex 5.  H of products. 6.  The Activation energy (arrow from H of

reactants to H of activated complex) 7.  ΔH: Arrow from (H of reactants to H of

products)

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An  Endothermic  PE  Diagram  

39  

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Note:  1. PE  increases.    Temperature  of  surroundings  will  decrease.  

2. The  Products  have  more  energy  stored  in  their  bonds  than  the  reactants  

An  Exothermic  PE  Diagram  

40  

Note:  1. PE  decreases.    Temperature  of  surroundings  will  increase.  

2. The  Products  have  less  energy  stored  in  their  bonds  than  the  reactants  

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The  Effect  of  A  Catalyst  Catalysts increase the rate of a reaction by lowering the activation energy. Less energy to get to the activated complex = higher frequency of effective collisions.

41  

Pgs.  24-­‐25  PRACTICE  

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43   44  

45  

Collision  Theory  POGIL  

46  

47  

What  now?  

Any  Ques)ons?  

3/14  Objective: SWBAT describe the properties of

chemical equilibrium. Do Now: 1) Hand in your PE Graph from HW.

2) Write your name on a separate sheet of paper, take out your reference tables, and a calculator. Close all other books.

HW: Ch 9 pgs 28-29 Questions.

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N2 (g) + 3 H2 (g) à 2 NH3

4.0 L of each reactant are used to produced 1.5g of NH3. 1.  Which is the limiting reagent?

2.  Calculate the percent yield.

49   50  

Pgs.  26-­‐27  Review  

51   52  

53  

⇋  Unit 9: Kinetics, Thermodynamics & Equilibrium Lesson 3: Equilibrium Systems 54  

3/23/11  

10  

Equilibrium  Strikes  Again  Equilibrium:

A state of rate balance between two opposing changes. The rate of the forward change is equal to the rate of the reverse change. Most reactions are reversible:

A + B ↔ C + D + energy The double arrow means that both reactions are occurring simultaneously

55  

Forward  and  Reverse  A Forward reaction:

N2 (g) + 3H2 (g) à 2NH3 (g) + 92kJ

A Reverse reaction: 2NH3 (g) + 92kJ à N2 (g) + 3H2 (g)

An Equilibrium:

N2 (g) + 3H2 (g) ↔ 2NH3 (g) + 92kJ

56  

Proper9es  of  Equilibrium  1.  Equilibrium is dynamic. The forward and reverse

reactions are occurring simultaneously. Since the rates of both are equal, it looks like nothing is happening, but don’t be fooled!

2.  Equilibrium can only happen in a closed system, isolated from its surroundings.

3.  Changing the conditions of the equilibrium system will change the equilibrium.

4.  Don’t fool yourself into thinking that just because the rates of the reactions are the same, that the concentrations of the reactants and products also have to be the same. They just have to be constant.

57  

Different  Equilibrium  Points  

58  

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Types  of  Equilibrium  1.  Chemical Equilibrium:

The rates of a forward and reverse reaction are equal.

2.  Solution Equilibrium: The rate of a solvent dissolving is equal to the rate of a dissolved solvent precipitating out of solution.

3.  Physical Equilibrium (aka “Phase Equilibrium”) The rate of a forward phase change is equal to the rate of a reverse phase change.

59  

Chemical  Equilibrium  Keq= the “equilibrium constant” The ratio of the amount of product to reactant at equilibrium Keq = [products]/[reactants]

Square brackets mean “concentration”

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The  Mass  Ac9on  Expression  The ratio of the concentrations of the products to the concentrations of the reactants. Tells us how far a reaction proceeded before reaching equilibrium. For a system:

aA + bB ↔ cC + dD

Keq = [C]c [D]d

[A]a [B]b

61  

Set  up  the  Mass  Ac9on  Expression!  

2SO2 (g) + O2 ↔ 2SO3 (g)

62  

Set  up  the  Mass  Ac9on  Expression!  

CaCl2 (s) ↔ Ca2+ (aq) + 2Cl- (aq)

63  

Set  up  the  Mass  Ac9on  Expression!  

2O3 (g) ↔ 3O2 (g)

64  

Keq  Keq is what we get when we solve the mass action expression. It has no units. It’s just a ratio. For a system:

A + B ↔ 2C + Energy The equilibrium concentrations are [A]=2M, [B]= 5M, and [C]=10M. What is Keq?

65  

What  does  Keq  mean  Keq tells us how much product was formed before the system reached equilibrium. There are 5 major points where Keq is informative

66  

Value  of  Keq   [Products]  :  [Reactants]   Meaning  0  <1  1  >1  ∞  

[Products]<<[Reactants]    [Products]  <  [Reactants]  [Products]  =  [Reactants]  [Products]  >  [Reactants]  [Products]>>[Reactants]  

Reac9on  didn’t  start  Reac9on  went  a  bit  

Reac9on  went  halfway  Reac9on  went  mostway  Reac9on  to  comple9on  

Bo-om  line:    Bigger  the  Keq,  the  closer  the  reac9on  gets  to  comple9on  before  reaching  equilibrium  

3/23/11  

12  

3/15  Objective: SWBAT calculate how chemical equilibrium can

be disturbed. Do Now: Find the Keq for the following reaction:

S2(g) + 2H2(g) 2H2S H2 = 2.16 M S2 = 0.30 M H2S = 0.50 M HW: Chapter 9 pg. 30 Questions.

67  

⇋  

68  

Solu9on  Equilibrium  There is only a certain amount of space for solute particles in a solution. Once that limit gets reached, the solution is “saturated” (more on that next unit) At the saturation point, every solvent molecule that enters solution means one has to precipitate out.

69  

Keq  works  for  solu9ons,  too!  It works exactly the same way that it does for chemical changes. Bottom line: The higher the Keq for a solution, the more solvent is dissolved.

70  

Ksp  Ksp- The solubility product constant Used for substances with very low solubility. We don’t divide by the amount of undissolved solute (it would make things ridiculously small). For a system

MA2 (s) ↔ M+ (aq) + 2A- (aq) Ksp = [M+] [A-]2

Bottom Line: The higher the Ksp, the more soluble the substance is.

71  

Honors  Reference  Table  A  A table of solubility product constants for nearly insoluble salts at 1 atm and 298 K

72  

3/23/11  

13  

Physical  Equilibrium  When two phase changes are in equilibrium. Happens at particular temperatures: Ex:

H2O (s) ↔ H2O at 273K

73  

Pgs.  28-­‐29  Review  

74  

75   76  

77  

Equilibrium  Demo  

78  

3/23/11  

14  

Bean  Simula9on  Lab  

79   80  

What  now?  

Any  Ques)ons?  

⇋  3/16  

Objective: SWBAT describe changes in entropy as favored or unfavored. Do Now: Did part I reach equilibrium? How do you know? HW: 1) Bring in design draft tomorrow. 2) Ch. 9 pg 31 Questions.

81  

Ready,  Set,  Go!  

Unit 9: Kinetics, Thermodynamics & Equilibrium Lesson 4: Will a Reaction Happen On Its Own…

82  

Spontaneity  Spontaneous Chemical Reactions:

Reactions that proceed on their own once initiated. The spontaneity of a reaction depends on two things: 1.  Enthalpy 2.  Entropy

83  

Enthalpy  (H)  Enthalpy:

The amount of heat energy stored in a system. Nature favors reactions that undergo a DECREASE in enthalpy (that release energy). Spontaneous reactions tend to release energy. Most exothermic reactions are spontaneous. Most endothermic reactions are non-spontaneous. Does a match keep burning once struck? Why?

84  

3/23/11  

15  

Entropy  (S)  Entropy:

The randomness (disorder) of a system. Nature favors reactions that undergo an INCREASE in entropy. As temperature increases, entropy increases. Entropy increases from solid àliquid àgas. Why?

85  

4  possible  situa9ons  1.  Exothermic reactions that increase entropy:

ALWAYS SPONTANEOUS 2.  Endothermic reactions that decrease entropy:

ALWAYS NONSPONTANEOUS 3.  Exothermic reactions that decrease entropy:

SPONTANEOUS AT LOW TEMPERATURES 4.  Endothermic reactions that increase entropy:

SPONTANEOUS AT HIGH TEMPERATURES

86  

Gibbs  Free  Energy  (G)  A measurement of the amount of energy available in a system to do work. Nature favors a decrease in free energy. If the change in free energy (ΔG) is negative in a system, the reaction will be spontaneous.

87  

Josiah  Willard  Gibbs  (1839  –  1903)  

ΔG  =  ΔH  –  (T  ΔS)    The Gibbs Free Energy Equation!

88  

ΔH   ΔS   ΔG   Spontaneous?  

-­‐   +   Always  nega9ve  

Always  

-­‐   -­‐   Nega9ve  if  T  is  low  

At  Low  Temp  

+   -­‐   Nega9ve  if  T  is  high  

At  High  Temp  

+   +   Never  nega9ve  

Never  

Fun  with  Gibbs  1. Calculating ΔG: For the fictitious reaction A + B à C + 30.0 kJ at 25oC (298 K): If ∆H for this reaction is –30.0 kJ and ∆S is –0.010 kJ/K, calculate ∆G and determine if this reaction is spontaneous or nonspontaneous.

89  

3/16  Objective: SWBAT apply the Gibbs free energy equation. Do Now: Is this reaction spontaneous? How do you know? Ammonia is synthesized from nitrogen and hydrogen gases at

a temperature of 475 degrees Celsius.  

N2(g)+3H2(g) --->2NH3(g)

If delta H = -92.2kJ and delta S= -0.1987kJ/K HW: Science Fair Design due tomorrow. Quiz tomorrow. Test next week.

90  

3/23/11  

16  

Using  Honors  Reference  Table  B    Table B- Energies of Formation of Compounds at 1 atm and 298K. Tells us about the spontaneity and ΔH associated with making compounds at specific conditions. For synthesis. Reverse values for decomposition.

91  

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+$##(!"#$%G()*(+)$,-.&)"((8!!!!D*M(&"(7FK,)C(

2:#FG"H-I" J4KLM" J4NO4"PQG"H0I" JMK" J4K";(RFM"H-I" J4ML#" J4GK4"S(HFQI#"H-I" J%ON" JO%O"SF#"H0I" JG%G" JG%M"SF"H0I" J44$D" J4GL"S?RFM"H-I" JLL4" JKK4"S#QK"H0I" JOM" JGG"S#QM"H0I" TN#" TKO"S#Q#"H0I" T##L" T#$%"QU"H0I" J#L4" J#LG"QV"H0I" T#K" T#"VS:"H0I" T4O" JN"W9F"H-I" J#4N" J4OO"'0F"H-I" JK$4" JNK%"PF"H0I" T%$D" TOL"PF#"H0I" TGG" TN4"XS:"H-I" JMGK" JM$%"P(S:"H-I" JM44" JGOM"RF#"H0I" J#%K" JG$$D"Q#F"H0I" J#M#" J##O"Q#F"H:I" J#ON" J#GL"

!

Determining  ΔS  Rearrange the equation to solve for ΔS:

ΔS = (ΔH-ΔG)/T Calculate ΔS for CO (use Table B values) Is it spontaneous always? Sometimes? Never?

92  

Calculate the ΔS for C2H4 Is it spontaneous always? Sometimes? Never?

93  

Calculate ΔS for ICl Is it spontaneous always? Never? Sometimes?

94  

Determine  the  Temperature  when  something  becomes  spontaneous  

The temperature at which a reaction becomes spontaneous is the equilibrium temperature. ΔG is 0 at this temperature.

ΔG = ΔH – (TΔS) 0 = ΔH – (TΔS)

TΔS =ΔH T=ΔH/ΔS

95  

Determine the temperature at which ICl reaches equilibrium

96  

3/23/11  

17  

3/18  Objective: SWBAT predict the direction an equilibrium will shift when stress is applied. Do Now: Quiz HW: Edited design due MondayUnit 9 Test next week.

97   98  

What  now?  

Any  Ques)ons?  

3/21  Objective: SWBAT predict the concentrations of reactants and products at equilibrium. Do Now: 1) Calculate the formula mass of AlBr3. Show your work. 2) Hand in Science Fair Designs. HW: pg 32 Questions. Chapter 9 Test Friday

99  

Like  a    SeeSaw  

Unit 9: Kinetics, Thermodynamics & Equilibrium Lesson 5: Changing Equilibrium 100  

Equilibrium  can  Shiq  LeChatelier’s Principle: If a system at equilibrium is subjected to a stress, the system will shift in a way that relieves the stress. This will cause a change in concentration of reactants and products until equilibrium is re-established.

101  

Henry  Louis                      Le  Chatelier  (1850  –  1936)  

Shiqing  Equilibrium  

“Forward Shift”: Shifting to favor the forward reaction (aka “shifting to the right”) “Reverse Shift”: Shifting to favor the reverse reaction (aka “shifting to the left”)

102  

Reactants  ↔  Products  Forward  Shiq  à  

ß  Reverse  Shiq  

3/23/11  

18  

A  Pile  of  Shiqs  You Must Learn These!

103  

Real  Big  Shiqs  

104  

N2  (g)  +  3H2(g)  ↔  2NH3  (g)  +  heat  

KNO3  (s)  +  34.89kJ  ↔  K+1  (aq)  +  NO3-­‐1(aq)  

What happens to [K+1] when the temperature is increased? What happens to [NO3

-1] when the temperature is decreased?

105  

Shiq!  

2CO  (g)  +  O2  (g)  ↔  2CO2  (g)  +  566kJ  

What happens to [CO2] when [CO] is added? What happens to [O2] when [CO2] is removed? What happens to [CO] when pressure is increased?

106  

Shiq!  

N2  (s)  +  2O2  +  66.4kJ  ↔  2NO(g)  List five things that can be done that will increase [NO]

107  

Shiqs!  

NaCl  (s)  +  3.88kJ  ↔  Na+(aq)  +  Cl-­‐(aq)  

List four things that can be done that will increase [NaCl]

108  

3/23/11  

19  

LeChatlier  Principle  POGIL  

109  

3/22  Objective: SWBAT predict how a given shift will affect dynamic equilibrium. Do Now: If the beaker started with 100 g of water, how would you know that the following reaction is at equilibrium?

H2O(l) ↔ H2O(g)

HW: Chap 9 pg 33 questions

110  

Pg.30  Review  

111  

112  

Pg  31  Q  Review  

113  

114  

3/23/11  

20  

115  

LeChatlier  Principle  POGIL  

116  

Systems  Thinking  Ac9vity  •  Make sure safety rules are on write-up

117  

118  

What  now?  

Any  Ques)ons?