Chemical KineticsUnit 12 Part 2

• Kinetics – Kinetics – the study of the speeds of chemical reactions and the mechanisms by which reactions occur.

• Also referred to as “Reaction RatesReaction Rates”

• Rates of chemical change usually are expressed as the amount of reactant forming products per unit time.

∆[reactants]∆ time

Kinetics Kinetics

• How do reactions really occur?How do reactions really occur?– Reactant particles MUST Reactant particles MUST collidecollide!!

• Rates of chemical reactions are described in a Rates of chemical reactions are described in a model called model called collision theory.collision theory.

– Studies show most molecular collisions do NOT Studies show most molecular collisions do NOT result in a reaction…why??result in a reaction…why??

• Atoms, ions, and molecules react to form Atoms, ions, and molecules react to form products only when they collide with the products only when they collide with the proper orientation proper orientation and and sufficient energy.sufficient energy.

Collision TheoryCollision Theory

• EffectiveEffective collisions are defined by two collisions are defined by two conditions: conditions: – exactly the right exactly the right orientationorientation to react (head on to react (head on

collision is best) collision is best) – enough enough energyenergy to to break bonds and form break bonds and form new new

ones ones • The minimum amount of energy required is The minimum amount of energy required is

called the reaction’s called the reaction’s activation energy.activation energy.– The activation energy is a The activation energy is a barrierbarrier that reactants that reactants

must get over to reactmust get over to react– The higher the barrier the larger the amount of The higher the barrier the larger the amount of

energy needed for the reaction to proceedenergy needed for the reaction to proceed

Collision TheoryCollision Theory

• During a reaction, a particle that is neither During a reaction, a particle that is neither reactant nor product forms momentarily, reactant nor product forms momentarily, called an called an activated complex activated complex

– if there is sufficient energy if there is sufficient energy – and if the atoms are oriented properlyand if the atoms are oriented properly

• An activated complex is a kind of An activated complex is a kind of transition molecule transition molecule which has similarities which has similarities to reactants & productsto reactants & products

– An activated complex is the arrangement of An activated complex is the arrangement of atoms at the peak of the activation-energy atoms at the peak of the activation-energy barrier.barrier.

Collision TheoryCollision Theory

• Collision theory Collision theory explains why some explains why some naturally occurring reactions are naturally occurring reactions are immeasurably slow at room temp.immeasurably slow at room temp.

– Carbon and Oxygen react when charcoal Carbon and Oxygen react when charcoal burns, but this reaction has a high burns, but this reaction has a high activation energyactivation energy

– At room temp, the collisions of oxygen and At room temp, the collisions of oxygen and carbon molecules aren’t energetic enough carbon molecules aren’t energetic enough to reactto react

– But the reaction can be helped along a But the reaction can be helped along a number of waysnumber of ways

Collision TheoryCollision Theory

• It is possible to It is possible to vary the conditionsvary the conditions of the of the reaction, the rate of almost any reaction reaction, the rate of almost any reaction can be can be modifiedmodified

o collision theory can help explain why the collision theory can help explain why the rates can be modifiedrates can be modified

• Several strategies can be used to speed Several strategies can be used to speed up reactions:up reactions:

o Increase the Increase the temperaturetemperatureo Increase the Increase the concentrationconcentrationo Decrease the Decrease the particleparticle sizesizeo Employ a Employ a catalystcatalyst

Reaction RatesReaction Rates

• IncreasingIncreasing the temperature the temperature speeds up speeds up the the reaction, while reaction, while loweringlowering the temperature the temperature slows down slows down the reactionthe reaction

• For every 10For every 10ooC increase in temperature, C increase in temperature, the reaction rate usually DOUBLES!the reaction rate usually DOUBLES!

• Recall, temperature is directly Recall, temperature is directly proportional to average kinetic energy: proportional to average kinetic energy:

– Average KE = ½ x mass x velocityAverage KE = ½ x mass x velocity22

• At a higher temperature, there are more At a higher temperature, there are more effective collisions effective collisions – more molecules – more molecules have the velocity / energy needed to react have the velocity / energy needed to react

TemperatureTemperature

• Just sitting out, charcoal Just sitting out, charcoal does not react does not react at a measurable rateat a measurable rate– However, when a However, when a starter flame starter flame touches the touches the

charcoal, the temperature is increased so charcoal, the temperature is increased so atoms of reactants collide with atoms of reactants collide with higher higher energyenergy and and greater frequencygreater frequency

– Some of the collisions are high enough in Some of the collisions are high enough in energy that the product COenergy that the product CO22 is formed is formed

o The energy The energy releasedreleased by the by the exothermic reaction then supplies exothermic reaction then supplies enough energy to get more C and enough energy to get more C and OO22 over the over the activation-energy activation-energy barrierbarrier• Remove the starter flame:Remove the starter flame: the the

reaction will continue on its own.reaction will continue on its own.

• The more reacting particles you have in a The more reacting particles you have in a given volume, the given volume, the higher the rate of higher the rate of reaction.reaction.

• Cramming more particles into a fixed Cramming more particles into a fixed volume increases the volume increases the concentrationconcentration of of reactants:reactants:

• Increasing the concentration, increases Increasing the concentration, increases the the frequencyfrequency of the collisions, and of the collisions, and therefore therefore increasingincreasing the reaction rate. the reaction rate.

• ***For gases, increasing ***For gases, increasing pressurepressure and/or and/or decreasing volume increases concentration.decreasing volume increases concentration.

ConcentrationConcentration

• The smaller the particle size, the The smaller the particle size, the larger larger the surface areathe surface area for a given mass of for a given mass of particles. Effectively is increasing the particles. Effectively is increasing the concentrationconcentration..

• An increase in surface area increases the An increase in surface area increases the amount of the reactant amount of the reactant exposed for exposed for collisioncollision to take place… to take place…

– Which increases the collision Which increases the collision frequencyfrequency and the reaction rate.and the reaction rate.

• Methods:Methods:– Grinding the reactants into a powderGrinding the reactants into a powder– Dissolving in a solvent.Dissolving in a solvent.

Particle SizeParticle Size

•A catalyst is often the best way to speed up an reaction.

• In fact, some reactions simply will not go forward measurably without one.

•A catalyst is a substance that increases the rate of a reaction without being changed or used up during the reaction

•The key is that they permit reactions to proceed at lower activation energy than is normally required

CatalystCatalyst

•Catalysts lower the required activation energy by providing an alternative path or arrangement for the molecules to react.

•Not always understood why certain catalyst behave the way the do…they just do!

•By lowering the Ea threshold, more molecules in the system will have the required energy to surmount the barrier.

•A negative catalyst has the opposite effect. It interferes with effective collisions and increases the Ea required.

CatalystCatalyst

• The type of reactant substances also plays a role in reaction rates.

• Weak bonds = easier to break, reactants with weak bonds will react faster.

– Essentially results in a low activation energy.

• Strong bonds = harder to break, reactants with strong bonds will react slower (high Ea)

• Electronegativity and ionization energy also plays a role.

Nature of ReactantsNature of Reactants

• Reaction Reaction Rate LawsRate Laws relate the speed of a relate the speed of a reaction to the reaction to the concentrationsconcentrations of the of the reactantsreactants..

– Determined from experimentDetermined from experiment

• The General Rate LawThe General Rate Law

– For reaction: xA + yB For reaction: xA + yB Products Products

– Rate = k[A]m[B]n

• Example: NExample: N22 + 2H + 2H22 2NH3 2NH3

– Rate Law: Rate = k[NRate Law: Rate = k[N22]]11[H[H22]]22

Reaction Rate LawsReaction Rate Laws

Reaction MechanismReaction Mechanism• Most chemical reactions consist of

complex multi-step sequences of two or more simpler (elementary) reactions.

• For example: Ozone in the upper atmosphere can decompose into Oxygen: 2O3 3O2

• Happens after UV light liberates Cl atoms from Happens after UV light liberates Cl atoms from chlorofluorocarbons and breaks down ozone:chlorofluorocarbons and breaks down ozone:

Elementary Step 1: Cl + OElementary Step 1: Cl + O33 ClO + O ClO + O22

Elementary Step 2: OElementary Step 2: O33 O + O O + O22

Elementary Step 3: ClO + O Elementary Step 3: ClO + O Cl + O Cl + O22

Overall Net Reaction: Overall Net Reaction: 2O3 3O2

UV

Reaction MechanismReaction Mechanism• Since ClO molecules and O atoms were formed

in one elementary step and then used up in another, they are called “intermediates”.

• Like catalysts, they do not appear in the net chemical equation.

• Intermediates live much longer than activated complexes

• Many intermediates exist only for a fraction of a second: often measured in femtoseconds (1 x 10-15 seconds).

• Caltech scientists have used laser flashes to record the formation of intermediates.

• For a complex reaction, its Energy Diagram For a complex reaction, its Energy Diagram has a series of has a series of hills & valleyshills & valleys

– The peaks correspond to the energies of the The peaks correspond to the energies of the activated complexesactivated complexes

– Each valley represents an Each valley represents an intermediate product intermediate product which becomes a reactant in the next stage of which becomes a reactant in the next stage of the reactionthe reaction

• For the complex reaction:For the complex reaction:

• The elementary steps are:The elementary steps are:

Reaction MechanismReaction Mechanism

H2(g) + 2ICl (g) I2(g) + 2HCl(g)

1) H2 + 2ICl(g) ICl(g) + HCl(g) + HI(g) 2) ICl(g) + HCl(g) + HI(g) I2(g)

+2HCl(g)

Energy Diagram with Energy Diagram with IntermediatesIntermediates

Reaction MechanismReaction Mechanism• Rate Determining Step – every

reaction has one step that is the slowest.

• A reaction can not go any faster than the slowest step, so it “determines” the rate.

• Complex reaction: 2NO + 2H2 N2 + 2H2O

• Step 1: 2NO N2O2

• Step 2: N2O2 + H2 N2O + H2O

• Step 3: N2O + H2 N2 + H2O

• Thus, the second step is the rate determining step.

Fast!

Fast!

Slow

Reaction MechanismReaction Mechanism• What would be the rate determining step for the What would be the rate determining step for the

reaction:reaction:

4A + 2B + 2C 4A + 2B + 2C 2D + 3E 2D + 3E

Step 1: 2A + B Step 1: 2A + B F + D F + D FASTFAST

Step 2: 2C + B Step 2: 2C + B D + 2F D + 2F SLOWSLOW

Step 3: 2A + 3F Step 3: 2A + 3F 3E 3E FASTFAST

Which is the rate determining step? (Step 2)Which is the rate determining step? (Step 2)

Which reactant concentrations have an effect on rate?Which reactant concentrations have an effect on rate?

(C and B only)(C and B only)