Module 4 Rates of Reactions This module roughly corresponds to Chapters 8, 9 and 10 of your textbook (but we will change the order of some items)

Module 4Module 4Rates of ReactionsRates of Reactions

This module roughly corresponds to This module roughly corresponds to Chapters 8, 9 and 10 of your textbookChapters 8, 9 and 10 of your textbook(but we will change the order of some (but we will change the order of some

items)items)

Module 4: Lesson 1Module 4: Lesson 1

Slow and Fast ReactionsSlow and Fast Reactions

Reaction Speeds (relative)

• Chemical reactions can take place at different relative rates• Slow• Moderate• Fast• Explosive

Examples of Relatively Slow Examples of Relatively Slow ReactionsReactions

Metal Rusting (oxidation, corrosion)Metal Rusting (oxidation, corrosion) Colours fading (oxidation)Colours fading (oxidation)

Moderate Speed ReactionsModerate Speed Reactions(relatively speaking)(relatively speaking)

Electrolysis of waterElectrolysis of water Smoldering (slow burning)Smoldering (slow burning) Neutralization reactionsNeutralization reactions

Fast Reactions Fast Reactions (Relatively) (Relatively)

BurningBurning

PrecipitationPrecipitation

Explosive Reactions(relatively very fast)

• Dynamite• Gun powder• Hydrogen burning• Fireworks

Problem with Relative RatesProblem with Relative Rates

These categories are relative. These categories are relative. They are based on the judgement of They are based on the judgement of

the observer. the observer. They do not have measurements, so it They do not have measurements, so it

is difficult to compare them is difficult to compare them accurately.accurately.

We need a QUANTITATIVE We need a QUANTITATIVE measurement of reaction rate.measurement of reaction rate.

Formula for Reaction RateFormula for Reaction Rate A Rate in general is the change in some A Rate in general is the change in some

quantity in a certain time quantity in a certain time

Reaction Rate =Reaction Rate =Units of amountUnits of amount = = ΔΔamountamount

Units of TimeUnits of Time ΔΔtimetime

t

Ar

How is How is Amount (Amount (ΔΔA)A) measured?measured?

Units of amount may be: Units of amount may be: molesmoles, the preferred unit if it is known. , the preferred unit if it is known. grams, sometimes used for convenience.grams, sometimes used for convenience. concentration, concentration, Volume. Used in experiments which invole Volume. Used in experiments which invole

gases. Volume may be proportional to gases. Volume may be proportional to molesmoles

partial pressure, in some gas experiments.partial pressure, in some gas experiments. It depends on which property can be It depends on which property can be

measured by your experiment.measured by your experiment.

How is How is TimeTime ((ΔΔt) t) measured?measured?

Units of time may be: Units of time may be: SecondsSeconds, the preferred unit of time, the preferred unit of time Minutes, for moderate speed reactions,Minutes, for moderate speed reactions,

Usually converted to seconds before calculations.Usually converted to seconds before calculations. Hours, for slower reactionsHours, for slower reactions Days, for quite slow reactionsDays, for quite slow reactions Years, for extremely slow reactionsYears, for extremely slow reactions

Seconds are the best to useSeconds are the best to use unless they are inconvenient in your unless they are inconvenient in your

experimentexperiment

The Rate of a Reaction The Rate of a Reaction ChangesChanges

Most reactions start off fast, and then Most reactions start off fast, and then slow down.slow down. A graph can show the changing rate of a A graph can show the changing rate of a

reaction.reaction.

A few reactions may start slowly, and A few reactions may start slowly, and then speed up.then speed up.

Graphing Reaction RatesGraphing Reaction Rates

Units of Time

Uni

ts o

f A

mou

nt (

chan

ge)

Product

Reactant

ΔTime

ΔAmount

ΔTime

ΔAmount

1

2

At time 1 the reaction is faster than at time 2. As the product is formed, the reactant is used up

The rate of change of the products is positive (+)

The rate of change of the reactants is Negative (-)

The preceding graph assumes that the The preceding graph assumes that the reactant is being used up at the same reactant is being used up at the same rate as the product is being produced. rate as the product is being produced. This is not always true. For example, in This is not always true. For example, in the reaction Hthe reaction H22 + I + I22 2 HI two moles of 2 HI two moles of hydrogen iodide are produced for every hydrogen iodide are produced for every mole of Iodine that is used up, so the mole of Iodine that is used up, so the product line would be steeper!product line would be steeper!

Rates for HRates for H22 + I + I22 2HI 2HI

Units of Time

Mol

es o

f S

ubst

ance

Amount of HI

Amount of I2

The Stoichiometry EffectThe Stoichiometry Effect

The difference in rates is related to The difference in rates is related to the stoichiometry of the reaction.the stoichiometry of the reaction.

The Stoichiometry EffectThe Stoichiometry Effect

t

D

dt

C

ct

B

bt

A

ar

1111

For an equation like:

aA + bB cC + dDWhere a,b,c and d are coefficients, and ABC and D are chemical formulasThe following relationship is true...

Module 4, Lesson 2Module 4, Lesson 2

Factors which affect the rate of Factors which affect the rate of reactionsreactions

There are many factors that can There are many factors that can increase or decrease the rate of a increase or decrease the rate of a chemical reaction, including:chemical reaction, including: The The naturenature of the substances reacting of the substances reacting The The surfacesurface area in contact area in contact The presence of a The presence of a catalystcatalyst The The temperaturetemperature The The concentrationconcentration of the substances of the substances

(aqueous solutions only)(aqueous solutions only) The The pressurepressure of the reactants of the reactants

(gaseous reactions only)(gaseous reactions only)

Nature of the SubstancesNature of the Substances(phase, number of bonds and type of bonds)(phase, number of bonds and type of bonds)

As a general rule:As a general rule: Homogeneous reactions Homogeneous reactions (all reactants in the same (all reactants in the same

phase) are usually faster than heterogeneous phase) are usually faster than heterogeneous reactions (reactants in different phases).reactions (reactants in different phases).

Phases: solid (s), liquid (l), gaseous (g), aqueous (aq)Phases: solid (s), liquid (l), gaseous (g), aqueous (aq) The The fewer bonds fewer bonds that need to be broken, the faster that need to be broken, the faster

the reaction is.the reaction is. Double bonds are much harder to break than single bondsDouble bonds are much harder to break than single bonds

GaseousGaseous substances usually react faster than substances usually react faster than liquids, liquids usually react faster then solidsliquids, liquids usually react faster then solids

Reactions involving Reactions involving ionicionic compounds are usually compounds are usually faster than reactions involving covalent compoundsfaster than reactions involving covalent compounds

ReactiveReactive elements (alkali metals, halogens) are elements (alkali metals, halogens) are faster to react than less reactive elementsfaster to react than less reactive elements

FasterFaster > Slower> Slower Homogeneous Homogeneous > Heterogeneous> Heterogeneous Fewer bonds Fewer bonds to break to break > More bonds > More bonds to to

breakbreak

Gases Gases > Liquids > Liquids > Solids> Solids IonicIonic > Covalent> Covalent Active ElementsActive Elements > Inactive > Inactive

ElementsElements

Some Kinetic Theory:Why are ionic compounds faster to

react?

• Ionic compounds are held together more loosely than covalent ones. It takes less energy to break them apart.

Effect of Surface AreaEffect of Surface Area

Grinding the reactants into smaller Grinding the reactants into smaller pieces increases their effective pieces increases their effective surface areas.surface areas.

Reactions involving powdered Reactions involving powdered reactants will be faster than reactants will be faster than reactions with solid chunks.reactions with solid chunks.One large cube

Surface area = 24cm2

Eight small cubesSurface area = 48cm2

Kinetic Theory:Why does surface area affect the rate of reactions?

• The more surface area a substance has, the more places there are for it to get hit by molecules.

• This means there are more collisions which means the reaction will be faster.

CatalystsCatalysts

A A catalystcatalyst is a substance which can is a substance which can speed up a reaction without being speed up a reaction without being consumed (used up) itself.consumed (used up) itself.

An An inhibitoinhibitor is a substance that can r is a substance that can slow or stop a reaction.slow or stop a reaction.

Note: Note: Heat is Heat is NOT NOT a catalyst. a catalyst. Catalysts must be made of matter, Catalysts must be made of matter, not energy.not energy.

Kinetic Theory:Catalyst

• Catalysts lower the “activation energy” of a reaction

• The collisions between molecules become more “effective”

Effect of TemperatureEffect of Temperature

In general:In general: Increasing the temperature of the reactants Increasing the temperature of the reactants

usually increases the rate of reactionusually increases the rate of reaction Often increasing by 10Often increasing by 10°°C will nearly double C will nearly double

the rate.the rate. But:But:

There are a few situations where increasing There are a few situations where increasing the temperature may slow the reaction the temperature may slow the reaction (perhaps by interfering with a catalyst)(perhaps by interfering with a catalyst)

Kinetic Theory: Why does temperature affect reaction rate?• According the kinetic theory, at

higher temperatures particles (such as atoms & molecules) move faster, and therefore collide more often and with greater energy.

• This means there are more chances to react.

ConcentrationConcentration In aqueous solutions, if the reactants In aqueous solutions, if the reactants

are more concentrated, they will react are more concentrated, they will react more quickly.more quickly.

Reactions slow down as the Reactions slow down as the concentration of reactants decreases.concentration of reactants decreases. This is because solute particles are closer This is because solute particles are closer

together in a concentrated solution and together in a concentrated solution and react quickly, but far apart in a dilute react quickly, but far apart in a dilute solution.solution.

This effect is less noticeable in reactions This effect is less noticeable in reactions involving solid solutions (alloys) than liquid involving solid solutions (alloys) than liquid or gaseous solutions. or gaseous solutions.

Pure solids have no concentration so there is no effect on them. Pure solids have no concentration so there is no effect on them.

Kinetic Theory:concentration

• In a concentrated solution there are more reactant particles, so they will collide with each other more often.

PressurePressure

In reactions involving gases, higher In reactions involving gases, higher pressure can increase the rate of pressure can increase the rate of reaction. reaction.

This is similar to increasing the This is similar to increasing the concentration for solutions.concentration for solutions.

Kinetic Theory:Pressure

• In a gas at high pressure, the particles are colliding with each other more often.

Summary Lesson 2

• Factors which affect rates– Nature of reactants (more reactive=faster)– Surface area (more surface area = faster)– Temperature (higher temperature = faster)– Catalyst (lower activation energy = faster)– Concentration of solutions (higher = faster)– Pressure of gases (higher pressure = faster)

• Definition of a catalyst– A substance that speeds up a chemical reaction

without being used up.

Module 4, Lesson 3Module 4, Lesson 3

Collision TheoryCollision Theory

Collision TheoryCollision Theory

This is the theory that chemical This is the theory that chemical changes proceed when reacting changes proceed when reacting molecules collide with sufficient molecules collide with sufficient energy to rearrange the atoms.energy to rearrange the atoms.

Activation energyActivation energy is the minimum is the minimum amount of energy with which amount of energy with which particles must collide in order for the particles must collide in order for the collision to be collision to be effectiveeffective,,

An An Effective CollisionEffective Collision* is one which can * is one which can result in the formation of new result in the formation of new molecules by rearranging the bonds.molecules by rearranging the bonds.

During an effective collision, an During an effective collision, an activated complexactivated complex is usually formed. is usually formed.

The activated complex is a temporary, The activated complex is a temporary, unstable arrangement of atoms unstable arrangement of atoms present during the collision (right at present during the collision (right at the point of highest potential energy.)the point of highest potential energy.)

*Your new textbook uses the term “*Your new textbook uses the term “inelasticinelastic” to ” to describe an effective collision. This is a different describe an effective collision. This is a different usage of the term inelastic than was used in the usage of the term inelastic than was used in the older textbook!older textbook!

Molecular CollisionsMolecular Collisions

2 H2 2 S

2 molecules of Hydrogen 2 atoms of Sulfur

Ineffective Collision

(molecules didn’t have enough energy)

Molecular CollisionsMolecular Collisions

2 H2 2 S

2 molecules of Hydrogen 2 atoms of Sulfur

Effective Collision

Formation of activated complex

Molecular CollisionsMolecular Collisions

H2 S

Molecule of hydrogen sulfide Molecule of hydrogen sulfide

Effective Collision

New molecules formed

H2 S

Enthalpy (Potential Energy) Enthalpy (Potential Energy) GraphGraph

shows enthalpy during reactionshows enthalpy during reaction

Enthalpy of Reactants

Enthalpy of Activated Complex

Enthalpy of Products

ΔH (Heat of reaction)

Activation Energy (EA)

Reaction Progress (time)

Pot

entia

l Ene

rgy

(ent

halp

y)

A written description of what this graph shows is found on the NEXT slide

Collisions and EnergyCollisions and Energy

As molecules collide, their kinetic As molecules collide, their kinetic energy is changed to potential energy is changed to potential energy and their bonds “stretch”energy and their bonds “stretch”

If the collision is effective, the bonds If the collision is effective, the bonds are broken, and the potential energy are broken, and the potential energy is changed back into kinetic as the is changed back into kinetic as the molecules separate.molecules separate.

This is one of the reasons there is a This is one of the reasons there is a “hump” in the enthalpy graph.“hump” in the enthalpy graph.

This description refers to the graph on the PREVIOUS slide

How a Catalyst WorksHow a Catalyst Works

A catalyst works by lowering the A catalyst works by lowering the activation energy of a reaction.activation energy of a reaction.

This means that less energy is This means that less energy is needed to initiate the reaction when needed to initiate the reaction when a catalyst is present.a catalyst is present.

It is like “lowering the hump” that the It is like “lowering the hump” that the reactants must get over. (see reactants must get over. (see diagram in Study guide, p.58)diagram in Study guide, p.58)

Graphic RepresentationGraphic RepresentationNone of the molecules have enough energy to climb over the activation energy barrier.NO REACTION!

Now let’s heat up the molecules so they move faster!

Some of the molecules now have enough energy to make it over the “activation energy barrier”These molecules REACT and form new products

The more molecules make it over the barrier, the faster the reaction is.

Graphic RepresentationGraphic RepresentationCool molecules again, but this time with a catalyst.

The Catalyst makes the

“barrier” lower so it is easier to

get over

Some of the molecules now have enough energy to make it over the “activation energy barrier”These molecules REACT and form new products

The more molecules make it over the barrier, the faster the reaction is.

Potential Energy GraphPotential Energy Graph withoutwithout and and withwith a catalyst a catalyst

Enthalpy of Reactants

Enthalpy of Activated Complex

Enthalpy of Products

ΔH (Heat of reaction)

Forward Activation Energy (EA)

(Without catalyst)

Reaction Progress (time) Forward

Pot

entia

l Ene

rgy

(ent

halp

y)

Reverse

Some reactions are reversible. If a reaction is reversed, then the Activation energy is read from the opposite side. The ΔH changes sign (but still has the same numeric value).

Reverse Activation Energy

(With catalyst)

Maxwell-BoltzmannMaxwell-Boltzmann Energy Distribution CurvesEnergy Distribution Curves

Not all molecules move at the same Not all molecules move at the same speed (ie. With the same kinetic speed (ie. With the same kinetic energy)energy)

Some are faster, some are slower, Some are faster, some are slower, most are average. If we could graph most are average. If we could graph their kinetic energies, we would get their kinetic energies, we would get something like this:something like this:

The Effects of Temperature and The Effects of Temperature and Catalyst on an Energy Distribution Catalyst on an Energy Distribution

GraphGraph A kinetic energy distribution graph A kinetic energy distribution graph

shows the energy of molecules at a shows the energy of molecules at a certain temperature. certain temperature. Most molecules have nearly average Most molecules have nearly average

kinetic energy (the “hump” energy), kinetic energy (the “hump” energy), Some molecules are slower than Some molecules are slower than

average (left of the hump)average (left of the hump) some molecules are faster than average some molecules are faster than average

(right of the hump). (right of the hump). The distribution is normally a “bell-The distribution is normally a “bell-

shaped” curve.shaped” curve.

Boltzmann-Maxwell Distribution Curve

Act

iva

tion

Ene

rgy

Ea The larger this shaded area is, the faster the reaction will be.

Effect of Temperature on Effect of Temperature on distribution curvedistribution curve

Kinetic Energy in kJ/mol

Num

ber

of m

olec

ules

Cold: no molecules reach activation energy

Warmer: some molecules reach activation energy

Hot: many molecules reach activation energy

Increasing temp.Moves peak

Forward.

Effect of a Catalyst on the Effect of a Catalyst on the Distribution CurveDistribution Curve

Kinetic Energy in kJ/mol

Num

ber

of m

olec

ules

Activation energy before catalyst.

No molecules reach activation.

Activation energy after catalyst.

Some molecules reach activation.

Catalyst movesActivation energy

back

Summary of Lesson 3

• Collision Theory– Reactions occur when molecules collide.– The collisions must be effective to make new

substances.– Activation energy is the minimum amount of energy

needed for a collision to be effective.

• There are 2 types of graph used to illustrate collision energy.– Enthalpy Graphs:

• Show enthalpy of reactants, activated complex and product

– Distribution Graphs• Show number of molecules at various energy levels.

Module 4, Lesson 4Module 4, Lesson 4

Rate Laws and Rate ConstantsRate Laws and Rate Constants

(for aqueous and gaseous (for aqueous and gaseous reactions)reactions)

In general, reactions proceed more In general, reactions proceed more quickly when the reactants are more quickly when the reactants are more concentratedconcentrated

TheThe concentration concentration of a reactant of a reactant is is represented byrepresented by putting putting square brackets square brackets around the formulaaround the formula of the reactant. of the reactant. Eg. [NaOH] means “the concentration of Eg. [NaOH] means “the concentration of

sodium hydroxide in moles per litre”sodium hydroxide in moles per litre”

Rates of Reaction and Rates of Reaction and ConcentrationConcentration

The Rate Constant for a The Rate Constant for a ReactionReaction

For a simple (first order) reaction with just For a simple (first order) reaction with just one reactant , the rate of the reaction can one reactant , the rate of the reaction can be found by multiplying a certain number by be found by multiplying a certain number by the concentration of the reactant.the concentration of the reactant.

This special number is called the This special number is called the rate rate constantconstant of the reaction, and is symbolized of the reaction, and is symbolized by by kk..

The rate of a simple reaction is therefore The rate of a simple reaction is therefore given by the formula: Rate= k[A] given by the formula: Rate= k[A] Where k is the rate constant, and [A] is the Where k is the rate constant, and [A] is the

concentration of the reactant.concentration of the reactant.

Example for Simple Reaction• A chemist has developed a process for

changing dissolved carbon dioxide into carbon and oxygen. Under a certain set of conditions, the k value for the reaction CO2(aq) C(s)+O2(g) is 0.955 L/s

• Find the rate of reaction when the concentration of carbon dioxide is:

a) [CO2]=2 mol/L

b) [CO2]=0.5 mol/L

Answer:1.91 mol/sec

Formula: rate = k [CO2]

Answer:0.478 mol/sec

Work 1: rate = 0.955 (2)

Work 2: rate = 0.955 (0.5)

Finding k for a 1st order reaction

• For a simple first order reaction

• Where:• r is the rate (given or determined by experiment)• [A] is the concentration of the one reactant (given

or determined by experiment)

][A

rk

Higher Order ReactionsHigher Order Reactions

In some reactions, the rate depends on the In some reactions, the rate depends on the product*product* of the concentration of two of the concentration of two reactants.reactants.

In other reactions, the rate may depend on In other reactions, the rate may depend on the the squaresquare of the concentration of a certain of the concentration of a certain reactant.reactant.

These are higher order reactions (2These are higher order reactions (2ndnd order, order, 33rdrd order etc.) order etc.) *I mean the mathematical product, the result of *I mean the mathematical product, the result of

multiplication, not the chemical product.multiplication, not the chemical product.

The Coefficients of the The Coefficients of the Equation Equation usuallyusually give the give the

“order”“order” The coefficient of a reactant is its order.The coefficient of a reactant is its order.

Note: there are exceptions to this rule!Note: there are exceptions to this rule! The sum of the coefficients is normally The sum of the coefficients is normally

the overall order of the reaction.the overall order of the reaction. In a rate law calculation, the In a rate law calculation, the

concentration of a reactant is raised to concentration of a reactant is raised to the power of its coefficient (or order)the power of its coefficient (or order)

PURE solids and liquids are not PURE solids and liquids are not

The Rate Law (formula)The Rate Law (formula) For the reaction: For the reaction:

a A + b B +… a A + b B +… p P p P Rate = k [A]Rate = k [A]aa[B][B]bb……

Where: Where: a,b are coefficients. a,b are coefficients. (numbers in front)(numbers in front) A, B are reactants. A, B are reactants. (chemical formulas)(chemical formulas) k is the rate constant of reaction.k is the rate constant of reaction. [A], [B] are reactant concentrations.[A], [B] are reactant concentrations.

Example for 2nd Order Reaction

• Under a certain set of conditions, the k value for the reaction 2H2O2 2H2O+O2 is 0.84 L2/mol-s

• Find the rate of reaction when the concentration of hydrogen peroxide is:

a) [H2O2]=2 mol/L

b) [H2O2]=0.5 mol/L

Answer:3.36 L/sec

Formula: rate = k [H2O2]2

Answer:0.21 L/sec

Work a): rate= 0.84(2)2

Work b): rate= 0.84(0.5)2

Finding k for a sample higher order reaction

• For sample reaction aA + bB +... cC +...

• Where:• r is the rate• [A] and [B] are the concentrations of reactants and

a and b are coefficients • The values of [A], [B] and r may be given in the

problem, or may be determined by experiment

...][][ ba BA

rk

Example of High Order Example of High Order ReactionReaction

2 2 HH22 + O + O22 22 H H220 0 the coefficients are in yellowthe coefficients are in yellow 2 2 HH22 + + 1 1 OO22 22 H H220 0 remember there is one not remember there is one not

shownshown

So the rate formula for this reaction So the rate formula for this reaction would be:would be: Rate= k [HRate= k [H22]]22 [O [O22]]1 1 Note: the 1 is not usually written!Note: the 1 is not usually written!

It is a 3It is a 3rdrd order reaction overall ( order reaction overall (2+1=2+1=33)) It is 2It is 2ndnd order with respect to hydrogen.[H order with respect to hydrogen.[H22]]22

It is 1It is 1stst order with respect to oxygen.[O order with respect to oxygen.[O22]]11

Example for 3rd Order Reaction

• Under a certain set of conditions, the k value for the reaction 2H2+O2 2H2O is 1.84 L/mol-s

• Find the rate of reaction when the concentrations are:

a) [H2]=3 mol/L,[O2]=2 mol/L

Formula: rate = k [H2]2 [O2]

Answer:33.1 mol/sec

work: rate = 1.84 [3]2 [2]

The Value of kThe Value of k

The value of the rate constant (k) is The value of the rate constant (k) is different for each reactiondifferent for each reaction, ,

However once it has been found for a However once it has been found for a reaction, it can be used whenever reaction, it can be used whenever that reaction occurs under the same that reaction occurs under the same conditions.conditions.

The value of k is set at specific The value of k is set at specific conditions (eg. temperature=25conditions (eg. temperature=25°°C, C, pressure = 101 kPa, etc.)pressure = 101 kPa, etc.)

The Units of kThe Units of k

The units of k vary depending on the problem. Units The units of k vary depending on the problem. Units are chosen to make the units of the rate work out as are chosen to make the units of the rate work out as desired usually to mol/sec (moldesired usually to mol/sec (molss-1-1) or mol/L) or mol/Lsec sec (mol(molLL-1-1ss-1-1))

Since the units are variable, they are frequently not Since the units are variable, they are frequently not given. Sometimes they are “reverse engineered” given. Sometimes they are “reverse engineered” from the units of the answer.from the units of the answer.

Common units are:Common units are: ss-1-1 (or per sec) (or per sec) 11stst order, answer in order, answer in

mol/Lmol/Lsecsec LLss-1-1 (or L/sec) (or L/sec) 11stst order, answer in mol/sec order, answer in mol/sec LL22molmol-1-1ss-1 -1 (or(or LL22/mol/molsec) sec) 22ndnd order, answer in mol/sec order, answer in mol/sec LL33molmol-2-2ss-1 -1 (or(or LL33/mol/mol22sec) sec) 33ndnd order, answer in mol/sec order, answer in mol/sec

My advice: Don’t worry too much about the units of My advice: Don’t worry too much about the units of k, think about what the units of the rate should be.k, think about what the units of the rate should be.

Summary of Lesson 4

• For the reaction aA+bB+…pP• Rate = k [A]a[B]b… where

• k = rate constant• [A] = concentration of reactant A (mol/L)• [B] = concentration of reactant B (mol/L)• a = order of reactant A• b = order of reactant B• P = the product(s) of the reactions

• Orders of a reactant are usually given by the coefficients of the equation.

Exceptions to the Rules!

• In some cases, the order of the reaction is NOT found from the coefficients!

• This happens when a reaction occurs in several steps, and one step is slower than the others.

• If this is the case, the question will give you a clue or information to that effect!

Problems Write rate law expressions for the following

reactions: 2Na+

(aq) + CaCO3 (aq) Na2CO3(aq) + Ca2+(aq)

3K+(aq) + PO4

3-(aq) K3PO4(aq)

2N2O(g) 2N2(g) + O2(g)

For the following reaction: NH4

+(aq) + NO2

-(aq)N2(g) + 2H2O(l)

the rate constant is .015 L2/mols. Calculate the initial rate if the concentration of NH4 is 3mol/L and the concentration of NO2

- is 4 mol/L Do the questions for module 4 in the student

study guide (p. 4-7 to 4-13.)

Answers

Rate=k[Na+]2[CaCO3]

Rate=k[K+]3[PO43-]

Rate=k[N20]2

Rate =k[NH4+][NO2]

=0.015L2/mols (3mol/L)(4mol/L)=0.18mol/s

Module 4, Lesson 5Module 4, Lesson 5

The usefulness of reaction ratesThe usefulness of reaction rates

Importance of Reaction Importance of Reaction Rates Rates

The study of reaction rates has led The study of reaction rates has led scientists to a better understanding scientists to a better understanding of the way in which reaction rates of the way in which reaction rates affect natural processes and the affect natural processes and the development of technologies to development of technologies to control reaction rates. This control reaction rates. This knowledge has affected our society, knowledge has affected our society, the economy and the environment. the economy and the environment. (Typical Ministry of Education Jargon…)

Examples of Useful Control of Examples of Useful Control of Reaction Rates in NatureReaction Rates in Nature

Chlorophyll acts as a catalyst during Chlorophyll acts as a catalyst during photosynthesis photosynthesis (example of natural catalyst)(example of natural catalyst)

Enzymes allow glucose to be metabolized at Enzymes allow glucose to be metabolized at reasonable temperatures reasonable temperatures (example of natural catalyst)(example of natural catalyst)

Strong HCl acid in stomach helps break up Strong HCl acid in stomach helps break up food food (example of concentration effect)(example of concentration effect)

Slower metabolism at low temperatures Slower metabolism at low temperatures allow some animals to conserve food and allow some animals to conserve food and oxygen during hibernation. oxygen during hibernation. (example of natural effect of (example of natural effect of temperature)temperature)

Examples of Technologies that Examples of Technologies that Use Control of Reaction Rates.Use Control of Reaction Rates.

Refrigeration slows the spoilage of foodRefrigeration slows the spoilage of food Preservatives inhibit the spoilage of foodPreservatives inhibit the spoilage of food Low temperature surgery reduces blood Low temperature surgery reduces blood

loss and oxygen use during operationsloss and oxygen use during operations Catalytic converters in cars reduce Catalytic converters in cars reduce

pollution by increasing oxidation of fuels.pollution by increasing oxidation of fuels. Enzymes in laundry detergents help Enzymes in laundry detergents help

remove stains.remove stains.

Additional ActivitiesAdditional Activities

Addison-Wesley ChemistryAddison-Wesley Chemistry Read pages 395 – 412Read pages 395 – 412 End of Chapter Exercises #17 – 25End of Chapter Exercises #17 – 25

Study GuideStudy Guide Pages 4-1 to 4-14 Pages 4-1 to 4-14

Module 4, Lesson 6Module 4, Lesson 6

EntropyEntropy

She’s Breakin’ Apart, Captain. I canna’ hold

‘er together.

Energy & Exothermic Energy & Exothermic ReactionsReactions

Exothermic reactions are driven by Exothermic reactions are driven by reducing reducing EnthalpyEnthalpy,, ie. the conversion ie. the conversion of potential energy into heat.of potential energy into heat. It is easier to reach the activation It is easier to reach the activation

energy for an exothermic reaction than energy for an exothermic reaction than for an endothermic one. for an endothermic one.

If energy was the only force driving If energy was the only force driving reactions, then only exothermic reactions, then only exothermic reactions would take place.reactions would take place.

• But, as we have seen in a demonstration, endothermic reactions DO occur.

• Sometimes they even occur spontaneously.

So what drives endothermic So what drives endothermic reactions?reactions?

There is another tendency, called There is another tendency, called entropy, that causes systems to entropy, that causes systems to become less orderly.become less orderly.

This tendency drives many This tendency drives many endothermic reactions. endothermic reactions.

Entropy = Entropy = D osi Rd E r

What is Disorder?What is Disorder? In nature it is easier for most systems In nature it is easier for most systems

to become disordered than it is for to become disordered than it is for them to become orderly. them to become orderly. Eg. If you spill loose LegoEg. If you spill loose Lego®® blocks on the blocks on the

floorfloor they are very unlikely to land in a neatly they are very unlikely to land in a neatly

stacked pyramid! Chances are they will stacked pyramid! Chances are they will spread out in a chaotic pattern.spread out in a chaotic pattern.

ProbabilityProbability(neatly stacked)(neatly stacked)

0.0000000000...0001%0.0000000000...0001%

ProbabilityProbability(a real mess)(a real mess)

99.99999...%99.99999...%

Entropy, a facetious Entropy, a facetious exampleexample

Don’t copyI don’t want you to show

this to your parents!

facetious [fuh-see-shuhs] –adjective

1. not meant to be taken seriously or literally: a facetious remark. 2. amusing; humorous. 3. lacking serious intent; concerned with something nonessential, amusing, or frivolous: a facetious person.

Clothes spread out on the floor have greater entropy Clothes spread out on the floor have greater entropy (or disorder) than clothes neatly placed in a hamper.(or disorder) than clothes neatly placed in a hamper.

Since nature follows the law of disorder, it is natural Since nature follows the law of disorder, it is natural for clothes to be spread out on the floor.for clothes to be spread out on the floor.

To put them neatly in the hamper requires you to To put them neatly in the hamper requires you to expend energy.expend energy.

And entropy tells you that they eventually will be And entropy tells you that they eventually will be spread out on the floor again.spread out on the floor again.

So when your mother asks why your clothes are on the So when your mother asks why your clothes are on the floor, you can blame ENTROPY—floor, you can blame ENTROPY— Its not just an excuse, it’s the law…Its not just an excuse, it’s the law…

……The The Second Law of Second Law of ThermodynamicsThermodynamics..

Yeah, that didn’t work with my mom eitherYeah, that didn’t work with my mom either!!

The Law of Disorder:The Law of Disorder:

Things tend to move spontaneously Things tend to move spontaneously in the direction of maximum in the direction of maximum disorder or chaos.disorder or chaos.

The only way to overcome chaos, is to The only way to overcome chaos, is to expend energy!expend energy!

One way to recover energy is to allow One way to recover energy is to allow a system to become more chaotic!a system to become more chaotic!

What is Entropy?What is Entropy? Entropy (S) is a measure of how Entropy (S) is a measure of how

disordered a system is.disordered a system is.

Examples:Examples: Scattered blocks have higher entropy than Scattered blocks have higher entropy than

neatly stacked blocks! neatly stacked blocks! (they are more (they are more disorderly)disorderly)

Gases have higher entropy than solids Gases have higher entropy than solids (their particles are more disorderly)(their particles are more disorderly)

A rusty wrecked car has more entropy than A rusty wrecked car has more entropy than a shiny new car. a shiny new car. (its pieces are more disorderly)(its pieces are more disorderly)

What Other Factors Affect What Other Factors Affect Entropy?Entropy?

The entropy of…The entropy of… a gas is greater than that of a liquid or a gas is greater than that of a liquid or

solid. ie:solid. ie: S Sgasgas > S > Sliquidliquid > S > Ssolidsolid. (so reactions . (so reactions favour gases)favour gases)

Lots of small particles is greater than a Lots of small particles is greater than a few large ones. (so reactions favour few large ones. (so reactions favour smaller particles) ie: smaller particles) ie: SSsmall particlessmall particles > S > Sbig particlesbig particles

Hot material is greater than a cool one Hot material is greater than a cool one since movement causes greater chaos, since movement causes greater chaos, ie:ie: S Shothot > S > Scoldcold

Standard Entropy TablesStandard Entropy Tables The symbol for total entropy is S.The symbol for total entropy is S. When the When the

entropy of one mole of substance is entropy of one mole of substance is calculated at a standard temperature* its calculated at a standard temperature* its called standard entropy and the symbol is called standard entropy and the symbol is SSo o

oror “S-naught“S-naught”” The standard entropy of different substances The standard entropy of different substances

has been calculated, and recorded in tables.has been calculated, and recorded in tables. There is a small standard entropy table on There is a small standard entropy table on

page 407 of your old text book.page 407 of your old text book. The unit of standard entropy is joules per The unit of standard entropy is joules per

kelvin per mol (J/Kkelvin per mol (J/Kmol). The mol). The total entropy is total entropy is given by:given by: S=(S S=(Soo)(T )(T (kelvins)(kelvins))(n )(n (mol)(mol) ) ) S= SS= SooTnTn

*whenever possible, the reactions are carried out at 25oC, 298 kelvins

Reviewing those symbolsReviewing those symbols

S = Total EntropyS = Total Entropy = total disorder of a = total disorder of a material.material.

SS00 = Standard Entropy = Standard Entropy = entropy of = entropy of one mole of substance. (based on its one mole of substance. (based on its state and reactions at 25C)state and reactions at 25C)

S = SS = S00TnTn total entropy = standard entropy total entropy = standard entropy ×× kelvin kelvin

temperature temperature ×× #moles #moles

To keep it simpler…To keep it simpler…

When the temperature constant we When the temperature constant we usually work with standard entropy, usually work with standard entropy, not total entropy not total entropy It’s easier since we don’t have to It’s easier since we don’t have to

multiply by the kelvin temperaturemultiply by the kelvin temperature Of course, if the temperature changes Of course, if the temperature changes

during a problem, you have to use total during a problem, you have to use total entropy instead.entropy instead.

Entropy and Enthalpy are the two forces Entropy and Enthalpy are the two forces that determine if a reaction can take that determine if a reaction can take place.place.

Sometimes enthalpy and entropy oppose Sometimes enthalpy and entropy oppose each other, sometimes they reinforce each other, sometimes they reinforce each other!each other!

Reinforce: Reinforce: Enthalpy lower Enthalpy lower Entropy higher Entropy higherEnthalpy higher Enthalpy higher Entropy lowerEntropy lower

Oppose: Oppose: Enthalpy lower Enthalpy lower Entropy lower Entropy lowerEnthalpy higher Enthalpy higher Entropy higher Entropy higher

Reactions which are known to take Reactions which are known to take place are called spontaneous place are called spontaneous reactions. They reactions. They WILLWILL happen easily happen easily when circumstances are right.when circumstances are right. Another interpretation is that Another interpretation is that

spontaneous reactionsspontaneous reactions start easily or start easily or automatically. automatically.

Non-spontaneousNon-spontaneous reactions are difficult or reactions are difficult or impossible to get started.impossible to get started.

Enthalpy, Entropy and Enthalpy, Entropy and Spontaneous ReactionsSpontaneous Reactions

EnthalpyEnthalpy(heat content)(heat content)

EntropyEntropy(disorder)(disorder)

SpontaneouSpontaneouss

Reaction?Reaction?

ReversibleReversible??

DecreasesDecreases(exothermic) (exothermic)

IncreasesIncreases

YESYES NONO

IncreasesIncreases (endothermic)(endothermic)

IncreasesIncreases

Perhaps*Perhaps* YESYES

DecreasesDecreases (exothermic) (exothermic)

DecreasesDecreases

Perhaps*Perhaps* YESYES

IncreasesIncreases (endothermic)(endothermic)

DecreasesDecreases

NONO YES.YES. (Spontaneous)(Spontaneous)

*Under the right conditions

Exercises: (refer to P.407 in text)

List the following from most disorderly to most orderly: Calcium carbonate, diamond, graphite,

ammonia gas, liquid water, chlorine Calculate the total entropy of 36

grams of water at 27°C Find the change in standard entropy

as one mole of water evaporates at 25°C: H2O(l)H2O(g)

Answer all parts of problem 8 on page 407

Chlorine (223),Ammonia (192), Calcium Carbonate (88.7), liquid water (69.9), Graphite(5.69), Diamond(2.43)

Total entropy is equivalent to 41964 J The change in standard entropy of water

is 118.8 J/Kmol #8a: 88.7 to 253.3 =164.6 J/Kmol #8b: 466.2 to 139.8= -326.4 J/Kmol #8c: 353.6 to 373.4 = +19.8 J/Kmol

The Laws of Thermodynamics The Laws of Thermodynamics (simplified versions)(simplified versions)

11stst Law Law: In any chemical or physical : In any chemical or physical process, the total energy remains the process, the total energy remains the same. same. (AKA—conservation of Energy)(AKA—conservation of Energy)

22ndnd Law Law: The entropy or disorder of an : The entropy or disorder of an isolated system will increase over time. isolated system will increase over time. (AKA—law of disorder)(AKA—law of disorder)

33rdrd Law: Law: As the temperature of a system As the temperature of a system approaches absolute zero, the entropy will approaches absolute zero, the entropy will approach a minimum. (ie: molecular approach a minimum. (ie: molecular motion stops at 0K and the material motion stops at 0K and the material becomes as orderly as it will ever get.)becomes as orderly as it will ever get.)

11stst Law Law: In any chemical or physical : In any chemical or physical process, the total energy remains the process, the total energy remains the same. same. (AKA—conservation of Energy)(AKA—conservation of Energy)

22ndnd Law Law: The entropy or disorder of an : The entropy or disorder of an isolated system will increase over time. isolated system will increase over time. (AKA—law of disorder)(AKA—law of disorder)

33rdrd Law: Law: As the temperature of a system As the temperature of a system approaches absolute zero, the entropy will approaches absolute zero, the entropy will approach a minimum. (ie: molecular approach a minimum. (ie: molecular motion stops at 0K and the material motion stops at 0K and the material becomes as orderly as it will ever get.)becomes as orderly as it will ever get.)

Enrichment

Warning: Depressing Content Ahead!(close your eyes if you don’t like gloomy predictions)

***Start this 2 minutes before end of class***

“The result would inevitably be a state of universal rest and DEATH, if the universe is finite and left to obey existing laws… running down like a clock, and stopping for ever.”

“The result would inevitably be a state of universal rest and DEATH, if the universe is finite and left to obey existing laws… running down like a clock, and stopping for ever.”

In 1862, after discovering the concept of entropy, Lord Kelvin contemplated its effect on the universe as a whole.

In 1862, after discovering the concept of entropy, Lord Kelvin contemplated its effect on the universe as a whole.

Next slide

ENTROPY & THE COLD DEATH OF THE UNIVERSE

If entropy continues, in 1014 years all the hydrogen in the Universe will be

used up. New stars will stop forming. The old ones will start to burn out,

collapse into black holes, or explode as supernovae

In 1015 years, the orbits of all the planets circling the

dead stars will decay.

The planets will crash into their suns, or be thrown

into deep space.

In 1016 years the stars will be sucked into black

holes or flung from orbit and torn apart, leaving only stellar debris and

black holes.

In 1040 years all the protons of all atoms will have decayed, leaving

only radiation and black holes

For the next 10100 years the black holes will

disintegrate leaving only Hawking Radiation

At this point only photons will remain. No matter, no black holes, nothing solid.

After 101000 years even the energy will fade away as

the universe becomes colder and more

disordered. When the Universe reaches -273°C

all that remains is…

ENTROPY

Have a Nice Day!

• Textbook Reading Assignment: – pp. 395 to 416 (first half of chapter 17)

• Textbook Assignment Exercises: – pp. 424 to 425, #17 to #33. – Due date: next class

• Study Guide Reading Assignment– Module 4 (pp.41 to 414)– Try all self-correcting exercises in module 4

Answers to Questions & ProblemsTextbook pages 424-425

17. Chemical reactions proceed when molecules collide with sufficient energy (activation energy) to rearrange the molecules and create new products.

18. Not all collisions lead to new products. The energy of the collision must be high enough to break the old bonds and allow new ones to form (activation energy)

19. The rate of a reaction is the change in the amount (or concentration) of a reactant over a change in time. ie. ΔA

Δt

20. How does each affect the rate?a. Temperature: an increase in temperature will

increase the rate of most reactions, because molecules will collide with greater energy.

b. Concentration: increasing the concentration of reactants will increase the rates of most reactions, because there will be more molecules to collide

c. Particle size: Smaller particles will react more quickly than larger particles because they will have greater surface area

d. Inhibitor: an inhibitor slows down a reaction by preventing a catalyst from working or by having the reverse effect of a catalyst.

21. The rate of a spontaneous reaction is increased by a catalyst. Catalysts allow more effective collisions by lowering the activation energy.

22. Which of these statements are true? Answer: only c is always true! (explanations below)

a. Most chemical reactions can be increased by raising the temperature, but ones involving enzymes may not, since enzymes have specific temperatures at which they are most effective.

b. All molecules must “climb” the activation energy barrier. If a reaction is exothermic, the energy it gives off will heat up the remaining molecules and help them to make it over, but they still have to do it.

c. Enzymes are biological catalysts.

23. At first the gas particles do not have enough kinetic energy to “climb” the activation energy barrier. When a flame is held near the gas, the gas molecules move faster as they heat up. Once enough of the have enough energy the reaction begins

24. Explain the terms:a. Specific rate constant: a proportionality constant

that relates the concentration of the reactants to the rate of a reaction.

b. First order reaction: is a reaction whose rate is proportional to the concentration of only one reactant

c. Rate law: an expression relating the rate of a reaction to the concentration of its reactants

25. Rate = k[NO(g)][O3(g)]26. The reaction will be 75% complete after 100 minutes

(or 1 hour 40 min)27. Define the terms:

a. Elementary reaction: the reactants are converted to products in a single step.

b. Intermediate: a product of the first step of a reaction that immediately becomes the reactant of another step of the reaction

c. Reaction mechanism: all the elementary reactions of a complex reaction.

28. Sketch:

the EA for the first reaction step

must be smaller than the EA

for the the 2nd reaction step..

Balanced Equation: 2NO + O2 2 NO2

29.Entropy is a measure of the disorder of a system.

30. If the products are more orderly, then the entropy change is unfavourable for the reaction.

31.The lower entropy is in…a. The completed jigsaw puzzle (not the pieces)

b. The ice (not the water)

c. The sodium chloride crystals (not the solution)

d. The house (not the pile of bricks and lumber)

32. The entropy change is:a. positive (or an increase, greater or forward)b. negative (or decrease, less, or backwards)c. positived. negative

33. ΔS0 = S0 products – S0 reactants

= 1mol (96.4J/kmol) + 3/2 mol(205.0J/kmol) – 1mol (149.2J/kmol

= 403.9 J/K – 149.2 J/K= 254.7 J/K

There were 32 answers • Key: any mistake, omission, missing unit, etc.

counts as an error, but count no more than one error per “lettered” question (up to 3 errors in question 22, 24, or 27, up to 4 errors in question 20, 31 or 32, just one error other questions)

– No errors at all 5+– One error or omission 5– 2 or 3 errors or omissions 4+– 4, 5, or 6 errors 4– 7, 8, 9 or 10 errors 3+– 11 to 16 mistakes 3– 17 to 24 mistakes 2– 25 or more mistakes 1– Did not do the exercises 0

Fighting FireFighting Fire

An example of how controlling An example of how controlling rates of reaction can be useful.rates of reaction can be useful.

• Identify the factors that can influence the rate of a combustion reaction

• Fuel: Carbon-rich compounds, in gaseous state

• Oxygen: higher O2 Concentration, faster combustion

• Temperature: required for evaporation/sublimation of fuel and for the activation energy.

ChainChemicalReaction

Fuel + Oxygen CO2 + other products(carbon source) (or other oxidizer) (carbon dioxide) (water vapour, pollutants)

Wood, oil, coal, gasoline, wax,

alcohol and other organic material. VAPOURIZED

Fuel burns faster!

In natural fires, oxygen is always the oxidizer.

In the laboratory fluorine, chlorine or nitrous oxide can be

used as oxidizers

Every substance has an ignition temperature. At that temperature the

material begins to vapourize, and there is enough heat to provide the required

activation energy to start combustion.The ignition temperature of paper =233°C

Fire fighting techniques

• Removing the heat (temperature)– Spraying water on the fire– Using other coolants or refrigerants

• Removing or interrupting the fuel– Backfires burn out underbrush– Dig ditches, moats create a fuel-free zone– Venturi tube separate fuel from heat– “wet jet” separates fuel from heat (water keeps

oxygen from fanning the flames)• Removing oxygen

– Using CO2 to lower O2 concentration– Using foam or other material to smother fire– Using an explosive to momentarily use up oxygen

Examples in Action

Water has a huge capacity to absorb heat, due to its high specific heat capacity (c=4.19). Also, its vaporization is an endothermic change, cooling materials below their kindling temperatures as it evaporates.

Foam does not cool materials well, but it can prevent air

from reaching a fire and then choke the fire of oxygen.

(provided the foam doesn’t burn away too quickly– its not much

good on really hot fires.)

Lighting a backfire burns away underbrush from the path of a larger fire. When the larger fire reaches the burnt area, there is no more fuel for it to burn

A small flame, like a candle or a match, can be blown out

Assignment

• Chapter 8, page 231 #1 to 10

• Chapter 9, pp. 243-244 #1 to 10

• Chapter 10, pp. 272-273 #1 to 10

Answers to Questions Page

1. The advantage to measuring the reaction in the first moments is that the reaction is usually at its fastest rate before the reactants have been used up

2. CaCO3 + 2 HCl CaCl2 + CO2 + H2O

a. You could measure the mass of the CaCO3 used up and divide by the time it took, or you could measure the volume of CO2 produced and divide it by the time.

b. The rate is 2.0 g ÷ 85 s = 0.024 g/s (or 2.4x10-2 g/s)

or 0.05 mol 85 s = 0.00024 mol /s (or 2.4x10-4 mol/s)

3. Two slow reactions are the oxidation of iron and the decomposition of plastic in the environment (or any two sensible answers). Two fast reactions are fireworks and the precipitation of (or any two sensible answers).

4. C3H8 + 5 O2 3 CO2 + 4 H2O

a. ..

b. ..

5. Since the concentration of a reactant decreases as the reaction progresses, Δ[reactant] is always negative. Since never goes backwards, Δt is always positive. In order to get a positive rate for Δ[reactant] / Δt we must multiply it by -1,

t

OH

t

CO

t

O

t

HCr

][

4

1][

3

1][

5

1][ 22283

OHCOOHC rrrrr22283 4

1

3

1

5

1

6. For the rate as a function of the decrease in the concentration of a reactant to be equal to the rate of increase in the concentration of the product, the stoichiometric ratios (ie. Coefficients) must be equal.

7. .

a. .

b. .

c.

8. Answer (b) is correct. The stoichoimetric ratio determines the rate ratio, so 2B2O3 4B + 3O2 means:

t

HI

t

I

t

Hr

][

2

1][][ 22

t

OH

t

O

t

Hr

][

2

1][][

2

1 222

t

OH

t

Br

t

H

t

BrO

t

Brr

][

3

1][

3

1][

6

1][][

5

1 223

...24...4

1

2

132

32 sorrsorr

BOBBOB