Chemical Kinetics and Collision Theory

Aim KE1

How do chemical reactions actually happen?

• Collision Theory - A chemistry theory that explains how chemical reactions occur through effective collision of particles– Effective collision - one in which the colliding

particles approach each other at the proper angle and with the proper amount of energy

• The greater the rate of effective collisions, the greater the reaction rate (speed of reaction)

• An Analogy - Collision theory is similar to bowling:

– You throw a ball very weakly – do you knock all the pins

down? Why?– You throw the ball at the

10 pin (back right corner)– do you knock down all the

pins? Why?

Reaction mechanisms• Chemical reactions where only two atoms/molecules

collide to make a product are rare• Most reactions go through a series of steps called

the reaction mechanism• Example of steps in a reaction: the formation of

ammonia (a process called the Haber Process)

3H2 + N2 2NH3 + 92.2 kJ

• There are three basic steps in the reaction:– Step 1: three H2 are broken up into six H atoms

– Step 2: one N2 is broken up into two N atoms

– Step 3: each N atom collides with three H atoms to make NH3 (three separate collisions)

Rates of Reaction Mechanisms• Each step requires collisions• Some happen quickly, some happen more slowly• In the Haber Process, 3H2 + N2 2NH3 + 92.2 kJ

• Step 1: the break up of H2 is FAST• Step 2: the break up of N2 is FAST• Step 3: the formation of NH3 from the H and N is

SLOW

• The slowest step of the reaction mechanism is called the rate determining step

• It determines how fast the overall reaction will take to happen

Transition State Theory As a chemical reaction progresses

»Atoms rearrange themselves forming intermediate products (not quite reactants, but not quite products)

»Example – In the reaction of hydrogen with oxygen to form water at left

These intermediate products are called transition state complexes or activated complexes

Transition State Theory• Activated complexes

– exist for only brief periods of time while the atoms rearrange themselves

– have high energy due to their formation by high energy collisions

– they are unstable but need to form in order to make the final product(s)

– the energy needed to form the activated complex is called the activation energy

• Activation energy – gives the chemical system

enough energy to reach the activated complex

– The reaction can than continue to completion, with new products formed

– Activation energy varies with • the nature of

substances in the reaction

• The type of reaction (endothermic vs exothermic reactions)

If reactions didn’t need activation energy, they would all simply happen

But this assumes we don’t have an activated complex

• Transition State Theory–The energy changes in a chemical reaction can be

shown in a graph called a potential energy diagram –the high energy product is called an activated

complex or a transition state complex–the energy needed to form the activated complex is

the activation energy–The reaction pathway

represents time or the progress of the reaction

–The potential energy shows the amount of energy the chemicals have at different points in the reaction

Reading a Potential Energy Diagram

• The PE Diagram shows the amount of energy that changes in a chemical reaction

• PEreactants = energy the reactants possess already

• PEproducts = the energy the products possess after the reaction occurs

• Ea = the activation energy needed to reach the activated complex

H = the heat of reaction – the heat change between the products and the reactants

Rates of Reaction

Aim KE 2What is the difference between exothermic and endothermic reactions?

Mr. Foley, may I be excused? My brain is full.

Role of Energy in Reactions • Enthalpy is the difference between the potential

energy (PE) of the products and the reactants

H = Hproducts – H reactants

• heat of reaction or enthalpy (H) – The heat energy (enthalpy) released or absorbed

during a chemical reaction – In an exothermic reaction

• heat is released• H is negative ( - H )

– In an endothermic reaction • heat is absorbed• H is positive ( + H )

53.0 kJ + H2 + I2 2 HI

H2 + I2

2 HI

Po

ten

tial

En

erg

y

Reaction Pathway

Activation EnergyH = + 53 kJ

• In an endothermic reaction – heat is added to reactants in order to make the

products– for the heat of reaction for an endothermic reaction

is + H (heat of reaction)– energy was added to the reactant side of the

reaction

Example2 H2 + O2 2 H2O + 484 kJ H2 + O2

2 H2OPo

ten

tial

En

erg

y

Reaction Pathway

H = - 484 kJ

Activation Energy

• In an exothermic reaction – heat is removed from the reactants in order to make

the products– for the heat of reaction for an exothermic reaction

is - H – energy was removed and comes out as a product

Potential energy diagram #1• Match each letter with the appropriate label using the word bank below:1. PE of Reactants = ____ 3. Activation Energy = _____2. PE of Products = ____ 4. Heat of reaction = _____5. Mark an X where the activated complex is located.6. What type of reaction is this, endothermic or

exothermic? ________________ 7. How do you know? _________________

_________________8. Is the heat of reaction positive or negative?

__________________

Potential energy diagram #2• Match each letter with the appropriate label using the word bank below:1. PE of Reactants = ____ 3. Activation Energy = _____2. PE of Products = ____ 4. Heat of reaction = _____5. Mark an X where the activated complex is located.6. What type of reaction is this, endothermic

or exothermic? ________________ 7. How do you know? _________________

_________________8. Is the heat of reaction positive or negative?

__________________

Table I – Heats of Rxn• Various reactions are

presented in this chart• All are at standard

pressure and room temperature (298 K)

• The H represents the energy absorbed (+) or released in the reaction (-)

• This energy represents the total energy of all the moles formed in each reaction

• Examples from Table I

Combustion of methaneCH4(g) + 2O2(q) CO2(g) + 2H2O(g) + 890.4 kJ

H = - 890.4 kJ; an exothermic reactionFormation of water vapor2H2(g) + O2(q) 2H2O(g) + 483.4 kJ

H = - 483.4 kJ; an exothermic reactionSynthesis of nitrogen dioxide66.4 kJ + N2(g) + 2O2(g) 2 NO(g)

H = + 66.4 kJ; an endothermic reactionDissolving (ionization) of a salt (cold pack salt)

H2O

14.78 kJ + NH4Cl(s) ---> NH4+

(aq) + Cl- (aq)

H = +14.78 kJ; an endothermic reaction

Using Table I to fill in the chart below:

Reaction Exothermic or Endothermic?

H value

C3H8(g)+5O2(g)3CO2(g) +4H2O(l) Exothermic -890.4 kJ

N2(g) + 3H2(g) 2NH3(g) Exothermic -91.8 kJ

H2O(g)

LiBr(s) Li+(aq) + Br-(aq)

Exothermic - 48.83 kJ

The formation of liquid water from hydrogen and oxygen gas

Exothermic -571.6 kJ

The dissolving of sodium chloride in water to form sodium ions and chlorine ions

Endothermic +3.81

Rates of Reaction

Aim KE 3What speeds up chemical reactions?

Factors affecting Rates of Chemical Reactions• Just like when we spoke about the rate of

dissolving and the effects on solubility• Various factors affect the speed or rate of a chemical

reaction– Concentration – the number of particles in a

given sample – Pressure – affects only gases– Nature of Particles – how large they are and how

well they react– Surface Area – the amount of molecular surface

exposed between the reactants– Temperature – the amount of kinetic energy

Increasing Chemical Reaction Rates

1 - Increase the concentration of the reactants– Increasing concentration could mean adding more

reactants (increasing molarity)– or increasing the density by making the volume of

the reaction vessel smaller– More collisions = increasing reaction rate

Increasing Rates of Chemical Reactions

2 - Increase the pressure of gaseous reactants– Squeezing together gases increases concentration– Increasing pressure forces reactants closer

together - but only for gases

Increasing Rates of Chemical Reactions3 - Nature of Reactants

– Chemical reactions occur by breaking and rearranging existing bonds.

– The less electrons that need to be rearranged, the faster the reaction is.

– As a result, reactions between ionic substances in aqueous solution, such as double replacement reactions, are rapid at room temperature

– Ex: KI (aq) + Pb(NO3)2(aq) KNO3 (aq) + PbI2(solid)

– On the other hand, reactions in which covalent bonds are broken, such as the decomposition of hydrogen peroxide, occur slowly at room temperature.

Increasing Rates of Chemical Reactions

4 - Increasing surface area– More particles available to react with each other

Increasing Rates of Chemical Reactions

5 - Increasing temperature– Increases average kinetic energy of particles– Increases energy and number of collisions

75oC

R 50oC

A 25oC

T

E

Time

Increasing Rates of Chemical Reactions

6 – Catalysts• Catalysts speed up reactions without being

permanently altered• Enzymes in biological systems are organic (carbon

based) catalysts• They change the

pathway the reaction mechanism follows

• so less activation energy is required to get to the activated complex or transition state

EquilibriumAim KE 4d – What is equilibrium in chemistry?

• Equilibrium– A balance between two opposing forces or objects– Tends to be static (doesn’t change once

reached)– Ex – gravity vs your leg bones or a building’s

supports

• In chemistry– Dynamic equilibrium occurs– So that the balancing forces are always

adjusting themselves– Chemical reactions are reversible– They can occur in either direction, so one

opposing reaction is exothermic, the other is endothermic

• Under conditions of dynamic equilibrium –the rates of the forward and reverse reactions

are equal–A balance point occurs as the two reactions

compete–But the amounts of

materials on each side of the reaction may NOT be equal

–and the point of equilibrium can change with

changing conditions

• Types of Equilibrium - Phase Equilibrium– We know phase changes are reversible– Ice in a container at 0oC - S L – Rate of melting equals rate of freezing– Water in a container at 25oC - L G– Rate of evaporation equals rate of condensation– There is a balance

or equilibrium point at each of these temperatures

– What happens if the temp changes?

• Types of Equilibrium - Phase Equilibrium–What happens if the temperature changes?–If we write the phase equilibrium as an equation:

H2O (l) H2O (g) • At a temperature of 25oC, there might be 100 mL

of liquid water and 5 mL of water vapor• As the temperature increases, more liquid will

change to gas and the equilibrium point shifts to the right

• At first, evaporation will occur faster than condensation and more water vapor will form; eventually, the rates will equal out as more vapor cools to from liquid

• A new balance point occurs where there 95 mL of liquid water and 10 mL of water vapor

• Types of Equilibrium – Solution Equilibrium – When a solid dissolves in a saturated solution– The rate of dissolving = the rate of crystallization

• Examples:– Mr. Foley’s Duncan Donuts Light and Sweet coffee

– C12H22O11 (s) C6H12O6 (aq)

where sugar is dissolving =

– C6H12O6 (aq) C12H22O11 (s)

where sugar is precipitating and settling on the bottom =

– The equilibrium equation is

C12H22O11(s) C6H12O6(aq)

• Types of Equilibrium – Solution Equilibrium – Remember – a saturated solution holds the

maximum amount of solute at that temperature– What happens if we lower the temp of Mr. Foley’s

coffee?– The rate of dissolving < the rate of crystallization

(at first)– Sugar precipitates faster than

the sugar dissolves– A new saturation point is reached

as the excess sugar at the new temp falls out of solution

– But a new equilibrium occurs with more sugar at the bottom

Solution Equilibrium – solids in liquids• So lowering the temperature of the coffee will shift the

equilibrium or balance point to the left as more sugar precipitates out and settles on the bottom:

C12H22O11(s) C12H22O11 (aq)

• Salts in solution will act the same wayNaCl (s) Na+

(aq) + Cl+(aq)

– An increase in temperature will shift the equilibrium point to the right as more salt dissolves at higher temperatures

– A decrease in temperature will shift the equilibrium point to the left as more salt precipitates out of solution

Solution Equilibrium – Gases in solution• Gases dissolved in Liquids

– In a closed system (closed container)– Equilibrium may exist between dissolved gas

and the gas dissolved in a solution

CO2(g) CO2(aq) – Note: gas equilibrium is affected by both

pressure AND temperature!– Increase the pressure – more CO2 is dissolved– Increase the temperature – less CO2 is dissolved– In each case the equilibrium is shifted either to

left or right in the equation– The rates will be equal, but the amounts may

differ

EquilibriumAim KE 5e – What happens in chemical equilibrium reactions?

Chemical Equilibrium• In any chemical reaction:

– Many chemical reactions go to completion– Examples

Pb(NO3)3(aq) + KI (aq) KNO3(aq) + PbI(solid)

• The reverse reaction cannot happen because PbI2 precipitates out of the solution and is no longer available to react

CaCO3(s) CaO(s) + CO2(g)

• The reverse reaction cannot happen because the CO2 gas is released and cannot react with the CaO to become CaCO3 again

Chemical Equilibrium• In many other chemical reactions the reactions are

reversible–Ex: the Haber Process (production of ammonia)

• In a reaction chamber at equilibrium, the H2 and N2 react to create NH3

3H2(g) + N2(g) 2NH3(g) + 91.8 kJ

• this is an exothermic rxn; heat is lost in the rxn

• NH3 breaks down into H2 and N2 at the same rate

2NH3(g) + 91.8 kJ 3H2(g) + N2(g)

• this is an endothermic rxn; heat is gained in the rxn

Chemical Equilibrium• A balance or equilibrium occurs between the two

reaction rates–The forward rate = the reverse rate

3H2(g) + N2(g) 2NH3(g) + 91.8 kJ –The amounts of H2, N2 and NH3 are constant at a

given temperature and pressure –If stress is applied to this system, the equilibrium of

the system will change as well• LeChatelier’s Principle:

–any system in equilibrium can be disturbed by adding a stress to it

–Stress will cause the system to reach a new equilibrium point

Le Chatelier’s and Chemical Systems• In the Haber process of ammonia synthesis

N2(g) + 3H2(g) 2NH3(g) + 91.8 kJ

• At equilibrium the concentrations of nitrogen, hydrogen, and ammonia are constant

• But a stress added to this system will cause a chemical shift

– Either the stress will favor the forward rxn or the reverse rxn

– This will change the concentrations of the substances involved

– And eventually make the rates of the forward and reverse reactions equal again

Le Chatelier’s Principle and Chemical Systems

1. Effect of Concentration on Equilibrium– In the Haber process

N2(g) + 3H2(g) 2NH3(g) + 91.8 kJ

– The concentration (molarity, M) is expressed in moles per liter

– [N2] = concentration of nitrogen,

– [H2] = concentration of hydrogen

– [NH3] = concentration of ammonia

– When the concentration of one of the above increases,

– the system will shift to reduce that amount and return the system to a new equilibrium

Le Chatelier’s and Chemical Systems

– In the Haber process

N2(g) + 3H2(g) 2NH3(g) + 91.8 kJ

– Increasing the [N2] makes more N2 molecules available to react with H2 molecules

• More NH3 is then produced, increasing [NH3] but also decreasing the [H2]

• A new equilibrium point is then reached

– Increasing the [NH3] makes more [N2] and [H2] as the increased number of NH3 molecules allows more to break down

Le Chatelier’s Principle and Chemical Systems

2 - Effect of a Change in Pressure on Equilibrium– In the Haber process

N2(g) + 3H2(g) 2NH3(g) + 91.8 kJ

– increasing the pressure brings all the molecules closer together and the system tries to relieve the stress

– A shift will occur in the direction that will produce the fewest number of molecules that take up less space

Le Chatelier’s Principle and Chemical Systems

N2(g) + 3H2(g) 2NH3(g) + 91.8 kJ

– Four molecules are on the left, and only two are on the right

– As the pressure increases and NH3 molecules are formed, there are fewer molecules in the system

– Less force against the container side due to less molecules

– Less pressure as the system shifts to the right• In the reaction below there is no effect on the system

by changing the volume… why?

H2(g) + Cl2(g) 2HCl (g)

• Two on the left, two on the right, already equal

Le Chatelier’s and Chemical Systems

3. Effect of Temperature– In the Haber Process

N2(g) + 3H2(g) 2NH3(g) + 91.8 kJ– The forward reaction is exothermic– The reverse reaction is endothermic– Which needs more heat to happen – the

endothermic or exothermic reaction?

• An increase in temperature always favors the endothermic reaction more

– In the Haber Process, therefore, the equilibrium shifts to the left as this adds required heat to the right side of the reaction

Le Chatelier’s and Chemical Systems

4. Effect of Catalysts on equilibrium– In the Haber Process

N2(g) + 3H2(g) 2NH3(g) + 91.8 kJ

– A catalyst lowers the activation energy of both forward and reverse reactions

– Therefore, there is an equal affect on both reactions and their rates

– Catalysts don’t affect equilibrium