Kinetics and Equilibrium Slide 3 / 119 Slide 4 / 119 Kinetics - - PDF document

SLIDE 1

Slide 1 / 119 Slide 2 / 119

Kinetics and Equilibrium

Slide 3 / 119

The reaction that occurs in an airbag must happen in 0.030 seconds. 2NaN3(s) --> 2Na(s) + 3N2(g)

Think About It... Slide 4 / 119 Kinetics

In kinetics we study the rate at which a chemical process occurs. Besides the speed at which reactions occur, kinetics also sheds light on the reaction mechanism (exactly how the reaction

• ccurs).

This is similar to the first chapter in 9th grade physics: kinematics, which also had to do with the rate of change.

Slide 5 / 119 Factors That Affect Reaction Rates

The Physical State of the Reactants The Concentrations of the Reactants Temperature The Presence of a Catalyst Surface area

Slide 6 / 119

In order to react, molecules must come in contact with each

• ther.

The more homogeneous the mixture of reactants, the faster the molecules can react.

The Physical State of the Reactants

This proves to us something that is fairly fundamental to our understanding of Chemistry. In order for the particles of two substances to react these particles need to be in actual physical proximity to one another.

SLIDE 2

Slide 7 / 119

As the concentration of reactants increases, so does the likelihood that reactant molecules will collide.

Air hole

The flame from the Bunsen burner shows the variation in color as more air is mixed with the natural gas. As more air is mixed with the natural gas, the rate of combustion reaction will increase and the flame gets more blue in color. Recall that blue light has more energy than red. Since the rate of reaction increases, the rate of energy being released also increases, hence the blue color.

The Concentrations of the Reactants Slide 8 / 119

At higher temperatures, reactant molecules: · have more kinetic energy · move faster · collide both more

• ften and with greater

energy

Temperature Slide 9 / 119

Catalysts speed up reactions by changing the mechanism of the reaction. Catalysts are very specific in their direction and reaction

• involved. A catalyst will

not effect the reactions equilibrium, only the speed at which that equilibrium is reached.

The Presence of a Catalyst Slide 10 / 119

Catalysts are not consumed during the course of the reaction. If they are chemically changed at some point in the reaction they are changed back to their

• riginal state during

some other part of the reaction.

The Presence of a Catalyst

Palladium is a rare metal and is more expensive than

• Platinum. However, because it is not "used up" in a

reaction, a single piece of palladium can be used again and again (without wasting this VERY expensive metal) as a catalyst in certain organic reactions.

Slide 11 / 119

When more surface area of a reactant is exposed, a reaction will

• ccur faster.

More surface area allows for more interaction between particles. A steel bar will not ignite if exposed to a flame. In contrast, a steel wool pad will burn and ignite quickly due to the steel being able to access oxygen more readily.

Surface Area

click here to watch steel wool burn

Slide 12 / 119

1 Why does a higher temperature cause a reaction to go faster?

A Only because there are more collisions per second. B Only because collisions occur with greater energy. C There are more frequent collisions and they are of greater energy.

answer

SLIDE 3

Slide 13 / 119

2 Why does a higher reactant concentration make a reaction faster?

A Only because there are more collisions per second. B Only because collisions occur with greater energy. C There are more frequent collisions and they are of greater energy.

answer

Slide 14 / 119

3 Why does a greater surface area cause a reaction to proceed faster?

A Because there are more collisions per second. B Because collisions occur with greater energy. C There are more frequent collisions and they are of greater energy.

answer

Slide 15 / 119

4 What happens to a catalyst in a reaction?

A It is unchanged. B It is incorporated into the products. C It is incorporated into the reactants. D It evaporates away.

answer

Slide 16 / 119 Reaction Rates

Rates of reactions can be determined by monitoring the change in concentration of either reactants or products as a function of time.

0 s 20 s 40 s 1 mol A 0.5 mol A 0.5 mol B

0.25 mol A 0.75 mol B

A B

A --> B

Slide 17 / 119

In this reaction, the concentration

• f butyl chloride,

C4H9Cl, was measured at various times.

Reaction Rates

C4H9Cl(aq) + H2O(l) --> C4H9OH(aq) + HCl(aq)

Time , t(s) [C

4H9Cl](M)

0.0 0.1000 50 0.0905 00.0 0.0820 50.0 0.0741 200.0 0.0671 300.0 0.0549 400.0 0.0448 500.0 0.0368 800.0 0.0200 10000 0

Slide 18 / 119

The average rate of the reaction over each interval is the change in concentration divided by the change in time: [C4H9OH] t average rate =

Reaction Rates

This is the same as velocity in physics;

• r the first derivative in calculus.

Time , t(s) [C

4H9Cl](M) Average Rate (M/s)

0.0 0.1000 50 0.0905 1.9 x10

• 4

00.0 0.0820 1.7 x10

• 4
• 50. 0 0.0741 1.6 x10
• 4
• 200. 0 0.0671 1.4 x10
• 4

300.0 0.0549 1.22 x 10

• 4

400.0 0.0448 1.01 x10

• 4

500.0 0.0368 0.88 x10

• 4

800.0 0.0200 0.560 x10

• 4

10000 0

C4H9Cl(aq) + H2O(l) --> C4H9OH(aq) + HCl(aq)

SLIDE 4

Slide 19 / 119

Note that the average rate decreases as the reaction proceeds. This is because as the reaction goes forward, there are fewer collisions between reactant molecules.

Reaction Rates

This is the same as a negative acceleration in physics;

• r a negative second derivative in calculus.

Time , t(s) [C

4H9Cl](M) Average Rate (M/s)

0.0 0.1000 50 0.0905 1.9 x10

• 4

00.0 0.0820 1.7 x10

• 4
• 50. 0 0.0741 1.6 x10
• 4
• 200. 0 0.0671 1.4 x10
• 4

300.0 0.0549 1.22 x 10

• 4

400.0 0.0448 1.01 x10

• 4

500.0 0.0368 0.88 x10

• 4

800.0 0.0200 0.560 x10

• 4

10000 0

C4H9Cl(aq) + H2O(l) --> C4H9OH(aq) + HCl(aq)

Slide 20 / 119

A plot of [C

4H9Cl] vs. time

for this reaction yields a curve like this. The slope of a line tangent to the curve at any point is the instantaneous rate at that time. In calculus this slope is called the derivative.

Reaction Rates

The slope of the line on the velocity vs. time curve is acceleration in physics; The tangent to the curve of the first derivative of a function is the second derivative in calculus.

Instantaneous rate at t=600s

[C4H9Cl] (M) x 10

• 2

Time (s)

Instantaneous rate at t=0s

C4H9Cl(aq) + H2O(l) --> C4H9OH(aq) + HCl(aq)

Slide 21 / 119

All reactions slow down

• ver time.

Therefore, the best indicator of the rate of a reaction is the instantaneous rate near the beginning of the reaction. At this time we assume we are looking at the rate of the forward reaction unneffected by the reverse reaction(s)

Reaction Rates

Instantaneous rate at t=600s

[C4H9Cl] (M) x 10

• 2

Time (s)

Instantaneous rate at t=0s

C4H9Cl(aq) + H2O(l) --> C4H9OH(aq) + HCl(aq)

Slide 22 / 119 Reaction Rates and Stoichiometry

In this reaction, the ratio of C4H9Cl to C

4H9OH is 1:1.

Thus, the rate of disappearance of C

4H9Cl is

the same as the rate of appearance of C

4H9OH.

Instantaneous rate at t=600s

[C4H9Cl] (M) x 10

• 2

Time (s)

Instantaneous rate at t=0s

C4H9Cl(aq) + H2O(l) --> C4H9OH(aq) + HCl(aq) The symbol indicates there is a small change being measured.

Slide 23 / 119

5 Which expression represents a reaction rate?

A time / mass B concentration / time C energy / time D time / energy

answer

Slide 24 / 119

6 Which set of units represents a reaction rate?

A s/ g B M/ s C s D J/ s E s/ J

answer

SLIDE 5

Slide 25 / 119 Concentration and Rate

One can gain information about the rate of a reaction by seeing how the rate changes with changes in concentration.

Slide 26 / 119 Slide 27 / 119 Slide 28 / 119 Slide 29 / 119 Rate Laws

A rate law shows the relationship between the reaction rate and the concentrations of reactants. The exponents tell the order of the reaction with respect to each reactant.

Slide 30 / 119 Rate Laws

the reaction is: First-order with respect to [NH4 +] and First-order with respect to [NO2 −]. Rate = k [NH4 +] [NO2 −]

SLIDE 6

Slide 31 / 119 Rate Laws

Rate = k [NH4 +](1)[NO2 −](1) The overall reaction order can be found by adding the exponents on the reactants in the rate law. 1+1=2 This reaction is second-order overall. This means that if the concentration of either reactant is doubled, the reaction rate will double. If the concentration of both reactants is doubled; the rate of reaction will quadruple.

Slide 32 / 119 Rate Laws

Rate = k [NO2]2 This means that if the concentration of NO2 is doubled; the rate of the reaction will quadruple. 2NO2(g) N 2O4(g) This reaction is second order in NO2 and second order overall.

Slide 33 / 119 Rate Laws

2NH3(g) N2(g) + 3 H2(g) This reaction is second order in NH3 and second order overall. Rate = k [NH3]2 This means that if the concentration of NH3 is doubled; the rate of the reaction will quadruple.

Slide 34 / 119

The decomposition of N2 O5 : 2N2 O5 --> 4NO2 + O2 The above reaction is first order with respect to N2 O5 *Note that the coefficients does not relate to the order of the reaction

• f a particular reactant. The order is determined experimentally only.

Rate Laws Slide 35 / 119 Rate Laws

2NO(g) + 2H2 (g) --> N2 (g) + 2H2 O(g) The reaction is first order with respect to H2 and second order with respect to NO. Even though two molecules/moles of each NO and H2 are reacting in the above reaction, the reaction is not 4th order overall. Recall the coefficients are not related directly to the order of a reactant or reaction.The order of a reactant or reaction is determined only experimentally. You cannot predict it without

• bserving, in some way, the actual physical manifestation of the

interactions of particles in the reaction.

Slide 36 / 119 Rate Laws

This means that if the concentration of NO is doubled; the rate of the reaction will quadruple If the concentration of H2 is doubled; the rate of the reaction will double If the concentrations of both reactants is doubled; the rate will increase 8x

2NO(g) + 2H2(g) --> N2(g) + 2H2O(g)

The reaction is first order with respect to H2 and second order with respect to NO

Rate = k [NO]2 [H2]

SLIDE 7

Slide 37 / 119 Temperature and Rate

Generally, as temperature increases, so does the reaction

• rate. This is because k is

temperature dependent.

Temperature ℃

x10-3 k(s-1)

The rate constant K will change only when the temperature changes

Slide 38 / 119

7 What will happen to the rate of a reaction if the molarity of a first-order reactant is doubled?

A rate will double B rate will triple C rate will quadruple D rate remains unchanged

answer

Slide 39 / 119

8 What will happen to the rate of a reaction if the molarity of a second-order reactant is doubled?

A rate will double B rate will triple C rate will quadruple D rate remains unchanged

answer

Slide 40 / 119 The Collision Model

In a chemical reaction, bonds are broken and new bonds are

• formed. Molecules can only react if they collide with each
• ther.

Cl + ClON ⇒

Cl2 + NO

Before collision Collision After collision

Furthermore, molecules must collide with the correct orientation and with enough energy to cause bond breakage and formation.

Slide 41 / 119 The Collision Model

We can think of this like meeting people in a crowd. In order to meet someone, you have to first shake hands with them. In

• rder to do that you must both be facing in the correct
• direction. If either of you has their back turned to the other,

you might collide, but you will not meet. Cl + ClON ⇒

Cl + ClON

Before collision Collision After collision NO reaction

Similarly, if the orientation of the molecules does not support the reaction, no reaction will occur.

Slide 42 / 119 Activation Energy

In other words, there is a minimum amount of energy required for reaction: the activation energy, Ea . Just as a ball cannot get

• ver a hill if it does not

roll up the hill with enough energy, a reaction cannot occur unless the molecules possess sufficient energy to get over the activation energy barrier.

SLIDE 8

Slide 43 / 119 Reaction Coordinate Diagrams

It is helpful to visualize energy changes throughout a process on a reaction coordinate diagram like this one for the rearrangement of methyl isonitrile.

H3CNC H3CCN Methyl isonitrile to Methyl nitrile

Slide 44 / 119 Reaction Coordinate Diagrams

In order to draw one of these diagrams we must consider the starting and ending formations of the molecule as well as all transition states needed to go from one to the other. Molecules don't simply jump from point A to point E, but must visit all transition states along the way.

H3CNC H3CCN Methyl isonitrile to Methyl nitrile

Slide 45 / 119 Reaction Coordinate Diagrams

We start with what we know is the shape of the starting molecule.

Slide 46 / 119 Reaction Coordinate Diagrams

Since we know that there must be an increase in energy, the molecule must actually shift to a less favorable formation. This conversion requires energy which is why we see the gradual increase on the energy level diagram.

Slide 47 / 119 Reaction Coordinate Diagrams

Notice how odd the highest energy formation looks. The bonds are so strained at this point that the maximum amount of energy is needed in

• rder to acheive this formation..

Slide 48 / 119 Reaction Coordinate Diagrams

As the energy comes down off the peak energy, the formation starts to resemble a formation we view as more "normal", meaning it requires less energy to maintain.

SLIDE 9

Slide 49 / 119 Reaction Coordinate Diagrams

Finally, the molecule reaches a state that requires less energy than was

• riginally required to establish the original molecule. Based on the

difference between the starting and ending energies we can determine that this process releases energy and is therefore exothermic.

Slide 50 / 119

The diagram shows the energy

• f the reactants and products

(and, therefore, E). The high point on the diagram is the transition state. The species present at the transition state is called the activated complex.The energy gap between the reactants and the activated complex is the activation energy barrier.

Reaction Coordinate Diagrams Slide 51 / 119 Maxwell–Boltzmann Distributions

Temperature is a measure of the average kinetic energy of the molecules in a sample. At any temperature there is a wide distribution of kinetic energies.

Fraction of molecules Speed / KE Ea

Slide 52 / 119 Maxwell–Boltzmann Distributions

As the temperature increases, the curve flattens and broadens. Thus at higher temperatures, a larger population of molecules has higher energy.

Fraction of molecules Speed / KE Ea

According to Kinetic Molecular Theory, higher temperatures mean more collisions. More collisions mean more exchange of energy. As more energy is exchanged, we end up with a "flattening" of the graph.

Slide 53 / 119

If the dotted line represents the activation energy, then as the temperature increases, so does the fraction of molecules that can overcome the activation energy barrier. As a result, the reaction rate increases.

Maxwell–Boltzmann Distributions

F r a c t i

• n
• f

m

• l

e c u l e s Speed / KE Ea

Slide 54 / 119 Potential Energy Diagrams

A potential energy diagram shows how PE changes as a reaction proceeds, i.e. over time. In this example, the PE of the products are LOWER than the PE of the reactants. Therefore, #H° is negative and the reaction is exothermic.

SLIDE 10

Slide 55 / 119

9 This potential energy diagram represents a reaction that is ____________ with a #H of _____.

A exothermic, -10 kJ B exothermic, -25 kJ C endothermic, +15 kJ D endothermic, +25 kJ

answer

Slide 56 / 119

10 This potential energy diagram represents a reaction that is ____________ with a #H of _____.

A exothermic, -5 kJ B exothermic, -20 kJ C endothermic, +5 kJ D endothermic, +25 kJ

answer

Slide 57 / 119

11 This reaction is _________ with a #H of _________.

A exothermic, -25 kJ B C endothermic, +25 kJ D endothermic, +40 kJ exothermic, -15 kJ

answer

Slide 58 / 119 Activation Energy

In order for a reaction to

• ccur, a minimum amount of

energy must first be applied to the reactants. This minimum energy input is needed to break bonds, usually. This minimum energy input is called the Activation Energy, or Ea.

Slide 59 / 119 Activated Complex

At the peak of the graph, the substance that exists is called the activated complex, or transition state. This activated complex is referred to as an intermediate, since it exists as an intermediate step between the existence of reactants and products. The activated complex is very unstable, due to its high PE.

Slide 60 / 119

12 The activation energy for the forward process in this reaction is:

A +15 kJ B

• 15 kJ

C +25 kJ D

• 25 kJ

answer

SLIDE 11

Slide 61 / 119

13 The activation energy for the reverse reaction is __________.

A +25 kJ B

• 25 kJ

C +40 kJ D

• 40 kJ

answer

Slide 62 / 119

14 What is another name for the activated complex?

A energy barrier B transition state C rate limiter D collision group E reactant or product

answer

Slide 63 / 119

15 At what stage of a reaction do atoms have the highest energy?

A reactant stage B product stage C transition state stage

answer

Slide 64 / 119

16 Activation energy is __________.

A the heat released in a reaction B the energy given off when reactants collide C generally very high for a reaction that takes place rapidly D an energy barrier between reactants and products

Slide 65 / 119 Catalysts

Catalysts increase the rate of a reaction by decreasing the activation energy of the reaction.

This graph shows the decomposition of glucose both with and without a catalyst. Notice that the energies of reactants and products are unchanged by the catalyst.

Catalysts are very specific and will work only in a given direction. Always read the equation left to right unless specified.

Slide 66 / 119

One way a catalyst can speed up a reaction is by holding the reactants together and helping bonds to break.

Catalysts

Shown here is hydrogen reacting with ethylene. Both molecules are adsorbed onto the metal (blue) surface which acts as the catalyst.

SLIDE 12

Slide 67 / 119

Automobiles are equipped with catalytic converters, which are part of their exhaust systems. The exhaust gases contain CO, NO, NO2, and unburned hydrocarbons that pass over surfaces impregnated with catalysts.

Catalytic converters

The catalysts promote the conversion of the exhaust gases into CO2, H2O, and N2.

Slide 68 / 119 Enzymes

Enzymes are catalysts in biological systems. The substrate fits into the active site of the enzyme much like a key fits into a lock.

Substrate entering active site of enzyme Enzyme/substrate complex Enzyme/products complex Products leaving active site of enzyme

Active site

substrate

Slide 69 / 119 Enzymes

Liver contains an enzyme called catalase which speeds up the decomposition of H2O2.

Slide 70 / 119

17 Why does a catalyst cause a reaction to proceed faster?

A There are more collisions per second. B Collisions occur with greater energy. C The activation energy is lowered. D There are more frequent collisions and they are of greater energy.

answer

Slide 71 / 119

18 Which of the following acts as catalysts in the body?

A carbohydrates B nucleic acids C lipids D enzymes E water

Slide 72 / 119

We have previously seen that when there is no change in Gibbs free energy, #G = 0, a system is at equilibrium. An equilibrium state can be either physical or chemical.

Equilibrium

SLIDE 13

Slide 73 / 119

Physical equilibrium occurs when there are changes of physical state (phase) occurring at the same rate. For example, a mixture of ice and water at 0°C is undergoing melting and freezing simultaneousely.

Physical Equilibrium Slide 74 / 119

Another example of physical equilibrium takes place between dissolved solute and undissolved solid in a saturated solution.

Physical Equilibrium

Two processes, the dissolving and precipitation of solute,

• ccur simultaneously and at the same rate.

Slide 75 / 119 Chemical Equilibrium

Chemical equilibrium occurs when a reaction and its reverse reaction proceed at the same rate. N2O4 (g)

2 NO2 (g)

When both the forward and reverse reactions are being carried

• ut, we write the equation with a double arrow.

Slide 76 / 119 Chemical Equilibrium

N2O4 is a colorless gas. It decomposes into NO2, a brown-colored gas. Eventually, the forward and backward reactions reach a state of

• equilibrium. This means that the colorless gas is decomposing at

the same rate at which the brown gas is combining. N2O4 (g)

2 NO2 (g)

Slide 77 / 119

When a reaction begins, the forward reaction

• ccurs quickly, at first then

begins to slow down. The reverse reaction is non-existent at the start, but then picks up speed as more product is created.

Chemical Equilibrium

Kf [N2O4] Kr [NO2]2 Equilibrium achieved (rates are equal)

R a t e Time

N2O4 (g)

2 NO2 (g)

Slide 78 / 119 The Concept of Equilibrium

As a system approaches equilibrium, both the forward and reverse reactions are occurring. At equilibrium, the forward and reverse reactions are proceeding at the same rate. Once equilibrium is achieved, the amount of each reactant and product remains constant.

Kf [N2O4] Kr [NO2]2 Equilibrium achieved (rates are equal)

Rate Time

* This does not imply that the amount of reactants = amount of products.

SLIDE 14

Slide 79 / 119 Chemical Equilibrium

It is important to note that equilibrium can be achieved regardless of whether you start with reactants or products, as long as there is sufficient material to get both processes going. N2 (g) + 3H2 (g) <--> 2NH3 (g)

Slide 80 / 119 Chemical Equilibrium

For the reaction N2 (g) + 3H2 (g) <--> 2NH3 (g), you can start with nitrogen and hydrogen and no ammonia, or you can start with

• nly ammonia. Either way, the same equilibrium concentrations
• f all three substances will eventually be reached.

Slide 81 / 119

19 The following reaction reaches equilibrium around_________ A 12 seconds B 22 seconds C 32 seconds D 42 seconds E 52 seconds

H2(g) + Cl2(g) <--> 2HCl(g) Pressure (atm) time (seconds)

0 10 20 30 40 50

H2(g) = Cl2(g) = HCl(g) =

answer

Slide 82 / 119

20 Equilibrium is reached when...

A The rates of the forward and reverse reaction become equal B The concentrations of products and reactants become equal C When G for the reaction is equal to zero D A and C E A, B, and C

Slide 83 / 119

21 Which of the following would, if TRUE, signal that the reaction or process has reached equilibrium?

A The concentrations of reactants and products have become equal B The concentrations of reactants and products stop changing C The reactants stop turning into products D The products stop turning into reactants E None of these indicate a reaction has reached equilibrium

answer

Slide 84 / 119

Therefore, we can write a rate law for the forward reaction: This rate law is And, we can write a rate law for the reverse reaction: This rate law is N2O4 (g)

2 NO2 (g)

Chemical Equilibrium

N2 O4 (g) ---> 2 NO2 (g) Rate = kf [N2 O4 ] 2 NO2 (g) ---> N2 O4 (g) Rate = kr [NO2 ]2 Recall that the rate of a reaction is proportional to the reactant concentration.

SLIDE 15

Slide 85 / 119 The Equilibrium Constant

Recall that the definition of chemical equilibrium is when the rate of the forward reaction equals the rate of the reverse reaction. Therefore, at equilibrium, we can set the two rate laws equal: Rate of forward reaction = Rate of reverse reaction kf [N2O4] = kr [NO2]2

Rewriting this, it becomes

kf kr [NO2]2 = [N2O4]

Slide 86 / 119 Slide 87 / 119 Slide 88 / 119 Slide 89 / 119

The equilibrium constant for the reaction is CH4 (g) + 2O2 (g) CO2 (g) + 2H2O(l) [CO2 ] [CH4 ][O2]2 K = slide for answer

The Equilibrium Constant Slide 90 / 119

The equilibrium constant for the reaction 2Cl2 (g) + 2H2O(g) 4HCl(g) + O2 (g) K= [HCl]4 [O2] [Cl2 ]2 [H2O]2

The Equilibrium Constant

slide for answer

SLIDE 16

Slide 91 / 119

The equilibrium constant for the reaction (CH3)4Sn(s) (CH3)4Sn (g) K = [(CH3)4Sn]

The Equilibrium Constant

slide for answer

Slide 92 / 119

22 What would be the correct equilibrium expression for the following reaction? A B C D

4Fe(s) + 3O2(g) --> 2Fe2O3(s) [Fe2O3]2 [O2]3[Fe]4 [Fe]4[O2]3 [Fe2O3]2

[O2]3

1

1

[O2]3

Slide 93 / 119 The Equilibrium Constant

Since pressure is proportional to concentration for gases in a closed system, the equilibrium expression can also be written in terms of pressure

P

A = partial pressure of

reactant A and so on Consider the generalized reaction aA + bB cC + dD

⇌

** Slide 94 / 119 Relationship Between Kc and Kp

From the Ideal Gas Law we know that

** PV = nRT

Rearranging it, we get

P = nRT V [M] = n V P = [M] x RT Slide 95 / 119 Relationship Between Kc and Kp

Plugging this into the expression for Kp for each substance, the relationship between Kc and Kp becomes

# #

where

Kp = Kc (RT)#n

#n = (moles of gaseous product) - (moles of gaseous reactant)

Slide 96 / 119

If K>>1, the reaction is product- favored; product predominates at equilibrium. If K<<1, the reaction is reactant-favored; reactant predominates at equilibrium.

Reactants

Products

Reactants

Products

K ≪ 1 #--# K≫1 #--#

What Does the Value of K Mean?

SLIDE 17

Slide 97 / 119

The direction from which equilibrium is achieved (starting with all products or all reactants) doesn't matter. It is only a convention that we calculate K by dividing the concentrations of products over reactants, since the reaction proceeds both ways. Unless specified, read the equation left to right (Left -reactants, right -products) This convention makes it easy to understand the meaning of Keq .

Equilibrium Slide 98 / 119

23 For a reaction at equilibrium, a value of K= 1 x 108 means that

A reactants are favored B products are favored C the reaction is spontaneous D the reaction is endothermic

answer

Slide 99 / 119

24 For a reaction at equilibrium, a value of K= 4 x 10-12 means that

A reactants and products are present in equal amounts B the forward reaction is occurring faster than the reverse reaction C the reverse reaction is occurring faster than the forward reaction D reactants are favored E products are favored

Slide 100 / 119 Equilibrium

Application The process of rust formation has a very high equilibrium constant yet if one lays out a piece of iron, one will not see the conversion to rust. Why? The size of the equilibrium constant has no bearing on the rate of

• reaction. Rust formation, like many processes, is favorable

thermodynamically, but for a number of reasons (solid state, high activation energy, etc.) does not occur quickly.

move for answer Slide 101 / 119 Le Châtelier’s Principle

“If a system at equilibrium is disturbed by a change in temperature, change in pressure,

• r change in concentration of one of the

components, the system will shift its equilibrium position so as to counteract the effect of the disturbance.”

Henri Louis Le Châtelier

Slide 102 / 119 Le Châtelier’s Principle

+

The Haber Process: Synthesis of ammonia from hydrogen and nitrogen

Fritz Haber, 1918

SLIDE 18

Slide 103 / 119 Le Châtelier’s Principle

Initial equilibrium

H2 NH3

N2 H2 added at this time

Equilibrium reestablished

Time P a r t i a l p r e s s u r e

+

If we change the concentration:

Slide 104 / 119

If H2 is added to the system, the added hydrogen needs to be consumed.

Le Châtelier’s Principle

"take it and go" policy

Initial equilibrium H2 NH3

N2

H2 added at this time Equilibrium reestablished

Time P a r t i a l p r e s s u r e

How? By reacting with nitrogen. More N2 will be consumed with the added H2 More NH3 will be produced

Slide 105 / 119

What happens when you remove NH3 ? The system has to replace it to bring back the equilibrium More NH3 will be produced. We need large amount of NH3 for making fertilizers.

Le Châtelier’s Principle Slide 106 / 119

25 What happens to a reaction at equilibrium when a reactant concentration is lowered?

A The reaction makes more products. B The reaction makes more reactants. C The reaction makes more of both products and reactants, so equilibrium is unaffected.

answer

Slide 107 / 119

26 What happens to a reaction at equilibrium when certain amount of a product is removed from the reaction system? A The reaction makes more products.

B The reaction makes more reactants.

C The reaction makes more of both products and reactants, so equilibrium is unaffected.

answer

Slide 108 / 119

27 What is the effect of adding more water to the following equilibrium reaction? CO2 (g) + H2O(l) <--> H2CO3 (aq)

A More H2CO3 (carbonic acid) is produced. B The CO2 concentration increases. C Equilibrium shifts left, toward the reactants. D Equilibrium is unaffected since water is a pure liquid.

SLIDE 19

Slide 109 / 119

28 What is the effect of adding more CO2 (g) to the following equilibrium reaction? CO2 (g) + H2O (l) <--> H2CO3 (aq)

A More H2CO3 (carbonic acid) is produced.

B The CO2 concentration increases. C Equilibrium shifts left, toward the reactants. D Equilibrium is unaffected since water is a pure liquid.

answer

Slide 110 / 119 Le Châtelier’s Principle

If we change volume/pressure: N2 + 3H2 <--> 2NH3 If the volume is decreased, it will increase the pressure of the system, so the system will try to reduce the pressure by producing less molecules. The reaction will favor producing fewer number of molecules. The equilibrium will shift to the right. This only applies to situations where the volume of the products or reactants are not fixed. In

• ther words, this matters with gases only.

Slide 111 / 119

If the volume is increased, it will reduce the pressure of the system, So the system will try to increase the pressure by producing more molecules. The reaction will favor producing more molecules. The equilibrium will shift to the side with more moles - in this case- the left.

Le Châtelier’s Principle

If we change volume/pressure: N2 + 3H2 <--> 2NH3

Slide 112 / 119 Le Châtelier’s Principle

Stress: Increasing pressure/ reducing volume Effect: Equilibrium shifts to the right Stress: Decreasing pressure/ increasing volume Effect: Equilibrium shifts to the left

N2 (g) + 3H2 (g) <--> 2NH3 (g) The Haber Process - Summary of pressure changes

Slide 113 / 119

29 How does an increase in pressure affect the following reaction at equilibrium? C2H2 (g) + H2 (g) <--> C2H4 (g)

A Equilibrium shifts to the right; the reaction makes more products. B Equilibrium shifts to the left; the reaction makes more reactants. C The reaction makes more of both products and reactants, so equilibrium is unaffected.

answer

Slide 114 / 119

30 Which of the changes listed below would shift the equilibrium of this reaction to the right?

4HCl(g) + O2 (g) <--> 2Cl2 (g) + 2H2O(g)

A

Addition of Cl 2 gas.

B

Removal of O 2 gas.

C

Increase in pressure.

answer

SLIDE 20

Slide 115 / 119

If we change the Temperature:

Le Châtelier’s Principle

Co(H2O)62+(aq) + 4 Cl(aq) CoCl42- (aq) + 6 H2O (l)

At high temp, the blue hexa aqua species is favored At low temp, the purple tetra chloro species is favored

at 100 0C at 25 0C

Many ions containing transition metals produce colored

• solutions. In the reaction above, the cation, Co(H2O)6 2+ , yields

a pink solution, while the anion CoCl4 2- produces a blue solution.

Pink blue

H = (+)

Click here for animation

Slide 116 / 119

If the temperature of the system is increased: PCl5 <--> PCl3 + Cl2 ΔH = 88 KJ endothermic Since the reaction is endothermic...view energy as a reactant Energy + PCl5 <--> PCl3 + Cl2 The system should accept (take in) the thermal energy supplied to favor the endothermic reaction, and the equilibrium will shift to the right (forward reaction). If the temperature of the system is lowered: The system will restore equilibrium by producing energy hence resulting in a shift of the reaction back to the left. If we change the Temperature:

Le Châtelier’s Principle

Note: The only change that will affect the magnitude of the equilibrium constant is a change in the temperature!!

Slide 117 / 119

31 For the reaction below, what is the effect of raising the temperature?

4A(g) + B(g) <--> 3C(g) #H = - 405 kJ A Equilibrium shifts to the right; the reaction makes more products. B Equilibrium shifts to the left; the reaction makes more reactants. C The reaction makes more of both products and reactants, so equilibrium is unaffected.

answer

Slide 118 / 119 Equilibrium

Application The reaction that occurs in a car engine is an exothermic

• process. Explain why keeping your engine cool might ensure that

all of the gasoline has been burned. gasoline + O2 --> CO2 + H2O + heat If heat is removed, the reaction will shift right ensuring full conversion to product

move for answer Slide 119 / 119