Unit 9: Kinetics and Equilibrium

Chapter 12

General Chemistry I

Edmond North High School

Collision Theory

• The following three statements summarize the collision theory. – 1. Particles must collide in order to react. – 2. The particles must collide with the correct orientation.– 3. The particles must collide with enough energy to form an activated

complex, which is an intermediate particle made up of the joined reactants.

• An effective collision is one that results in a reaction; new products are formed.

The Collision Model

• The Collision Model is used to explain current known characteristics of reaction rates.

• According to the Collision Model:– Molecules must collide to react.– Concentration affects rates because collisions are more likely.– Must collide with enough energy.– Temperature and rate are related.– Only a small number of collisions produce reactions.

Activated Complex

• The Activated Complex is an intermediate particle formed when reactants collide and stick together.– Old bonds are breaking while

new bonds are forming.• The minimum amount of energy

colliding particles must have in order to form an activated complex is called the activation energy. – Particles that collide with less

than the activation energy cannot form an activated complex.

Reaction Coordinate Diagrams

• It is helpful to visualize energy changes throughout a process on a reaction coordinate diagram.– It shows the energy of the

reactants and products (and, therefore, H).

– The high point on the diagram is the transition state.

– The species present at the transition state is called the activated complex.

Exothermic Reactions Review

• An exothermic reaction releases heat, and an endothermic reaction absorbs heat.

Endothermic Reactions Review

• The endothermic reaction absorbs heat because the products are at a higher energy level than the reactants.

Orientations

Kinetics

• Kinetics is the study of reaction rates and the factors that effect them.

• Why is it important?– When we understand

reaction rates we can control chemical reactions and use them for specific purposes.

Kinetics

• As a reaction occurs:– concentration of reactants decreases– concentration of products increases

• Kinetics studies the rate at which a chemical process occurs.• Besides information about the speed at which reactions

occur, kinetics also sheds light on the reaction mechanism (exactly how the reaction occurs).

Reaction Rate

• Reaction Rate is the number of atoms, ions or molecules that react in a given time to form products– We measure either the rate of disappearance of the

reactants or the rate of appearance of one of the products

– In simpler terms:• The average rate is the change in a given quantity

during a specific period of time.

Measuring Reaction Rates

• Change in Electrical Conductivity• Change in Color• Change in Pressure• Change in Volume

Reaction Rate Graph

• For the reaction H2 + I2 → 2HI– What is happening to the concentration of the

reactants?– What is happening to the concentration of the

products?

Reaction Rates & Stoichiometry

• We use coefficients to give us a ratio of rates between species in a reaction.– In the reaction 3H2 + N2 2NH3, since N2 has a

coefficient of 1 and NH3 has a coefficient of 2, the rate N2 is used = 2(rate NH3 is created)

– Likewise, since H2 has a coefficient of 3 and N2 has a coefficient of one, it is used 3 times as fast as N2.

Factors Affecting Reaction Rates

• The reaction rate for almost any chemical reaction can be modified by varying the conditions of the reaction. – An important factor that affects the rate of a chemical

reaction is the reactive nature of the reactants. As you know, some substances react more readily than others. • The more reactive a substance is, the faster the

reaction rate.

Factors That Affect Reaction Rates

• Physical State of the Reactants– In order to react, molecules must come in contact with each

other.– The more homogeneous the mixture of reactants, the faster the

molecules can react.

• Concentration of Reactants– As the concentration of reactants increases, so does the

likelihood that reactant molecules will collide.

Factors that Effect Reaction Rate

• Particle Size (surface area)– For solids, breaking up

big pieces increases surface area, increasing rate by having more places for the molecules to interact

– Ionic compounds have more surface area than covalent

Factors that Effect Reaction Rate

• Pressure – Gases only!– As pressure increases

the concentration increases, so you will have more collisions

Factors That Affect Reaction Rates

• Temperature– At higher temperatures, reactant molecules have

more kinetic energy, move faster, and collide more often and with greater energy.

• Presence of a Catalyst– Catalysts speed up reactions by changing the

mechanism of the reaction.• Catalysts are not consumed during the course of the

reaction.

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Catalysts

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

mechanism by which the process occurs.

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

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.

Factors that Effect Reaction Rate

• Inhibitors–Compounds

added to a reaction that slow it down• Examples: Lead in

diesel, preservatives in food.

Reaction Rate Laws

• Rate Laws relate the reaction rate and reactant concentration– We write the rate laws for elementary

reactions from the equation.– They include only the gaseous or aqueous

reactants!• Solids & liquids are not included because they

have a constant concentration that does not increase or decrease

Rate Law

• A Rate Law is an expression that shows how the rate depends on concentration of reactants

• Example:– Rate laws use reactants only– k is proportionality constant (called the rate constant)– n is the order (exponent) of the reactant determined

experimentally• n is generally the coefficient of the chemical in the balanced

equation HOWEVER experimental data overrides the coefficient.– Reaction is xth order in A– Reaction is yth order in B– Reaction is (x +y)th order overall

Rate = k [A]x[B]yaA + bB cC + dD

Order of Reaction

• Order of Reaction– For the reaction with the rate = k[A]2[B]

• Rate is 2nd order with respect to A• Rate is 1st order with respect to B• Rate is 3rd order overall (add exponents)

• How does concentration affect reaction rate?– If [A] doubles what happens to the rate?– If [B] triples, what happens to the rate?

Temperature and Rate

• Generally, as temperature increases, so does the reaction rate.

• This is because k is temperature dependent.

Indicators of Completed Reactions

• So how do you know when a reaction has gone to completion?– A gas is produced and

escapes (open container)

– A precipitate forms from two aqueous solutions.

– A covalent product is formed (usually water)

What is Equilibrium?

• When a reaction results in almost complete conversion of reactants to products the reaction goes to completion.

• Most reactions, however, do not go to completion. They appear to stop. – The reason is that these

reactions are reversible.

• A reversible reaction is one that can occur in both the forward and the reverse directions.

Chemical Equilibrium

• Chemical equilibrium occurs when two opposite reactions occurring at the same time and rate

A + B ↔ AB• Forward & Reverse reactions don’t stop at

equilibrium, just looks that way because concentration remains constant– Rateforward reaction = Ratereverse reaction

Equilibrium Example

• You have a bridge between 2 cities. The number of cars going in either direction on the bridge are equal (at equilibrium)

• The populations of the cities on either side of the bridge do not have to be equal!

Equilibrium Position

• The equilibrium position of a reaction is determined by:– Initial concentrations– Energy of reactants and products– Degree of organization of reactants and products

Progression of Equilibrium

• As a reaction progresses:– [A] decreases to a constant,– [B] increases from zero to a constant.– When [A] and [B] are constant, equilibrium is

achieved.

• In a system at equilibrium, both the forward and reverse reactions are being carried out; as a result, we write its equation with a double arrow.

A B

A System at Equilibrium

• Once equilibrium is achieved, the amount of each reactant and product remains constant.– Rates are equal; concentrations are not.– The concentrations do not change at

equilibrium.

Equilibrium Positions

• If equilibrium lies “to the left”:– There are more reactants and less products.

• If equilibrium lies “to the right”:– There are less reactants and more products.

• If reactants are mixed and concentrations do not change:– The reaction could already be at equilibrium.– Reaction rates are so slow that change is too difficult

to detect.

Equilibrium Expressions

• Equilibrium expressions relate concentrations of reactants to those of the products– These can be written from balanced equations– They look similar to rate laws! BE CAREFUL!

• Uses both reactants & products• Include only the gaseous or aqueous phases

as with rate laws

Deriving an Equilibrium Expression

• Forward reaction: N2O4 (g) 2NO2 (g)– Rate law: Rate = kf [N2O4]

• Reverse reaction: 2NO2 (g) N2O4 (g)– Rate law: Rate = kr [NO2]2

• Therefore, at equilibrium: Ratef = Rater

– kf [N2O4] = kr [NO2]2

• Rewriting this, it becomes

Keq =kfkr

[NO2]2

[N2O4]=

General Equilibrium Expressions

• A (s) + 2B (g) ↔ 2C (g) + 3D (g)

Keq = [C]2[D]3

[B]2

• Keq = equilibrium constant (capital K)– Numerical value of the ratio of product

concentrations to reactant concentrations

• [ ] = concentration in M (mol/L)• Order = coefficient becomes exponent

Solids and Liquids Are Constant

• Both can be obtained by dividing the density of the substance by its molar mass—and both of these are constants at constant temperature.– Therefore, the concentrations of solids and

liquids do not appear in the equilibrium expression.

Kc = [Pb2+] [Cl−]2

PbCl2 (s) Pb2+ (aq) + 2 Cl−(aq)

What Are the Equilibrium Expressions for These Equilibria?

Equilibrium Constant, K

• If we know the value of K, we can predict:– The tendency of a reaction to occur.– If a set of concentrations could be at equilibrium.– The equilibrium position, given initial concentrations.

• K will always have the same value at a certain temperature.– No matter what amounts are added, the ratio at

equilibrium will always be same

Equilibrium Constant (Keq)

• Tells you whether the products or reactants are favored– Keq > 1

• Products are favored (forward rxn); lots of product is made

– Keq < 1• Reactants are favored (reverse

rxn); not much product made

• The size of K and time needed to reach equilibrium are NOT related

The Units for K

• Are determined by the various powers and units of concentrations.

• They depend on the reaction.• At any temperature.

– Temperature affects rate.– Equilibrium position is a set of concentrations

at equilibrium.

The Reaction Quotient (Q)

• To calculate the reaction quotient or Q, one substitutes the initial concentrations on reactants and products into the equilibrium expression.– Q gives the same ratio the equilibrium

expression gives, but for a system that is not at equilibrium.• It is used to tell if a reaction is at equilibrium or

not.

Reaction Quotient: Comparing Q and K

• The relationship between Q and K tells which way the reaction will shift– Q = K: at equilibrium,

no shift– Q > K: too large, forms

reactants, shift to left– Q < K: too small, forms

products, shift to right

Le Chatelier’s Principle

• Le Chatelier’s Principle states that if stress is added to a system at equilibrium, the reaction will speed up in the direction that will relieve the stress.– Once the stress is relieved,

equilibrium is re-established and no further changes are noticed.

• 4 Types of stress– Concentration– Temperature– Pressure– Volume

Stress Factors

1. Concentration– Shifts away from an increase or addition– Shifts toward a decrease or subtraction

Stress Factors

2. Temperature (treat heat as a reactant if endothermic or product if exothermic)– Shifts away from an addition– Shifts toward a subtraction

Change in Temperature

• Temperature affects the rates of both the forward and reverse reactions.– Doesn’t just change the equilibrium position,

changes the equilibrium constant.

Stress Factors

3. Pressure (only effects gases)– Increase pressure,

shifts to side with lower total moles of gas

– Decrease pressure, shift to side with higher total moles of gas

CO (g) + H2 (g) ↔ CH4 (g) + H20 (g)

Stress Factors

4. Changes in the volume • Suppose the volume of the reaction vessel for the

system is decreased, resulting in an increase in pressure.– The equilibrium will shift to relieve the stress of

increased pressure.

Change in Volume

• Decrease V:– Decreases # gas molecules– Shifts towards the side of the reaction with less gas

molecules

• Increase V:– Increase in # of gas molecules– Shifts towards the side of the reaction with more gas

molecules

Catalysts

• Adding a Catalyst– Does not change K– Does not shift the position of an equilibrium system– System will simply reach equilibrium sooner

Stress Factor Practice

N2 (g) + 3H2 (g) ↔ 2NH3 (g) + heat

• Which way will the reaction shift if you:– Add N2

– Decrease pressure– Remove NH3

– Heat– Add a catalyst

The End.

Be Prepared for Unit 9 Test!