Download Topic 6 Kinetics File

Topic 6
Kinetics
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Definitions
Average bond enthalpy: The average enthalpy change of breaking one mole of a bond in a
gaseous atom into its constituent gaseous atoms.
Born-Haber cycle: Energy cycles for the formation of ionic compounds. If there is little
agreement between the theoretical and experimental values, this could indicate a degree
of covalent character.
Electron affinity: Enthalpy change when an electron is added to an isolated atom in the
gaseous state.
Endothermic: A reaction in which energy is absorbed. ΔH is +. Reactants more stable than
products.
Enthalpy: The internal energy stored in the reactants. Only changes in enthalpy can be
measured.
Entropy: A measure of the disorder of a system. Things causing entropy to increase: 1)
increase of number of moles of gaseous molecules; 2) change of state from solid to liquid
or liquid to gas; 3) increase of temperature
Exothermic: A reaction in which energy is evolved. ΔH is –. Products more stable than
reactants.
Gibb’s free energy: Must be negative for reaction to be spontaneous. ΔG = ΔH – TΔS
Hess’ law: Enthalpy change for a reaction depends only on difference between enthalpy of
products and enthalpy of reactants. It is independent of pathway.
Lattice enthalpy: The endothermic process of converting a crystalline solid into its
gaseous ions, or the reverse exothermic process. The lattice enthalpy increases with
decreasing size of the ions and increasing charge.
Spontaneous: A reaction that has a natural tendency to occur.
Standard conditions: 298 K and 1 atm.
Temperature: A measure of the average kinetic energy.
Standard enthalpy of vaporisation: The energy required to vaporise one mole of a liquid.
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Enthalpy of atomisation: The energy required to produce one mole of gaseous atoms from
an element in its standard state.
Bond dissociation enthalpy: The energy change when one mole of a specific bond is broken
or created under standard conditions.
Enthalpy of Combustion: The energy released when one mole of a compound is burned in
excess oxygen.
Standard enthalpy of formation: The energy change when one mole of a compound is
formed under standard conditions from its constituent elements in their standard states.
Standard enthalpy of solution: The energy change when one mole of a substance is
dissolved in an infinite amount of water under standard conditions.
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Practice Questions for Reaction Rates-1
Part 1
Please indicate the correct answer by circling the letter corresponding to the correct answer
Rates of reaction
1. The reaction between hydrogen and sulfur below is exothermic.
H2(g) + S(s)  H2S(g)
Which one of the following factors will increase the rate of this reaction?
A. Decreasing the temperature
B Decreasing the volume of the container
C Adding more sulfur
D Removing H2S
2. Raising the temperature speeds up most chemical reaction. This is due to an increase in which of
the following?
I. The number of collisions occurring in a given time
II. The energy of each collision
A. I only
B. II only
C. Both I and II
D. Neither I nor II
3. From question 2. above, which factor has a greater effect in speeding up a chemical reaction? i
A. I
B. II
C. I and II have equal effect
D. Depends on the reaction
4. Reaction rates generally increase in response to a decrease in
A. catalyst concentration.
B. reagent concentration
C. particle size.
D. temperature.
5. When 100 cm3 of 1.0 mol dm-3 HCl are added to 1 g of granular zinc at 25°C, hydrogen is evolved.
All of the following will increase the initial rate of hydrogen evolution EXCEPT
A. substituting 2.0 mol dm-3 HCI for 1.0 mol dm-3 HCl.
B. using 200 cm3 of 1.0 mol dm-3 HCI instead of 100 cm3.
C. substituting powdered zinc for granular zinc.
D. increasing the temperature of the 1.0 mol dm-3 HCl to 50°C.
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6. It has been found that the rates of many reactions are doubled by a 10-degree C rise in
temperature. The main reason for this is that
A. the energy of activation decreases with temperature.
B. the energy of activation increases with temperature.
C. the speed of molecules is dramatically increased with a 10 ºC increase in temperature.
D. the fraction of high-energy molecules increases exponentially with temperature.
7. A catalyst will
A. alter the mechanism (pathway) of a reaction.
B. increase ΔH for the reaction.
C. decrease ΔH for the reaction.
D. decrease Ea for the forward reaction only.
8. Which one of the following factors does not affect the rate of a chemical reaction?
A. The amount of reactants.
B. The concentration of reactants.
C. The temperature of the reagents.
D. The presence of a catalyst.
9. In most chemical reactions, the rate of reaction decreases as the reaction proceeds. The usual
reasons cited for this is that
A. The energy for the reaction is sunning out.
B. The concentration of the reactants are becoming lower.
C. The temperature is falling as the reaction proceeds.
D. The activation energy becomes greater.
10. In which of the following situations would you expect the rate of reaction between marble
(calcium carbonate) and nitric acid to be the greatest?
A. Powdered marble and 2 mol dm-3 acid at 40 ºC.
B. Powdered marble and 0.5 mol dm-3 acid at 40 ºC.
C. Powdered marble and 2 mol dm-3 acid at 20 ºC.
D. Marble chips and 0.5 mol dm-3 acid at 40 ºC.
11. In which one of the following reactions would surface area not be a factor affecting the rate?
A. Zinc and sulfuric acid.
B. Carbon dioxide gas with limewater.
C. Vegetable oil and aqueous sodium hydroxide.
D. Aqueous oxalic acid and aqueous potassium permanganate.
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Short Answer
1. A student wishes to produce hydrogen in the laboratory by the reaction of magnesium with
hydrochloric acid.
Mg (s) + HCl (aq)  MgCl2 (aq) + H2 (g)
Discuss the factors that will affect the rate of production of hydrogen. Use the collision
theory of reaction rates to account for any changes in the rate.
2. The Bunsen burners in our lab uses propane as a fuel to produce heat energy according to the
following equation:
C3H8 (g) + 5 O2  3 CO2 + 4 H2O
Explain, using collision theory, how the air in-take valve affects the rate of this reaction.
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Reaction Diagrams and Reaction Rates
The potential energy of substances involved in a reaction can be plotted versus the progress of the reaction, as
the process moves from initial reactants, through activated complex, to final products.
On the grids below, plot energy diagrams for 1–3 below, given the following information, and answer the
questions. For number 4, study the energy diagram and answer the questions.
1. On the grid to the right, plot the energy diagram given
the following information.
Potential Energy of reactants: 250
Potential Energy of activated complex: 350
Potential Energy of products: 300
Is the reaction exothermic or endothermic?
_______________________________________
How can you tell?
What is the value of ΔH ? ______________
If a catalyst were added, what would happen to the diagram?
If a catalyst were added, what would happen to the energies of reactants, products, and activated complex,
and to the rate?
Explain the effect on rate.
2. On the grid to the right, plot the energy diagram given the following information.
Potential Energy of reactants: 350
Activation energy (energy needed to
form activated complex from reactants): 100
Potential Energy of products: 250
Is the reaction exothermic or endothermic?
_______________________________________
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How can you tell?
What is the value of ΔH ? ______________
What is the potential energy of the activated complex? ______________
If the concentration of the reactants were increased, what would happen to the diagram?
If the concentration of the reactants were increased, what would happen to the energies of the reactants,
products, and activated complex, and to the rate?
Explain the effect on the rate
3. On the grid to the right, plot the energy diagram given the following information.
Potential Energy of reactants: 200
Activation energy: 200
ΔH : +150
Is the reaction exothermic or endothermic?
_______________________________________
How can you tell?
What is the potential energy of the products ______________
What is the potential energy of the activated complex? ______________
If the temperature were increased, what would happen to the diagram?
If the temperature were increased, what would happen to the energies of the reactants, products, and
activated complex, and to the rate?
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Explain the effect on rate.
4. Potential Energy of reactants: ____________
Potential Energy of activated complex ____________
Activation energy: ___________
Potential Energy of products: ____________
ΔH : __________
Is the reaction exothermic or endothermic?
_______________________________________
How can you tell?
Factors Affecting Rate
The rates of chemical reactions depend upon a number of factors. These factors can be controlled by chemists
in order to cause processes to proceed at desired rates. For each of the following factors, write its probable
effect (increase, decreases,no effect) on rate,and then explain the effect on the basis of collision theory.
Factor
decreased concentration
Effect on Rate
_______________
increased gas pressure
_______________
decreased pressure
_______________
decreased surface area
_______________
addition of catalyst
_______________
addition of inhibitor
_______________
Explanation
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Practice Questions for Reaction Rates–2
1) Write an expression for the rate of reaction in terms of a changing concentration for each of the
following reactions:
a) C4H8 (g)  2C2H4(g)
b) H2(g) + I2(g)  2HI(g)
c) 2NO2(g) + O3(g)  N2O5(g) + O2(g)
2) Nitrogen monoxide reacts with chlorine to form nitrosyl
chloride.
NO(g) + 1/2Cl2(g)  NOCl(g)
The figure shows the increase in nitrosyl chloride
concentration under appropriate conditions.
(The concentration of nitrosyl chloride starts at 0
although this fact is difficult to see in the graph).
a) Write an expression for the rate of reaction in
terms of a changing concentration.
b) Calculate the average rate of reaction between 40 and 120 sec.
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c) Calculate the instantaneous rate of reaction after 80 sec.
3) The equation for a reaction is: 4NO2(g) + 2 H2O(g) + O2(g)  4 HNO3(g)
Which one of the following is not numerically equal to the others? (Explain why).
4) Which one of the curves on the following graph shows the greatest initial reaction rate? (Explain
why).
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5) Iodate(V) ions oxidize iodide ions in acidic solution to form iodine and water according to the
equation:
IO31-(aq) + 5I1-(aq) + 6H1+(aq)  3 I2(aq) + 3 H2O(l)
If the number of moles of each reactant consumed after one minute was measured, which would have
been consumed least? (Explain why).
a) IO31c) H1+
1b) I
d) They would all have been consumed to the same extent.
6) The rate of reaction between zinc and sulfuric acid is measured by weighing a zinc plate, which is
then placed into a beaker of the acid. Every 10 minutes it is removed, rinsed, dried and reweighed.
This is continued until all of the acid is consumed.
a) Sketch the graph you would expect for the mass of the zinc plate against time in the acid.
b) At what point is the reaction rate the greatest? How can you tell?
c) Suggest another way that the rate of this reaction could have been measured.
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Practice Questions for Reaction Rates–3
Introduction
An enthusiastic chemistry student thought it would be interesting to see how particle
size could affect the rate of a chemical reaction. She took equal amounts of marble
chips (5.00 g) and reacted them in 50.0 mL of 2.0 M HCl according to the following
reaction.
CaCO (s) + 2 HCl(aq) CaCl2 (aq) + CO2(g) + H2O(l)
In one case she used the marble chips as rather large chunks whereas in another case
she used a mortar and pestle to finely grind the chips. She monitored the reaction at
25 °C by measuring the mass of reaction mixture and Erlenmeyer every 20 s for a period
of about 5 minutes. The following table shows that data that she collected.
Use the above reaction conditions and experimental data to do the following data manipulation and to
answer the following questions.
1. Prepare a properly labeled table showing the loss of mass for the both the large and
small marble chips. (Watch for significant figures) [5 points]
2. Now make a plot of loss of mass (y-axis) and time (x-axis) on the same graph for
both types of marble chips. Be sure to carefully label your graph and give it a title.
[10 points]
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3. The rate of reaction is monitored by measuring loss of mass from the reaction flask.
This loss of mass can be equated to the change in mass of which one of the
reactants or products? Explain your answer. [5 points]
4. Which reaction (large or small marble chips) has the greater initial rate? Explain,
using collision theory, why this happened. [5 points]
5. What happens to the rate over time for each reaction? Explain why this happens. [5
points]
6. Is there a difference between the two reaction conditions in the amount of carbon
dioxide formed by the end of the reaction (within experimental error)? Explain why
or why not? Calculate the % yield for both reaction conditions. [10 points]
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7. The student decides to test other factors that affect rate of reaction. She
repeats the experiment using the same quantity of uncrushed CaCO3 but in separate
experiments makes the following changes:
i. Changes 50.0 mL of 2.0 M HCl to 50.0 of 4.0 M HCl
ii. Changes 50.0 mL of 2.0 M HCl to 100.0 of 2.0 M HCl
iii. Keeps the reactant quantities the same but increases the temperature to 50°C.
Reproduce your graph for the loss of mass vs time (50.0 mL of 2.0 M HCl,
uncrushed CaCO3 treatment only). Now, for each of the above changes, indicate
the sort of graph line you would expect to see. Note that this question is meant to
be more qualitative than quantitative – that is - if you double the concentration of
the reactants, will the new rate be faster, the same, or slower than the rate for
uncrushed CaCO3 treatment only. Show all your answers on the same graph. [10
points]
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Reaction Rates
Learning Goal:
To be able to determine how temperature, concentration and activation energy
change the rate of a reaction.
Procedure:

Open the internet browser and enter the address: http://phet.colorado.edu

Click on “Play with Sims” and select “Chemistry” from the menu on the left.

Open the “Reaction & Rates” Simulation and select “Run Now”
Investigation and Analysis:
1. Explore the “Single Collision” Simulation and complete the table below.
Which variables did you
Record observations
Explain this change
change?
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2. You have learned that temperature, concentration and activation energy will change the rate
of a reaction. Complete the table below by comparing your predictions to your observations.
a. Predict what will happen to the rate of the reaction for each variable listed below.
b. Test your prediction with the simulation and record observations.
c. Explain your observations. Were all of your predictions correct? Which tests changed your
thinking about reaction rates? How did the simulation change your prediction?
Variable
Predict
Test
Explain
Increase
temperature
Decrease
temperature
Increase
concentration of
the reactants
Decrease
concentration of
the reactants
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3. Review your observations from questions #1 & 2. Write a summary paragraph, which includes
drawings, which demonstrate you have mastered the learning goal.
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Unit 10: Kinetics REVIEW
1. Which of the following factors should NOT affect the rate of a chemical reaction?
A. The concentration of the reagents
B. The time of day the experiment was done
C. The temperature of the reagents
D. The presence of a catalyst
2. In most chemical reactions, the rate of reaction decreases as the reaction proceeds. The usual
reason for this is that
A. the energy for the reaction is running out
B. the concentration of the reactants are becoming lower
C. the temperature is falling as the reaction proceeds
D. the activation energy becomes greater
3. In which of the following situations would you expect the rate of reaction between marble (calcium
carbonate) and nitric acid to be the greatest?
A.
B.
C.
D.
Powdered marble and 2.0 M acid at 40ºC
Powdered marble and 0.5 M acid at 40ºC
Powdered marble and 2.0 M acid at 20ºC
Marble chips and 0.5 M acid at 40ºC
4. In which one of the following reaction would surface area NOT be a factor affecting the rate?
A.
B.
C.
D.
Zinc and sulphuric acid
Carbon dioxide gas with limewater
Vegetable oil and aqueous sodium hydroxide
Aqueous oxalic acid and aqueous potassium permanganate
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5. The rate of decomposition of an aqueous solution of hydrogen peroxide can be followed by recording
the volume of gas displacing water in a measuring cylinder, against time.
a. Sketch the graph of volume against time you would expect for the complete decomposition of a sample of
hydrogen peroxide.
b. On the same axes, use a dotted line to sketch the curve that you would expect to find if the experiment
were
repeated using a smaller volume of a more concentrated solution of hydrogen peroxide so that the number of
moles of hydrogen peroxide remains constant.
c. When lead dioxide is added, the rate at which oxygen is evolved suddenly increases, even though at the
end of
the reaction, the lead dioxide remains unchanged. Explain this observation.
d. In what other way, could the rate at which the hydrogen peroxide decomposes be increased?
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6. Consider this reaction: 2Al + 3CuSO4  Al2(SO4)3 + 3Cu
ΔH = 350 kJ
a. Describe what must happen in order for this reaction to occur.
b. Now that the reaction is occurring, what are the ways to control the rate of the reaction?
c. Based on the collision theory, explain why the reaction rate changes when you increase the concentration
of the reactants.
c. Write the rate expressions for this reaction.
d. If the reaction rate is 2.5 x 10-3 mol cm-1s-1, determine the rate of formation of the aluminum
sulfate.
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e. Draw the energy diagram and show with a dotted line, the affect of adding a catalyst
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Collision Theory Worksheet
1)
Explain why all reactions have an activation energy, using your knowledge of collision theory.
2)
Describe how the activation energy of a reaction affects the overall rate of the chemical
reaction.
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3)
A rule of thumb used by organic chemists is that the rate of a chemical reaction can be
doubled by increasing the reaction temperature by ten degrees Celsius. Explain this drastic
increase in reaction rate using your knowledge of collision theory.
4)
It has been observed that more gas station fires occur on hot days than on cold days.
Explain this phenomenon using your knowledge of collision theory. (Hint: It’s not just the
temperature increase that causes this!)
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5)
It has been observed with one variety of paint that the rate of paint drying can be
drastically increased by adding a small amount of “accelerant”. Based on what you know of
catalysts, is it reasonable to think of this accelerant as being a catalyst? Explain.
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Kinetics Assignment
1. Explain the following situations in terms of collision theory. Be sure your explanation focuses on
what is happening at the particle level.
a. A reaction between powdered magnesium and 1.00 M hydrochloric acid is much faster than a
reaction between magnesium ribbon and the same amount of hydrochloric acid. (2 pts)
b. The reaction between magnesium ribbon and hydrochloric acid at 5°C is slower than the reaction
where the acid is at 55°C. (3 pts)
c. The reaction between magnesium ribbon and hydrochloric acid is faster with 2.00 M hydrochloric
acid than with 1.00 M HCl. (2 pts)
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2. An experiment is performed in which 1.00 g of magnesium metal is reacted with an excess volume
1.00 mol-1 hydrochloric acid.
a. Write the balanced chemical equation for this reaction. (1 pt)
b. Express the rate of the reaction in terms of both the reactants and products. (1 pt)
c. Suggest 2 ways in which the progress (rate) of this reaction could be measured. (2 pts)
d. Sketch the shape of a graph showing the volume of hydrogen gas given off with time.(2 pts)
e. On the same axes in a different colour, sketch the shape of graph would have if the
concentration of HCl were 2.00 M instead of 1.00 M. Explain the shape of your line. (2 pts)
f. On the same axes in a different colour sketch the shape of the graph if only 41.0 mL of 1.00 M
HCl were used instead of an excess. Explain the shape of your line. (2 pts)
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3. This graph shows how the mass of a reaction mixture decreases with time.
a. Suggest a reason why the mass of the reaction mixture could be decreasing. (1 pt)
b. Sketch how the graph would look if the temperature of the reaction were increased by 10°C
keeping everything else constant. Explain the shape of your graph. ( (2 pts)
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Maxwell-Boltzmann Distribution
The three lines on the graph are
the same gas (both type and
amount) for different
temperatures. I will refer to them
as either solid, (-100), dashed (20)
and dotted (600). The Y-axis is
the number of molecules that have
a given speed.
1. Why is speed equivalent to
temperature?
2. Do all molecules at the same temperature…
a. have the same speed? Y/N
b. have the same average speed? Y/N
3. State three differences in the size/shape of the curve for higher temperatures.
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4. How many molecules have the following speeds:
Speed
Number of gas
Number of gas
Number of gas
particles (SOLID
particles
particles
(DASHED LINE)
(DOTTED LINE)
LINE)
200
400
600
5. Will the area under the solid line = the area under the dashed line? Will either equal the
area under the dotted line?
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6. Draw slanted lines in the area that corresponds to all the molecules that have speeds of 600
m/s or greater for the solid line graph.
7. Using a different kind of shading (slanted lines the other direction, etc.), show all the
molecules that have speeds of 600 m/s or great for the dashed lines and dotted lines.
8. Which temperature has the most molecules with speeds above 600 m/s?
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Catalysts Web Research
Website: http://resources.schoo lscience.co.uk/JohnsonMatthey/index.htm
Introduction:
1. 4 reasons catalysts are used industrially
Principles of Catalysis:
2. Define catalyst
3. Distinguish between Heterogeneous and Homogeneous catalysts
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4. Draw a potential energy curve showing
a reaction that is not catalyzed.
5. Describe or show how the potential
energy curve changes for a reaction
that is catalyzed.
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6. Explain, using the Maxwell-Boltzmann distribution, why a reaction with a lower activation energy
has a higher reaction rate.
How catalysts work:
7. Reactant molecules are _____________________________ at ______________________
onto the surface of the catalyst. This involves the formation of _______________________
between reactant molecules and the catalyst which causes other bonds in the reactant molecule to
be _______________________________________________________. The weakened
structure is converted to another complex that is essentially the
___________________________________to the catalyst. Finally, this complex breaks down to
________________________________________ which moves away to leave the catalyst
surface ready to interact with another reactant molecule.
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Choosing a catalyst:
8. The choice of suitable catalyst for a particular reaction often depends on the
_________________________ _________________________________ formed between
___________________ and ________________ and _____________________ and
_____________________. They must be _______________ enough to form and provide an
______________________ pathway to the uncatalysed reaction but they must not be
_________ ___________________ as this would lead to an increase in the
_______________________________and would _____________________ the rate of
reaction.
Industrial Catalysts:
9. Briefly describe the industrial processes (identify products and reactants and classify the kind of
reaction occurring) and the list the catalyst used:
10. Haber Process
11. Margarine production
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12. Contact Process (Sulfuric Acid)
Catalysts All Around Us
13. List three non-industrial uses of catalysts:
Homogenous Catalysts
14. Summarize why transition metals are often used as homogenous catalysts.
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15. Homogenous catalysts are responsible for the breakdown of ozone in the atmosphere. Briefly
summarize the problem:
Economic Significance Of Catalysts
Based on what you have learned about catalysts, what impact do you think they have on economies?
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Kinetics Review Questions
9. Define rate of reaction
10. What are the units of the rate of reaction?
11. Draw a graph that shows the concentration of products with time as a reaction goes to completion.
Explain the shape of the graph.
12. On the same graph, using a different color or dotted line, show how the concentration of products
changes with time when a catalyst is added.
13. Draw a graph that shows the concentration of reactants with time as a reaction goes to completion.
Explain the shape of the graph.
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14. On the same graph, using a different color or dotted line, show how the concentration of reactants
changes with time if a lower concentration (through an increased volume so that the total number
of moles is the same) is used.
15. True or false: All collisions result in a reaction. Explain.
16. What are the 4 parts of collision theory and why is it important?
17. Give two reasons why a collision would not result in a reaction
18. The reaction between nitrogen and oxygen in the atmosphere under normal conditions is extremely
slow. Which statement best explains this?
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A. The concentration of oxygen is much lower than that of nitrogen
B. The molar mass of nitrogen is less than that of oxygen
C. The frequency of collisions between nitrogen and oxygen molecules is lower than that between
nitrogen molecules themselves
D. Very few nitrogen and oxygen molecules have sufficient energy to react
19. List 4 factors that affect reaction rates
20. Define activation energy
21. Sketch a Maxwell-Boltzmann Energy Distribution Curve for two different temperatures.
22. Does activation energy change with temperature?
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23. How does the number of molecules with the required activation energy change at higher
temperatures? Use energy distribution to explain.
24. Give two reasons why a temperature increase, increases reaction rate and identify the more
important reason.
25. Explain, using the Maxwell-Boltzmann Energy distribution, why adding a catalyst increases reaction
rate.
26. Will the following changes increase the reaction rate of the following reaction when 50 cm3 of 1.0
mol dm-3 HCl is added to 1.0 g calcium carbonate? Explain.
CaCO3(s) + 2 HCl (aq) --> CaCl2 + CO2(g) + H2O(l)
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a. increasing volume of HCl
b. increasing mass of CaCO3
c. decreasing size of CaCO3 particles
d. decreasing amount of CO2 present
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e. increasing temperature
f. adding water to the reaction
g. removing water from the reaction vessel
h. adding a catalyst
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Labeling the Kinetics Revision Map
You have been given a copy of the
kinetics revision map. This shows
some of the important ideas you
have met when you studied
chemical kinetics and related
concepts in the IB Chemistry class.
Each line on the map stands for an
idea that could be put into a
sentence.
The links are not explained on the
map. Read through the statements
below, and work out which link on
the map each sentence is about.
Label each line on the map with the number of the statement – eg
1. The temperature’s effect on rate of reaction is expressed mathematically by k= A e-Ea/RT.
2. The reaction rate depends on surface area.
3. The rate expression can be used to test reaction mechanism.
4. The molecularity can be bimolecular.
5. The slowest elementary step in a reaction mechanism is called the rate determining step.
6. The reaction order can be zero order.
7. The rate expression has the form: Rate = k [reactant] x
8. The half life of a first order reaction is expressed mathematically by t1/2 = 0.693/k
9. The reaction mechanism consists of one or more elementary steps.
10. A catalyst can be homogeneous.
11. The reaction rate can be measured by determining the ! [ ] / !t.
12. The reaction rate depends on activation energy.
13. The molecularity can be unimolecular.
14. The reaction order can be second order.
15. The activation energy’s effect on rate of reaction is expressed mathematically by k = A e-Ea/RT.
16. An elementary step is characterized by its molecularity.
17. The molecularity of the rate determining step is consistent with the rate expression.
18. The reaction rate depends on concentration.
19. The activation energy can be lowered by a catalyst.
20. A catalyst can be heterogeneous.
21. The reaction rate depends on temperature.
22. The reaction order is shown for each reactant in the rate expression.
23. The reaction mechanism consists of a series of elementary steps.
24. The molecularity can be, but is not likely to be, termolecular.
25. The effect of changing concentration can be determined by the initial rate method.
26. The reaction order can be first order.
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Kinetics IB Exam Questions
1.
What is the best definition of rate of reaction?
A.
The time it takes to use up all the reactants
B.
The rate at which all the reactants are used up
C.
The time it takes for one of the reactants to be used up
D.
The increase in concentration of a product per unit time
(Total 1 mark)
2.
Which factors can affect reaction rate?
I.
The state of the reactants
II.
The frequency of the collisions between particles
III.
The average kinetic energy of the particles
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
3.
Equal masses of powdered calcium carbonate were added to separate solutions of
hydrochloric acid. The calcium carbonate was in excess. The volume of carbon dioxide
produced was measured at regular intervals. Which curves best represent the evolution of
carbon dioxide against time for the acid solutions shown in the table below.
25 cm3 of 2 mol dm–3 HCl
50 cm3 of 1 mol dm–3 HCl
25 cm3 of 1 mol dm–3 HCl
A.
I
III
IV
B.
I
IV
III
C.
I
II
III
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D.
II
I
III
(Total 1 mark)
4.
What is the function of iron in the Haber process?
A.
It shifts the position of equilibrium towards the products.
B.
It decreases the rate of the reaction.
C.
It provides an alternative reaction pathway with a lower activation energy.
D.
It reduces the enthalpy change of the reaction.
(Total 1 mark)
5.
Hydrochloric acid is reacted with large pieces of calcium carbonate, the reaction is then
repeated using calcium carbonate powder. How does this change affect the activation
energy and the collision frequency?
Activation energy
Collision frequency
A.
increases
increases
B.
stays constant
increases
C.
increases
stays constant
D.
stays constant
stays constant
(Total 1 mark)
6.
Which statement is true about using sulfuric acid as a catalyst in the following reaction?

H ( aq )
CH3–CO–CH3(aq) + I2(aq) 

 CH3–CO–CH2–I(aq) + HI(aq)
I.
The catalyst increases the rate of reaction.
II.
The catalyst lowers the activation energy for the reaction.
III.
The catalyst has been consumed at the end of the chemical reaction.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
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7.
(a)
A solution of hydrogen peroxide, H2O2, is added to a solution of sodium iodide, NaI, acidified
with hydrochloric acid, HCl. The yellow colour of the iodine, I2, can be used to determine the
rate of reaction.
H2O2(aq) + 2NaI(aq) + 2HCl(aq) → 2NaCl(aq) + I2(aq) + 2H2O(l)
The experiment is repeated with some changes to the reaction conditions. For each
of the changes that follow, predict, stating a reason, its effect on the rate of reaction.
(i)
The concentration of H2O2 is increased at constant temperature.
(2)
(ii)
The solution of NaI is prepared from a fine powder instead of large crystals.
(2)
(b)
Explain why the rate of a reaction increases when the temperature of the system
increases.
(3)
(Total 7 marks)
8.
(a)
Define the term activation energy, Ea.
......................................................................................................................................
......................................................................................................................................
(1)
(b)
State two conditions necessary for a reaction to take place between two reactant
particles.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(2)
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(c)
Sketch an enthalpy level diagram to describe the effect of a catalyst on an
exothermic reaction.
(3)
(Total 6 marks)
9.
Graphing is an important method in the study of the rates of chemical reaction. Sketch a
graph to show how the reactant concentration changes with time in a typical chemical
reaction taking place in solution. Show how the rate of the reaction at a particular time can
be determined.
(Total 4 marks)
10.
(i)
Define the term rate of reaction.
(1)
(ii)
State an equation for the reaction of magnesium carbonate with dilute hydrochloric
acid.
(1)
(iii)
The rate of this reaction in (ii), can be studied by measuring the volume of gas
collected over a period of time. Sketch a graph which shows how the volume of gas
collected changes with time.
(1)
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(iv)
The experiment is repeated using a sample of hydrochloric acid with double the
volume, but half the concentration of the original acid. Draw a second line on the
graph you sketched in part (iii) to show the results in this experiment. Explain why
this line is different from the original line.
(4)
(Total 7 marks)
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Kinetics IB Exam Question Markscheme
1.
D
[1]
2.
D
[1]
3.
C
[1]
4.
C
[1]
5.
B
[1]
6.
A
[1]
7.
(a)
(i)
(ii)
(b)
increases rate of reaction;
molecules (of H2O2) collide more frequently / more collisions
per unit time;
No ECF here.
2
no effect / (solution) remains unchanged;
solid NaI is not reacting / aqueous solution of NaI is reacting / surface
area of NaI is not relevant in preparing the solution / OWTTE;
kinetic energy/speed of reacting molecules increases;
frequency of collisions increases per unit time;
greater proportion of molecules have energy greater than
activation energy/Ea;
Accept more energetic collisions.
2
3 max
[7]
8.
(a)
(b)
(c)
(minimum) energy needed for a reaction to occur / (minimum) energy
difference between reactants and transition state;
particles must collide;
appropriate collision geometry/orientation;
E ≥ Ea;
1
2 max
Diagram showing:
correct labelling of axes (enthalpy/H/(potential) energy for y-axis and
time/progress/course of reaction/reaction coordinate for x-axis) and
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H (products) line shown below H (reactants) line;
correct labelling of the two curves, catalysed and uncatalysed;
correct position of Ea shown with lines for a catalysed and uncatalysed
reaction;
the correct label ∆H /change in enthalpy;
Do not penalize if reactants and products are not labelled.
If an endothermic reaction is shown, award [2 max] if all other parts are
shown correctly.
3 max
[6]
9.
labelled axes (including appropriate units);
correctly drawn curve;
correctly drawn tangent;
y
at time t;
x
[3 max] for straight line graph or graph showing product formation.
rate equal to slope/gradient of tangent (at given time) / rate =
4
[4]
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10.
(i)
(ii)
decrease in concentration/mass/amount/volume of reactant with time /
increase in concentration/mass/amount/volume of product with time /
change in concentration/mass/amount/volume of reactant/product with time;
1
MgCO3(s) + 2HCl(aq) → MgCl2(aq) + CO2(g) + H2O(l);
Ignore state symbols.
1
Plot starts at the origin and levels off.
No mark awarded if axes are not labelled.
1
new curve reaches same height as original curve;
new curve less steep than original curve;
volume of gas produced is the same because the same amount of
acid is used;
reaction is slower because concentration is decreased;
4
(iii)
(iv)
[7]
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