Kinetics Workbook

for

Chemistry 12

Period Worksheets Quiz

• 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 Q 4
• 5 Mechanisms WS 5 Q5
• 6 Lab: The Iodine Clock Reaction Internet Review
• 7 Review Practice Test 1

• 8 Review Practice Test 2
• 9. Test

This workbook will allow you to 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 do not quite understand and bring them to class. Your instructor will review them. 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 cannot 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.

WS #1 Monitoring and Calculating Reaction Rates

1. Read the chapter from your textbook on Kinetics 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?

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

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)

iii)

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

Mass of copper (g) 3.26 2.93 2.61

Time (min) 5.0 7.0 9.0

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

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

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

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

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

Volume of NO₂ (mL) 10.0 11.5 12.7

Time (sec) 0.00 5.00 10.00

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

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

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

12. Calculate the rate of the following reaction:

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

_0.080

_0.060

_0.040

_0.020

_{0.00 2.0 4.0 6.0 8.0 10.0 12.0}

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

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

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

d) What is the rate in grams N₂ per hour?

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

i.

ii.

iii.

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

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 (min) 1.0 2.0 3.0

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

2. Ag⁺_(aq) + Cl^-_(aq) → AgCl _(s)

3. H_2(g) + F_2(g) → 2HF_(g)

4. 2Al_(s) + 3I_2(s) → 2AlI_3(s)

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

1.

2.

3.

4.

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

5.

6.

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)

2. HCl_(aq) + NaOH_(aq) → NaCl_(aq) + H₂O _(l)

3. Zn_(s) + 2HCl_(aq) → H_2(g) + ZnCl_2(aq)

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

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

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

2H₂O_2(aq) → 2H₂O_(l) + O_2(g)

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) + 2HCl_(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)

Solid reactants are slower than gases, which are slower than aqueous.

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

a) 2Al_(s) + 3I_2(s) → 2AlI_3(s)

b) Ag⁺_(aq) + Cl^-_(aq) → AgCl_(s)

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

a) 2Al_(s) + 3I_2(s) → 2AlI_{3 (s)}

b) 2Na⁺_(aq) + 2ClO^-_(aq) → 2Na⁺_(aq) + 2Cl^-_(aq) + O_2(g)

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

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

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

Ws # 3 Collision Theory

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

2. A successful collision requires _____________________ and __________________ .

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

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

_______ _______________________________________________________

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

_______ _______________________________________________________

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

_______ _______________________________________________________

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

Zn_(s) + 2 HCl_(aq) → ZnCl_2(aq) + H_2(g)

________________________________________________________________

ii) List four ways to increase the rate.

________________________________________________________________

________________________________________________________________

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)

________________________________________________________________

________________________________________________________________

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

i) Increasing temperature i) _________________

_________________

ii) Increasing [reactants] ii) _________________

iii) Increasing surface area (solid) iii) _________________

iv) Agitation of a heterogeneous reaction iv) _________________

v) Adding a catalyst v) _________________ _________________

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

Before Collision After Collision

A)

B)

________________________________________________________________

________________________________________________________________

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

9.

A candle is not burning at room temperature

A match lights the candle

The candle continues to burn

10.

H₂O₂ decomposes slowly at 20^o C

KI is added and rapid decomposition begins

The temperature increases

11.

H₂ and O₂ in a balloon do not react

A spark ignites the balloon

An explosion results

12.

CH₄ and O₂ in a balloon do not react

A platinum gauze ignites the balloon

An explosion results

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)

ii)

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

i)

ii)

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

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

i)

ii)

iii)

iv)

v)

16. C_(s) + O_2(g) → CO_2(g)

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

i)

ii)

iii)

iv)

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

Rate

Activation Energy

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

Rate

Temperature

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

Rate

Molarity

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

a)

b)

c)

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

a)

b)

c)

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

2 H₂O_2(aq) → 2 H₂O_(l) + O_2(g)

23. Describe how you could 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.

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

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)

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.

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.

Ea

26. 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.

Ea

Kinetics - Descriptions

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

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

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.

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

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

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

6. Water puts out a fire.