Kinetics Worksheets Questions

Power Point Lesson Notes- double click on the lesson number.

Worksheet Answers Quiz Answers

1 Monitoring Reaction Rates WS 1 Q1

2 Factors that Change the Rate WS 2 Q2

3 Collision Theory WS 3 Q3

4 PE Diagrams WS 4 Q4

5 Mechanisms WS5 Q5

6 Lab: The Iodine Clock Reaction Web Review Lab Handout

7 Review Practice Test 1 Practice Test 2

8. Review Quizmebc

9. Test

Text book Hebden Read Unit I

The following workbook is designed to ensure that you can demonstrate your understanding of all aspects of the kinetics unit. The minimum expectation is that you do all of these questions by the due dates given by your teacher. Do the questions. Use your notes from class to assist you. Then after you have finished go to the web site to evaluate your work. Make a list of those things that you don’t quite understand and bring them to class. I will go over them as best as I can. There are other things that you should do to prepare for the test at the end of the unit. Remember, what you put into this course is what you will get out. There is no substitute for consistent effort and hard work. If you can’t do a question, get some help before the end of the unit, you need to know, understand, and remember everything. Good luck! I know you can do well in this unit. Web Site Address: iannonechem.com

Ws #1 Monitoring and Calculating Reaction Rates

1. Read unit 1 of Hebden over the next week. “A” students should read it twice.

2. a) When measuring a property associated with a reactant in a reaction, does it increase or decrease?

Decrease as reactants are converted into products

2. b) When measuring a property associated with a product in a reaction, does it increase or decrease?

Increase as reactants are converted into products

3. Give three ways to measure the rate of the following reaction. State the specific properties that you would monitor and include units (amount is not a specific property). State if each property would increase or decrease. Describe in each case how you would calculate the reaction rate.

2HNO_3(aq) + Cu_(s) → NO_2(g) +H₂O_(l) + CuNO_3(aq)

The first one is done for you.

i) Mass of Cu Grams Decrease Rate = mass/time

ii) [HNO₃] M Decrease Rate = M/time

iii) Volume of NO₂ L Increase Rate = L/time

iv) [CuNO₃] M increase Rate = M/time

v) Mass of open container Grams Decrease Rate = mass/time

vi) Pressure of closed container KPa Increase Rate = Pressure/time

Any three of the above answers are fine.

Mass of copper (g) 3.26 2.93 2.61

Time (min) 5.0 7.0 9.0

4. Calculate the rate in units of (g Cu/min).

3.26 - 2.61 g Cu = 0.16g/min

9.0 - 5.0 min

5. Calculate the rate in units of (mole Cu/min).

0. 16g Cu x 1 mol = 0.0026 mol/min

min 63.5 g

6. Calculate the rate in moles HNO₃ consumed per second (remember that 2 moles are consumed per 1 mole of Cu).

0.0026 mol Cu x 1 min x 2 moles HNO₃ = 8.5 x 10^-5 moles/s

min 60s 1 mole Cu

7. Calculate the rate in units of (g/sec) for HNO_3.

8.5 x 10^-5 moles HNO₃ x 63.0 g = 0.0054 g/s

s 1 mole

Volume of NO₂ (mL) 10.0 11.5 12.7

Time (sec) 0.00 5.00 10.00

8. Calculate the rate in units of (mL NO₂/sec).

Rate = 12.7 - 10.0 ml = 0.27 ml/s

10.00 - 0.00 s

9. Calculate the rate in units of (L NO₂/min).

0.27 ml x 60s x 1L = 0.016 L/min

s 1 min 1000 ml

10. Calculate the rate in units of (moles NO₂/min) at STP.

0.0162 L x 1 mol = 7.2 x 10^-4 mol/min

min 22.4 L

11. Calculate the rate in units of (moles HNO₃/min) at STP (remember that 2 moles are consumed per 1 mole of NO₂)

7.23 x 10^-4 mol NO₂ x 2 moles HNO₃ = 0.0015 moles/min

min 1 mole NO₂

12. Calculate the rate of the following reaction:

2NO _(g) + 2H_{2 (g)} → N_{2 (g)} + 2H₂O _(g)

Rate = (0.080 – 0.020) moles = 0.0060 moles/s

(12.0 – 2.0) s

a) What is the rate in moles NO per second? 0.0060 moles/s

b) What is the rate in moles N₂ per second? 0.0030 moles/s

c) What is the rate in grams NO per min? 11 g/min

d) What is the rate in grams N₂ per hour? 3.0 x 10²g/h

13. Choose three properties that you could measure in order to monitor the rate of the following reaction.

Cu _(s) + 2AgNO_{3 (aq)} → 2 Ag _(s) + Cu(NO₃)_{2 (aq)}

Property Unit of Measurement Change

1. Mass Cu grams decrease

2. Mass Ag grams increase

3. Intensity [Cu⁺²] M increase

14. Calculate the rate of the following reaction in units of M/s:

Zn (s) + 2HCL (aq) → ZnCl_{2 (aq)} + H_{2 (g)}

Molarity of HCL (M) 0.612 0.813 1.05

time (seconds) 21.0 25.0 29.0

Rate = (1.05 - 0.612) M = 0.055 M/s

(29.0 - 21.0) s

15. Calculate the rate of the following reaction in L/min:

Zn (s) + 2HCL (aq) → ZnCl_{2 (aq)} + H_{2 (g)}

Volume of H₂ (L) 0.255 0.550 0.790

time (minutes) 1.0 2.0 3.0

Rate = (0.790 - 0.255) M = 0.27 L/min

(3.0 - 1.0) s

16. If 0.369g of HCl is neutralized with 0.250M NaOH in 25.0 seconds, what is the reaction rate in moles HCL /min.

0.369g x 1 mole

Rate = 36.5g = 0.0243 mole/min

0.41666 min

WS # 2 Factors That Change The Reaction Rate

Homogeneous reactions

- reactants are in the same phase (aq), (g) , or (l) and are thoroughly mixed.

Heterogeneous reactions

- reactants are in the two or more phases and are not thoroughly mixed (two solids do not mix).

Classify as Homogeneous or Heterogeneous:

1. Zn _(s) + 2 HCl _(aq) → H_{2 (g)} + ZnCl_{2 (aq)} heterogeneous

2. Ag⁺ _(aq) + Cl^- _(aq) → AgCl _(s) homogeneous

3. H_{2 (g)} + F_{2 (g)} → 2HF _(g) homogeneous

4. 2Al _(s) + 3I_{2 (s)} → 2AlI_{3 (s)} heterogeneous

The following four factors will increase the rate of a chemical reaction that is homogeneous:

1. Increasing the temperature.

2. Increasing the reactant concentration.

3. Adding a catalyst

4. Changing the nature of the reaction.

5. Increasing the pressure for gases

The above four factors as well as the two below will increase the rate of a heterogeneous reaction:

6. Increasing the surface area of a solid.

7. Agitation

Which factor will only increase the rate of a gaseous reaction?

8. Pressure

For each reaction specifically describe all of the ways to increase the reaction rate

(i.e.. increase[H₂]).

1. H_{2 (g)} + F_{2 (g)} → 2 HF _(g) This reaction is homogeneous so the first four factors will work.

Increasing the temperature

Increasing the pressure

Increasing [H₂] or [F₂]

Adding a catalyst

2. HCl _(aq) + NaOH _(aq) → NaCl _(aq) + H₂O _(l) This reaction is homogeneous so the first four factors will work.

Increasing the temperature

Increasing [HCl] or [NaOH]

Adding a catalyst

3. Zn _(s) + 2 HCl _(aq) → H_{2 (g)} + ZnCl_{2 (aq)} This reaction is heterogeneous most of the factors will work, except pressure- need a gaseous reactant..

Increasing the temperature

Increasing [HCl]

Adding a catalyst

Increasing the surface area of Zn_(s)

Agitation

4. State three examples of chemical reactions that are desired to be slow.

Food spoiling

Metal corrosion

Erosion

5. Give three examples of chemical reactions that are desired to be fast.

Combustion of gasoline in automobiles

Industrial chemical production

Cooking food

The combustion of gasoline in a car engine; while accelerating.

6. List all of the ways to increase the rate of the following reaction:

H₂O_{2 (aq)} → H_{2 (g)} + O_{2 (g)}

Increasing the temperature

Increasing [H₂O₂]

Adding a catalyst

I. Homogeneous reactions are generally faster than heterogeneous- the reactants are mixed better and therefore there are more collisions between reactant particles.

HCl _(aq) + NaOH _(aq) → NaCl _(aq) + H₂O _(l)

is faster than

Zn _(s) + 2 HCl _(aq) → H_{2 (g)} + ZnCl_{2 (aq)}

II. Simple ionic reactions (where there are no bonds to break) are generally faster than more complex ionic reactions (where there are bonds to break).

Pb⁺² _(aq) + 2Cl^- _(aq) → PbCl_{2 (l)}

is faster than

2Na⁺ _(aq) + 2ClO^- _(aq) → 2Na⁺ _(aq) + 2Cl^-_(aq) + O_{2 (g)}

1. Indicate the faster and slower reaction and explain why.

a) 2Al _(s) + 3I_{2 (s)} → 2AlI_{3 (s)}

Heterogeneous reaction with bonds to break will be slow.

b) Ag⁺_(aq) + Cl^-_(aq) → AgCl _(s)

Homogeneous reaction with no bonds to break will be fast.

2. Indicate the faster and slower reaction and explain why.

a) 2Al _(s) + 3I_{2 (s)} → 2AlI_{3 (s)}

Slow. The reaction is heterogeneous (two solid do not mix) with bonds to break.

b) 2Na⁺ _(aq) + 2ClO^- _(aq) → 2Na⁺ _(aq) + 2Cl^-_(aq) + O_{2 (g)}

Faster. The reaction is homogeneous.

3. Indicate the faster and slower reaction and explain why.

a) 3Ba⁺²_(aq) + 2PO₄^-3 _(aq) → Ba₃(PO₄)_2(aq)

Faster. The reaction is homogeneous and simple ionic with no bonds to break.

b) Cu_(s) + 2Ag⁺_(aq) → Cu⁺² _(aq) + 2Ag _(s)

Slow. The reaction is heterogeneous and the Cu_(s) bonds need to be broken.

Ws # 3 Collision Theory

1. Chemical reactions are the result of collisions between reactant particles, where bonds are broken and new ones form.

2. A successful collision requires sufficient energy and favorable geometry.

3. Describe as fast, medium or slow. Explain!

i) 2 H_{2 (g)} + O_{2 (g)} → 2 H₂0 _(l) (room temperature)

Slow. Gas reactions are slower than aqueous.

ii) 2 Ag⁺ _(aq) + CO₃^2- _(aq) → Ag₂CO_{3 (s)}

Fast. Homogeneous reaction simple ionic- there are no bonds to break

iii) 2 HCl _(aq) + Na₂CO_{3 (aq)} → CO_{2 (g)} + 2 NaCl _(aq) + H₂0 _(l)

Medium. Homogeneous complex reaction - there are bonds to break.

4. i) Describe how you would measure the rate of the reaction :

Zn _(s) + 2 HCl _(aq) → ZnCl_{2 (aq)} + H_{2 (g)}

Measure the decrease in Zn mass.

Measure the increase in H₂ gas volume.

Measure the mass of an open container which decreasing due H₂ escaping.

ii) List four ways to increase the rate.

Increasing the temperature

Increasing [HCl]

Adding a catalyst

Increasing the surface area of Zn_(s)

Agitation

5. A 10 °C temperature increase frequently doubles the rate of a slow reaction because:

a) The temperature has doubled.

b) The PE of the colliding particle has doubled.

c) The KE of the colliding particle has doubled.

d) The fraction of particles with sufficient KE to react has doubled.

6. Both collisions A and B have the same KE. Which collision is successful and explain why.

Before Collision After Collision

A)

B)

Collision B was successful due to favourable geometry.

7. Use the collision theory to explain how each factor increases the reaction rate.

i) Increasing temperature i) more collisions and harder collisions

ii) Increasing [reactants] ii) more collisions

iii) Increasing surface area (solid) iii) more collisions

iv) Agitation of a heterogeneous reaction iv) more collisions

v) Adding a catalyst v) lower Ea & low energy collisions are successful

8. Explain why collision A was successful while collision B was unsuccessful.

Before Collision After Collision

A)

B)

Collision A was successful because it had sufficient energy. The geometry is the same for both collisions.

Explain each of the following using the collision theory. You need to explain each statement.

9.

· a candle is not burning at room temperature Ea is too high

· a match lights the candle Provides Ea

· the candle continues to burn Exothermic

10.

· H₂O₂ decomposes slowly at 20^o C Ea is too high

· KI is added and rapid decomposition begins Catalyst- lowers Ea

· The temperature increases Exothermic

11.

· H₂ and O₂ in a balloon do not react Ea is too high

· A spark ignites the balloon Provides Ea

· An explosion results Exothermic

12.

· CH₄ and O₂ in a balloon do not react Ea is too high

· A platinum gauze ignites the balloon Catalyst lowers Ea

· An explosion results Exothermic

13. N_2(g) + O_2(g) → 2NO_(g)

Even though there are more than four billion collisions per second between N and O the amount of product after a year is too small to detect. Using the collision theory, give two reasons why this reaction might be slow.

i) Low Temperature

ii) High Ea

14. Give two reasons why some collisions will not result in a chemical reaction.

i) Insufficient energy

ii) Poor geometry

15. Give five reasons that might account for the following reaction having a high rate.

Ca _(s) + 2 HCl _(aq) → CaCl_{2 (aq)} + H_{2 (g)}

i) High surface area of Ca

ii) High concentration of HCl

iii) High temperature

iv) Agitation

v) Nature of the reactant

16. C_(s) + O_2(g) → CO_2(g)

List four ways the rate of the reaction could be increased.

i) Increase temperature

ii) Increase [ O₂ ]

iii) Increase pressure

iv) Increase SA of C

(add catalyst or agitate)

17. State the relationship between Activation energy and the rate of a reaction. Graph the relationship.

Inverse because decreasing the activation energy increases the rate.

Rate

Activation Energy

18. State the relationship between Temperature and the rate of a reaction. Graph the relationship.

Direct, because increasing the temperature increases the rate.

Rate

Temperature

19. State the relationship between Concentration and the rate of a reaction. Graph the relationship.

Direct, because increasing the concentration increases the rate.

Rate

20. Give three examples of reactions that are desired to be slow.

a) food spoiling

b) corrosion of metal

c) the fading of the colour in paint

21. Give three examples of reactions that are desired to be fast.

a) explosions

b) the combustion of gasoline in your car when you are passing someone on the freeway

c) the commercial production of chemicals

Molarity

22. List all of the ways to increase the rate of the reaction:

2 H₂O_{2 (aq)} → 2 H₂O _(l) + O_{2 (g)}

Increase the H₂O₂ concentration

Increase the Temperature

Add KI catalyst

23. Describe how you would measure the rate of the reaction above. State the property you would measure and describe how it changes. Draw a diagram to illustrate your answer.

Mass of an open container Decreases

or

Volume of O₂ See notes for diagram.

Pressure of O₂ in a closed system. See notes for diagram.

24. Pick the fastest and the slowest reaction at 20 °C.

Slowest gases are slower than aqueous a) H_2(g) + I_2(g) → 2 HI_(g)

b) 2 HCl_(aq) + Na₂CO_3(aq) → CO_2(g) + 2 NaCl_(aq) + H₂O_(l)

Fastest- simple ionic or double replacement c) Hg²⁺_(aq) + 2 I ^-_(aq) → HgI_2(s)

25. H₂ and O₂ can exist at 20 °C for years without reacting. But when a small spark ignites the mixture it reacts explosively. Explain using the Collision Theory.

The activation energy is too high at room temperature so there are no successful collisions.

A spark provides the kinetic energy required to overcome the Ea.

Exothermic reactions produce energy.

26. Draw a collision energy distribution diagram for a reaction where the y-axis is fraction of collisions and the x-axis is collision energy. Draw the Ea line showing about 10% of the collisions having sufficient energy. Draw the Ea line for the catalyzed reaction where 20% have sufficient energy.

27. Shade in the area of the collision energy distribution diagram showing those collisions that do not have the required energy to be successful at the temperature below.

28. Shade in the area of the collision energy distribution diagram showing those collisions that do have the required energy to be successful at the temperature below. Redraw the curve at a higher temperature.

Collision Energy

Kinetics - Descriptions

Use the collision theory to explain the following. Each sentence must be explained with a statement from the collision theory.

1. An unlit candle does not burn. It burns after being lit with a match. It continues to burn.

Ea is too high.

Match is energy and provides Ea.

Exothermic

2. A solution is reacting very slowly to produce bubbles. KI is added and although it is not consumed in the reaction , it speeds up the reaction rate. The temperature increases. The rate increases even more.

Ea is high

KI is a catalyst and lowers Ea and more collision are successful.

Exothermic→ temperature increases → rate increases

3. Iron reacts slowly with HCl. Iron is replaced with Zn and a much more vigorous reaction rate occurs.

Nature of reactant Fe → high Ea Zn → low Ea

4. H₂ and O₂ can exist together for years at room temperature without reacting. A spark begins the reaction. An explosion results.

High Ea → collisions are not successful

A spark provides the Ea

Exothermic → explosion

5. Dilute nitric acid shows little reaction with copper. Concentrated nitric acid vigorously reacts.

Low concentration → few collisions

High concentration → many collisions

6. Water puts out a fire.

Lowers temperature so there are less collisions

The collisions have less energy.

7. Paint prevents rusting.

There are fewer collisions between reactant molecules.

8. A preservative in food slows rotting.

The preservative is an inhibitor; which increases the Ea.

Ws # 4 Potential Energy Diagrams Worksheet

1. Draw the PE diagram showing the PE changes that occur during a successful collision of the exothermic reaction:

H₂ + I₂ → 2 HI + 250 KJ

The PE of the reactants = 400 KJ

The activation energy of the forward reaction = 200 KJ

Reaction path

2. Draw the PE diagram showing the PE changes that occur during a successful collision of the endothermic reaction:

A + B + 200 KJ → C

The PE of the reactants = 200 KJ

The Activation Energy in the forward direction = 250 KJ

Reaction Path

3. Write the following reaction in ΔH notation.

A + B + 200 kJ → C

A + B -----> C ΔH= +200kJ

4. Write the following reaction in Standard Notation.

H₂ + I₂ → 2 HI ΔH = -250 kJ

H₂ + I₂ → 2HI + 250 kJ

5. Write in Standard Notation.

2NI₃ + 3BaCl₂ → 2NCl₃ + 3BaI₂ ΔH = 175 kJ

2NI₃ + 3BaCl₂ + 175 kJ → 2NCl₃ + 3BaI₂

6. Write in ΔH notation.

2AlBr₃ + 3BaF₂ → 2AlF₃ + 3BaBr₂ + 276 kJ

2AlBr₃ + 3BaF₂ → 2AlF₃ + 3BaBr₂ ΔH= -267 kJ

Draw the potential energy diagram for the following reactions.

7. Potential energy of reactants = 250 kJ

Potential Energy of activated complex = 350 kJ

Potential Energy of the products = 300 kJ

a) How does the potential energy change as the reaction proceeds? Increases

b) How does the kinetic energy change as the reaction proceeds? Decreases

c) Is the reaction exothermic or endothermic? Endothermic

d) What is the value of ΔH? ΔH= +50kJ

If a catalyst was added, what would happen to the energies of the:

e) Reactants? Nothing

f) Products? Nothing

g) Activated Complex? Decrease

h) If a catalyst was added what would happen to the rate? Increase

Draw the potential energy diagram for the following reactions.

8. Potential energy of reactants = 350 kJ

Activation Energy = 100 kJ

Potential Energy of the products = 250 kJ

a) How does the potential energy change as the reaction proceeds? Decreases

b) How does the kinetic energy change as the reaction proceeds? Increases

c) Is the reaction exothermic or endothermic? Exothermic

d) What is the value of ΔH? ΔH= -100kJ

If the concentration of the reactants was increased, what would happen to the energies of the:

e) Reactants? Nothing

f) Products? Nothing

g) Activated Complex? Nothing

h) What would happen to the rate? Increase

Draw the potential energy diagram for the following reactions.

9. Potential energy of reactants = 200 kJ

Potential Energy of activated complex = 400 kJ

ΔH = 150 kJ

a) How does the potential energy change as the reaction proceeds? Increases

b) How does the kinetic energy change as the reaction proceeds? Decreases

c) Is the reaction exothermic or endothermic? Endothermic

d) What is the value of ΔH? ΔH= 150 kJ

If the temperature was increased, what would happen to the energies of the:

e) Reactants? Nothing

f) Products? Nothing

g) Activated Complex? Nothing

h) What would happen to the rate? Increase

10. Potential energy of products = 50 kJ

Potential Energy of activated complex = 400 kJ

ΔH= -50 kJ

Reaction Path

a) How does the potential energy change as the reaction proceeds? Decreases

b) How does the kinetic energy change as the reaction proceeds? Increases

c) Is the reaction exothermic or endothermic? Exothermic

d) What is the value of ΔH? ΔH= -50kJ

If the surface area of the reactants was increased, what would happen to the energies of the:

e) Reactants? Nothing

f) Products? Nothing

g) Activated Complex? Nothing

h) What would happen to the rate? Increase

11. What is the only thing, other than changing the reaction that will change the potential energy diagram? Describe how it will effect the diagram and the rate.

Catalyst Lowers Ea alloys more low energy collisions to be successful and increase the rate.

12. Label each interval on the potential energy diagram. a b c d e

a) Ea (forward) (catalyzed)

PE

b) Ea (reverse)(catalyzed)

c) ΔH

d) Ea (forward) (uncatalyzed)

Reaction Path

e) Ea (reverse) (uncatalyzed)

12. Label each interval on the potential energy diagram.

a b c d e

a) Ea (forward) (uncatalyzed)

b) Ea (forward) (catalyzed)

c) ΔH

d) Ea (reverse) (uncatalyzed)

e) Ea (reverse) (catalyzed)

Ws # 5 Mechanisms

1. OCl^- + H₂O → HOCl + OH^-

HOCl + I^- → HOI + Cl^-

HOI + OH^- → H₂O + OI^-

i) The net chemical equation is: OCl^- + I^- + → Cl^- +OI^-

ii) The reaction intermediates are: HOCl HOI OH^-

iii) The catalyst is: H₂O

2. Br₂ → 2Br fast

Br + OCl₂ → BrOCl + Cl slow

Br + Cl → BrCl fast

i) The net chemical equation is: Br₂ + OCl₂ → BrOCl + BrCl

ii) The reaction intermediates are: Cl & Br

iii) The catalyst is: None

iv) The rate determining step is 2

v) If the concentration of Br₂ is increased will the rate of the reaction increase? Explain your answer.

No because it is not in the rate determining step.

vi) If the concentration of OCl₂ is increased will the rate of the reaction increase? Explain your answer.

Yes because, OCl₂ is in the rate determining step.

3. The mechanism for the catalytic decomposition of formic acid is shown below.

step 1 HCOOH + H⁺ → [HCOOHH]⁺

step 2 [HCOOHH]⁺ → [HCO]⁺ + HOH

step 3 [HCO]⁺ → CO + H⁺

The potential energy diagram is:

190

180

170

160

150

Reaction Path

i) The catalyst is H⁺ Crosses out from left to right

ii) The rate determining step is Two Highest Ea

iii) ΔH = +10 kJ From start to end

iiv) The forward activation energy is 40 kJ Reactants to the highest point

iv) The reverse activation energy is 30 kJ Products to the highest point

v) The enthalpy of [HCOOHH]⁺ is 160 kJ After one hump

vi) Is the reaction exothermic or endothermic? Endo Uphill

vii) Which chemical formula has the greatest potential energy? (HCO)⁺ + HOH Highest point on graph

viii) Which chemical formula has the greatest kinetic energy? HCOOH + H⁺ Lowest point on graph

ix) Does this reaction absorb or release kinetic energy? Absorb because it is endothermic (uphill)

4. Define and remember the following definitions.

mechanism A sequence of steps that determines the overall reaction.

activation energy The minimum energy required in a successful collision.

rate determining step The slowest step in a reaction mechanism.

catalyst A substance that increases the rate of a chemical reaction by providing a alternate mechanism with lower activation energy. reaction intermediate A chemical species produced in a reaction mechanism and then consumed in a later step.

endothermic A reaction that absorbs energy

exothermic A reaction that produces energy

activated complex A unstable reaction intermediate with high potential energy and low kinetic energy.

ΔH The change in enthalpy or heat content for a reaction.

reaction rate The change in a reactant or product per unit of time.

5. The catalyzed decomposition of acetaldehyde has an overall reaction of:

CH₃CHO → CH₄ + CO . Determine step 2 of the reaction mechanism.

A proposed mechanism is:

step 1 CH₃CHO + I₂ → CH₃I + HI + CO

step 2 HI + CH₃I → I₂ + CH₄ This is the only step 2 that will give the overall reaction below.

overall CH₃CHO → CH₄ + CO

6. The following reaction has an overall reaction of:

2Ce⁴⁺ + Tl⁺ → 2Ce³⁺ + Tl³⁺

Determine step 2 of the reaction mechanism.

A proposed mechanism is:

step 1 Ce⁴⁺ + Mn⁺² → Ce³⁺ + Mn³⁺

step 2 Ce⁴⁺ + Mn³⁺ → Ce³⁺ + Mn⁴⁺ This is the only step 2 that will give the overall reaction below

step 3 Mn⁴⁺ + Tl⁺ → Tl³⁺ + Mn²⁺

overall 2Ce⁴⁺ + Tl⁺ → 2Ce³⁺ + Tl³⁺

7. A reaction has a overall equation of: Br₂ + OCl₂ → BrOCl + BrCl . Determine step 3 of the mechanism.

step 1 Br₂ → 2Br

step 2 Br + OCl₂ → BrOCl + Cl

step 3 Br + Cl → BrCl This is the only step 3 that will give the overall reaction below

overall Br₂ + OCl₂ → BrOCl + BrCl

List two intermediates: Br Cl

8. Complete the following mechanism.

step 1 NO + Pt → NOPt needed for next step

step 2 NOPt + NO → O₂Pt + N₂ O₂Pt needed for next step and N₂ needed to be a product

step 3 O₂Pt → O₂ + Pt

overall 2NO → N₂ + O₂

Identify the catalyst Pt Crosses out from left to right

Identify the two intermediates NOPt O₂Pt Crosses out from right to left

9. Draw a collision energy distribution diagram for a reaction where the y-axis is fraction of collisions and the x axis is collision energy. Draw the Ea line showing about 10% of the collisions having sufficient energy. Draw the Ea line for the catalyzed reaction where 20% have sufficient energy.

10. Shade in the area of the collision energy distribution diagram showing those collisions that have the required energy to be successful at the low temperature shown below. Draw the curve that represents the distribution at a higher temperature with a different color. Shade in the area representing the successful collisions at the higher temperature with a new color.