Chemical Kinetics Chapter 13 Chemical Kinetics Thermodynamics does

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Chemical Kinetics Chapter 13

Chemical Kinetics Chapter 13

Chemical Kinetics Thermodynamics – does a reaction take place? Kinetics – how fast does

Chemical Kinetics Thermodynamics – does a reaction take place? Kinetics – how fast does a reaction proceed? Reaction rate is the change in the concentration of a reactant or a product with time (mol/(Ls)). A B D[A] rate = Dt D[A] = change in concentration of A over time period Dt D[B] rate = Dt D[B] = change in concentration of B over time period Dt Because [A] decreases with time, D[A] is negative. 13. 1

A B time D[A] rate = Dt D[B] rate = Dt 13. 1

A B time D[A] rate = Dt D[B] rate = Dt 13. 1

Br 2 (aq) + HCOOH (aq) 2 Br- (aq) + 2 H+ (aq) +

Br 2 (aq) + HCOOH (aq) 2 Br- (aq) + 2 H+ (aq) + CO 2 (g) time 393 nm light Detector D[Br 2] a DAbsorption 13. 1

Br 2 (aq) + HCOOH (aq) 2 Br- (aq) + 2 H+ (aq) +

Br 2 (aq) + HCOOH (aq) 2 Br- (aq) + 2 H+ (aq) + CO 2 (g) slope of tangent [Br 2]final – [Br 2]initial D[Br 2] average rate = =Dt tfinal - tinitial instantaneous rate = rate for specific instance in time 13. 1

Reaction Rates and Stoichiometry 2 A B Two moles of A disappear for each

Reaction Rates and Stoichiometry 2 A B Two moles of A disappear for each mole of B that is formed. 1 D[A] rate = 2 Dt a. A + b. B D[B] rate = Dt c. C + d. D 1 D[A] 1 D[C] 1 D[D] 1 D[B] rate = =a Dt b Dt c Dt d Dt 13. 1

Write the rate for the following reaction: CH 4 (g) + 2 O 2

Write the rate for the following reaction: CH 4 (g) + 2 O 2 (g) CO 2 (g) + 2 H 2 O (g) D[CH 4] D[CO 2] 1 D[H 2 O] rate = = == Dt Dt Dt 2 13. 1

A+B Exothermic Reaction + + AB C+D Endothermic Reaction The activation energy (Ea )

A+B Exothermic Reaction + + AB C+D Endothermic Reaction The activation energy (Ea ) is the minimum amount of energy required to initiate a chemical reaction. 13. 4

Kinetic Molecular Theory • Model of what happens to gas particles during experimentation –

Kinetic Molecular Theory • Model of what happens to gas particles during experimentation – Large numbers of molecules in continuous motion – Attractive and repulsive forces are negligible – Energy is transferred between molecules during collisions, but average kinetic energy is unchanged (as long as temp is constant) – Kinetic energy of molecules is proportional to the absolute temperature. (in K)

COLLISION THEORY Collision theory states that. . . • particles must COLLIDE before a

COLLISION THEORY Collision theory states that. . . • particles must COLLIDE before a reaction can take place • not all collisions lead to a reaction • reactants must possess at least a minimum amount of energy - ACTIVATION ENERGY plus • particles must approach each other in a certain relative way - the STERIC EFFECT

COLLISION THEORY Collision theory states that. . . • particles must COLLIDE before a

COLLISION THEORY Collision theory states that. . . • particles must COLLIDE before a reaction can take place • not all collisions lead to a reaction • reactants must possess at least a minimum amount of energy - ACTIVATION ENERGY plus • particles must approach each other in a certain relative way - the STERIC EFFECT According to collision theory, to increase the rate of reaction you therefore need. . . more frequent collisions increase particle speed have more particles present more successful collisions give particles more energy or lower the activation energy or

A Reaction Profile CO(g) + NO 2(g) CO 2(g) + NO(g)

A Reaction Profile CO(g) + NO 2(g) CO 2(g) + NO(g)

Particles must collide with the proper geometry or orientation for atoms to come in

Particles must collide with the proper geometry or orientation for atoms to come in direct contact and form the chemical bonds of the products. (steric factor)

 • If both of these conditions are not met, particles will merely collide

• If both of these conditions are not met, particles will merely collide and bounce off one another without forming products.

http: //www. mhhe. com/physsci/chem istry/essentialchemistry/flash/collis 1 1. swf • Although, the percentage of successful

http: //www. mhhe. com/physsci/chem istry/essentialchemistry/flash/collis 1 1. swf • Although, the percentage of successful collisions is extremely small, chemical reactions still take place at a reasonable rate because there are so many collisions per second between reactant particles.

INCREASING THE RATE The following methods may be used to increase the rate of

INCREASING THE RATE The following methods may be used to increase the rate of a chemical reaction • INCREASE THE SURFACE AREA OF SOLIDS • INCREASE TEMPERATURE • ADD A CATALYST • INCREASE THE PRESSURE OF ANY GASES • INCREASE THE CONCENTRATION OF REACTANTS

INCREASING SURFACE AREA • Increasing surface area increases chances of a collision - more

INCREASING SURFACE AREA • Increasing surface area increases chances of a collision - more particles are exposed • Powdered solids react quicker than larger lumps • Catalysts (e. g. in catalytic converters) are in a finely divided form for this reason + In many organic reactions there are two liquid layers, one aqueous, the other nonaqueous. Shaking the mixture improves the reaction rate as an emulsion is often formed and the area of the boundary layers is increased giving more collisions. 1 1 CUT THE SHAPE INTO SMALLER PIECES 3 1 1 3 SURFACE AREA 9+9+3+3 = 30 sq units SURFACE AREA 9 x (1+1+1+1) = 54 sq units

INCREASING TEMPERATURE Effect increasing the temperature increases the rate of a reaction particles get

INCREASING TEMPERATURE Effect increasing the temperature increases the rate of a reaction particles get more energy so they can overcome the energy barrier particle speeds also increase so collisions are more frequent ENERGY CHANGES DURING A REACTION As a reaction takes place the enthalpy of the system rises to a maximum, then falls A minimum amount of energy is required to overcome the ACTIVATION ENERGY (Ea). Only those reactants with energy equal to, or greater than, this value will react. If more energy is given to the reactants then they are more likely to react. Typical energy profile diagram for an exothermic reaction

NUMBER OF MOLECUES WITH A PARTICULAR ENERGY INCREASING TEMPERATURE MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY

NUMBER OF MOLECUES WITH A PARTICULAR ENERGY INCREASING TEMPERATURE MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY Because of the many collisions taking place between molecules, there is a spread of molecular energies and velocities. This has been demonstrated by experiment. It indicated that. . . no particles have zero energy/velocity some have very low and some have very high energies/velocities most have intermediate velocities.

NUMBER OF MOLECUES WITH A PARTICULAR ENERGY INCREASING TEMPERATURE MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY

NUMBER OF MOLECUES WITH A PARTICULAR ENERGY INCREASING TEMPERATURE MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY T 1 T 2 TEMPERATURE T 2 > T 1 MOLECULAR ENERGY Increasing the temperature alters the distribution • get a shift to higher energies/velocities • curve gets broader and flatter due to the greater spread of values • area under the curve stays constant - it corresponds to the total number of particles

NUMBER OF MOLECUES WITH A PARTICULAR ENERGY INCREASING TEMPERATURE T 3 MAXWELL-BOLTZMANN DISTRIBUTION OF

NUMBER OF MOLECUES WITH A PARTICULAR ENERGY INCREASING TEMPERATURE T 3 MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY T 1 TEMPERATURE T 1 > T 3 MOLECULAR ENERGY Decreasing the temperature alters the distribution • get a shift to lower energies/velocities • curve gets narrower and more pointed due to the smaller spread of values • area under the curve stays constant - it corresponds to the total number of particles

NUMBER OF MOLECUES WITH A PARTICULAR ENERGY INCREASING TEMPERATURE T 3 MAXWELL-BOLTZMANN DISTRIBUTION OF

NUMBER OF MOLECUES WITH A PARTICULAR ENERGY INCREASING TEMPERATURE T 3 MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY T 1 T 2 TEMPERATURE T 2 > T 1 > T 3 MOLECULAR ENERGY REVIEW no particles have zero energy/velocity some particles have very low and some have very high energies/velocities most have intermediate velocities as the temperature increases the curves flatten, broaden and shift to higher energies

NUMBER OF MOLECUES WITH A PARTICULAR ENERGY INCREASING TEMPERATURE MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY

NUMBER OF MOLECUES WITH A PARTICULAR ENERGY INCREASING TEMPERATURE MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY Ea NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER MOLECULAR ENERGY ACTIVATION ENERGY - Ea The Activation Energy is the minimum energy required for a reaction to take place The area under the curve beyond Ea corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react.

INCREASING TEMPERATURE NUMBER OF MOLECUES WITH A PARTICULAR ENERGY TEMPERATURE MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR

INCREASING TEMPERATURE NUMBER OF MOLECUES WITH A PARTICULAR ENERGY TEMPERATURE MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY T 2 > T 1 T 2 Ea EXTRA MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER MOLECULAR ENERGY Explanation increasing the temperature gives more particles an energy greater than E a more reactants are able to overcome the energy barrier and form products a small rise in temperature can lead to a large increase in rate

ADDING A CATALYST • Catalysts provide an alternative reaction pathway with a lower Activation

ADDING A CATALYST • Catalysts provide an alternative reaction pathway with a lower Activation Energy (Ea) • Decreasing the Activation Energy means that more particles will have sufficient energy to overcome the energy barrier and react • Catalysts remain chemically unchanged at the end of the reaction. WITHOUT A CATALYST WITH A CATALYST

ADDING A CATALYST NUMBER OF MOLECUES WITH A PARTICULAR ENERGY MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR

ADDING A CATALYST NUMBER OF MOLECUES WITH A PARTICULAR ENERGY MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER MOLECULAR ENERGY Ea The area under the curve beyond Ea corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react. If a catalyst is added, the Activation Energy is lowered - Ea will move to the left.

ADDING A CATALYST NUMBER OF MOLECUES WITH A PARTICULAR ENERGY MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR

ADDING A CATALYST NUMBER OF MOLECUES WITH A PARTICULAR ENERGY MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY EXTRA MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER MOLECULAR ENERGY Ea The area under the curve beyond Ea corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react. Lowering the Activation Energy, Ea, results in a greater area under the curve after Ea showing that more molecules have energies in excess of the Activation Energy

CATALYSTS - A REVIEW • work by providing an alternative reaction pathway with a

CATALYSTS - A REVIEW • work by providing an alternative reaction pathway with a lower Activation Energy • using catalysts avoids the need to supply extra heat - safer and cheaper • catalysts remain chemically unchanged at the end of the reaction. Types Uses Homogeneous Catalysts same phase as reactants e. g. CFC’s and ozone Heterogeneous Catalysts different phase to reactants e. g. Fe in Haber process used in industry especially where an increase in temperature results in a lower yield due to a shift in equilibrium (Haber and Contact Processes)

INCREASING THE PRESSURE • increasing the pressure forces gas particles closer together • this

INCREASING THE PRESSURE • increasing the pressure forces gas particles closer together • this increases the frequency of collisions so the reaction rate increases • many industrial processes occur at high pressure to increase the rate. . . but it can adversely affect the position of equilibrium and yield The more particles there are in a given volume, the greater the pressure The greater the pressure, the more frequent the collisions The more frequent the collisions, the greater the chance of a reaction

INCREASING CONCENTRATION Increasing concentration = more frequent collisions = increased rate of reaction Low

INCREASING CONCENTRATION Increasing concentration = more frequent collisions = increased rate of reaction Low concentration = fewer collisions Higher concentration = more collisions However, increasing the concentration of some reactants can have a greater effect than increasing others

Concentration • A higher concentration of reactants leads to more effective collisions per unit

Concentration • A higher concentration of reactants leads to more effective collisions per unit time, which leads to an increasing reaction rate • We are not increasing the amount being made for a given balanced equation with limiting reactants, we are only speeding up how quickly those products are made.

Cu. CO 3(s) + 2 HCl(aq) Cu. Cl 2(aq) + CO 2(g) + H

Cu. CO 3(s) + 2 HCl(aq) Cu. Cl 2(aq) + CO 2(g) + H 2 O(l) blue Describe seven different ways to monitor the rate of the above reaction. State how each property would change as the reaction proceeds. 1. Mass of Cu. CO 3(s) over time decreases 2. [HCl] over time decreases 3. [Cu. Cl 2] over time increases 4. Volume of CO 2 over time increases 5. Mass of a open beaker over time decreases 6. Pressure of a closed beaker over time increases 7. Colour of the blue Cu 2+ over time increases

Cu. CO 3(s) + 2 HCl(aq) Cu. Cl 2(aq) + CO 2(g) + H

Cu. CO 3(s) + 2 HCl(aq) Cu. Cl 2(aq) + CO 2(g) + H 2 O(l) blue Describe five different ways to increase the rate of the above reaction. 1. Increase the temperature 2. Increase [HCl] 3. Add a catalyst 4. Increase the surface area of Cu. CO 3(s) 5. Agitate We can't change the nature of the reactant because then we wouldn't have the same reaction. Replacing HCl with H 2 SO 4 would be faster but a different reaction.