SHS
General Chemistry 2
Quarter 3: Week 7 - Module 7
Reaction Rates, Factors Affecting
Reaction Rates, and Collision Theory
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STEM – GENERAL CHEMISTRY 2
Grade 11/12 Quarter 3: Week 7 - Module 7- Reaction Rates, Factors
Affecting Reaction Rates, and Collision Theory
First Edition, 2021
Copyright © 2021
La Union Schools Division
Region I
All rights reserved. No part of this module may be reproduced in any form
without written permission from the copyright owners.
Development Team of the Module
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General Chemistry 2
Quarter 3: Week 7 - Module 7
Reaction Rates, Factors Affecting Reaction
Rates, and Collision Theory
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Target
Background
Chemical kinetics is the area of chemistry concerned with the speeds, or
rates, at which a chemical reaction occurs.
REACTION RATES
Figure 1: Reaction Rates
The rates of chemical reactions span a range of time scales. For example,
explosions are rapid, occurring in seconds or fractions of seconds; corrosion
can take years; and the weathering of rocks can takes place over thousands
or even millions of years. (images credit: Wikipedia))
After going through with this learning material, you are expected to:
a. Describe how various factors influence the rate of a reaction
STEM_GC11CKIIIi-j-130
b. Differentiate zero, first-, and second-order reactions
STEM_GC11CKIIIi-j-132
c. Explain reactions qualitatively in terms of molecular collisions
STEM_GC11CKIIIi-j-136
Specifically, you will be able to:
1.
2.
3.
4.
5.
6.
define rate of reaction;
identify factors that affect the rate of a reaction;
describe how each factor can affect the rate of a reaction;
determine the order of a reaction;
state the premises of the Collision Theory; and
explain the effect of each factor on the rate of a reaction using the
Collision Theory.
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Jumpstart
Our goal in this module is to understand how to determine the rates at
which reactions occur and to consider the factors that control these rates.
Activity 1: Chemical Reactions in Everyday Life
For each of the chemical reactions taking place from the picture, identify
the occurrence of each reaction (fast or slow) and the factors that control the
rate of the reaction.
Figure 2: Examples of Chemical Reactions in Everyday Life
(https://www.thoughtco.com/examples-of-chemical-reactions-in-everyday-life)
Activity 2. Answer the following questions
a. How long does it take an iron nail exposed to the rain to rust?
______________________________________________________________________________
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b. Compare the rusting of iron to how fast milk curdles when an acid like
vinegar or calamansi juice is added to it. (Try doing this at home if you have
milk, vinegar or calamasi juice.)
______________________________________________________________________________
c. Do you think you could light a log with a single matchstick? How about twigs
or smaller pieces of wood?
______________________________________________________________________________
d. Why do we keep food in the refrigerator?
______________________________________________________________________________
e. How do particles move at high temperatures compared at low temperatures?
______________________________________________________________________________
Discover
The Rate of a Chemical Reaction
The study of the rate of a chemical reaction is called chemical kinetics.
What is the “rate" of a reaction?
The rate of a reaction is the speed of the reaction. It is not <how much= of a
product is made, but instead <how quickly= a reaction takes place.
Reaction Rate = Reactants → Products
Time
Reaction Rate is defined as the rate at which reactants change into
products over time.
How can we measure the rate?
➢ The rate of a chemical reaction can be studied by measuring either:
• How quickly a reactant is used up, or
• How quickly a product is formed.
Consider the reaction:
REACTANTS
zinc + hydrochloric acid 4>
0
PRODUCTS
zinc chloride + hydrogen
0
There are two possible ways of measuring the rate:
1) measure how quickly one of the products (e.g. the hydrogen) is made
2) measure how quickly one of the reactants (e.g. the zinc) disappears
Factors That Affect Rate of a Chemical Reaction
Almost any reaction rate can be changed by varying the conditions under
which the reaction takes place
1. Nature of the reactants
• Substances vary in their chemical reactivity. It is a major factor that
determines the rate of a reaction.
• Chemists believe that the type, strength and number of chemical bonds
or attractions between atoms determines with what speed the particles
have to collide with each, to create an effective reaction.
a. Reactions involving simple ions are most often instantaneous. This is due to
the fact that the positive and negative charges attract each other and no bonds
have to be broken in creating the new substances.
IONS = FAST
Cr 2+ (aq)+Fe 3+(aq)-->Cr 3+(aq)+Fe 2+(aq)
b. Reactions between molecules are usually slower than ions. In molecules
bonds have to be broken and new bonds reformed. This slows down reaction
rates.
MOLECULES = SLOW
2CO(g)+O2(g)-->2CO2(g)
c. Gases tend to react faster than solids or liquids:
➢ It takes energy to separate particles from each other. In order to burn
candle wax, the solid wax has to be melted and then vaporized before it
reacts with oxygen. Methane gas is already in the gas state so it burns
faster than wax.
Fast
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
Slow
C25H52(s) + 38O2(g) →
25CO2(g) + 26H2O(g)
d. Aqueous ions tend to react faster than species in other states of matter.
➢ Solid Pb(NO3)2 will react with solid KI, but the reaction is really, really slow.
That's because the ionic bonding in each reactant is strong and the ions in
each compound are hard to separate from each other.
➢ When aqueous solutions of these compounds are mixed, the formation of
PbI2 iodide is rapid. In aqueous solutions, the ions of each compound are
dissociated. When the two the solutions are mixed together, all that is
required for a reaction to occur is contact between the lead(II) ions and the
iodide ions.
Fast
Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + 2KNO3(aq)
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0
Slow
Pb(NO3)2(s)
+ 2KI(s) → PbI2(s) + 2KNO3(s)
e. Reactions involving the breaking of fewer bonds per reactant proceed faster
than those involving the breaking of a larger number of bonds per reactant.
➢ Kerosene burns more slowly than methane because there are more bonds
to be broken per molecule of kerosene than there are per molecule of
methane. Kerosene is a larger molecule.
Fast
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
methane
Slow
C13H28(g)
+
20 O2(g)
→ 13 CO2(g)
+
14 H2O(g)
kerosene
2. Concentration of reactants
• Concentration refers to how much solute is dissolved in a solution.
• When the concentration of all the reactants increases, more molecules or
ions interact to form new compounds, and the rate of reaction increases.
• When the concentration of a reactant decreases, there are fewer of that
molecule or ion present, and the rate of reaction decreases.
In
a
less
concentrated acid,
the number of
collisions is low,
so the rate of the
reaction is slower.
3M HCl
6M HCl
With
a
more
concentrated acid,
the number of acid
particles is greater,
so the number of
collisions is greater
and the rate of the
reaction is higher
(faster).
There are more particles of acid per unit volume
in the 6M acid than there are in the 3M acid.
3. Temperature of reaction system
• Generally speaking, an increase in temperature will increase the reaction
rate.
• Similarly, decreasing temperature will decrease the reaction rate.
4. Presence of a catalyst
• A catalysts is a substance that affects the reaction rate without being
used up in the reaction.
• The presence of a catalyst usually increases reaction rate.
5. Surface Area
• Surface area is the exposed matter of a solid substance.
0
0
•
•
An increase in surface area will result in an increase of the exposure of
reactants to one another.
The greater the exposure, the greater the reaction rate.
8Inner9 particles are protected
and cannot collide with other
particles until they become
8exposed9.
The surface
particles are
8exposed9 and
can react.
There is now a
greater surface area
with more exposed
particles so more
collisions can
occur, hence faster
reaction.
If we break up this
8lump9 into smaller
pieces the number of
particles has not
changed but there
are now more
8surface9 particles.
The Rate Law
➢ The rate of a reaction often depends on the concentration of one or more
of the reactants.
when A is a single reactant
Rate Law
A → products
r = k[A]n
Where k is a constant of proportionality called the rate constant and n
is the reaction order
Example:
Reaction
Rate Law
2H2O2 (a) → 2H2O(l) + O2(g)
r = k[H2O2]
• when A and B are reactants and C is the product
Rate Law
A + bB → C
r = k[A]x[B]y
Where [A] and [B] express the concentrations of A and B, in units of moles/liter.
• The exponents x and y reaction orders
• k is known as the rate constant of the reaction.
Example:
Reaction
2NO(g) + O2(g) → 2NO2(g)
Rate Law
r = k[NO]2 [O2]
Note that the reaction order is unrelated to the stoichiometric coefficients of
the reaction; it must be determined experimentally.
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Reaction order
From the rate law of given concentration of reactants, r = k[A]n and r =
k[A]x[B]y, the exponents n, x and y are not from the balanced equation, they
must be determined experimentally. Their values determine how the rate
depends on the concentration of the reactants.
•
•
•
If the exponent n, and x and y = 0, the reaction is zero order and the rate
is independent of the concentration of the reactants.
If exponent n, and x and y = 1, the reaction is first order and the rate is
directly proportional to the concentration of the reactants.
If exponent n, and x and y = 2, the reaction is second order and the rate
is proportional to the square of the concentration of A.
Figure 3: Reaction concentration as a
function of time for different reaction
orders
(Image credit: Chemistry: A Molecular Approach)
A. Zero-Order Reaction
In a zero-order reaction, the rate
of the reaction is independent of the
concentration of the reactant.
r = k[A]0
•
•
• For
a
zero-order
reaction,
increasing the concentration of the reacting species will not speed up the
rate of the reaction.
A reaction is zero-order if concentration data is plotted versus time and
the result is a straight line. The concentration of the reactant decreases
linearly with time.(See Figure 3)
The slope of the line is constant indicating a constant rate, because the
reaction does not slow down as the concentration of the reactant
decreases.
B. First-Order Reaction
In a first-order reaction, the rate of the reaction is directly proportional to
the concentration of the reactant.
r = k[A]1
0
0
•
•
A reaction that depends on the concentration of only one reactant (a
unimolecular reaction). Other reactants can be present, but each will be
zero-order.
For a first-order reaction the rate slows down as the reaction proceeds
because the concentration of the reactant decreases.
C. Second-Order Reaction
In the second-order reaction, the rate of the reaction is proportional to
the square of the concentration of the reactant.
r = k[A]2 or r = k[B]2
• A second-order reaction will depend on the concentration (s) of one
second-order reactant or two first-order reactants.
How to determine the order of a reaction experimentally?
• The method of determining the order of a reaction is known as the
method of initial rates.
Method of Initial Rates is measured by running the reaction several times
with different initial reactant concentrations to determine the effect of the
concentration on the rate.
Example 1: First-Order Reaction
Let9s say that at 25 °C, we observe that the rate of decomposition of N 2O5
is 1.4×10-3 M/s when the initial concentration of N 2O5 is 0.020 M. Then, let9s
say that we run the experiment again at the same temperature, but this time
we begin with a different concentration of N 2O5, which is 0.010 M. On this
second trial, we observe that the rate of decomposition of N2O5 is 7.0×10-4 M/s.
The balanced chemical equation for the decomposition of dinitrogen pentoxide
is
2N2O5(g) → 4NO2(g) + O2(g)
Since there is only one reactant, the rate law for this reaction has the general
form:
r = k[A]n
substituting the given reactant for A: r = k[N2O5]n
Step 1: Tabulate the given data
[N2O5] (M)
Initial Rate (M/s)
0.020
1.4×10-3
0.010
7.0×10-4
Step 2: Set up a ratio of the first rate to the second rate
Rate 1 = k[N2O5]n1
Rate 2
k[N2O5]n2
Step 3: Substitute and simplify the equation.
1.4×10-3 = k[0.020]n
7.0×10-4 k[0.010]n
0
0
Notice that the left side of the equation is simply equal to 2, and that the rate
constants cancel on the right side of the equation. Everything simplifies to:
2 = 2n
Clearly, then, n =1, and the decomposition is a first-order reaction.
rate = k[N2O5]1
Determining the Rate Constant k
Once we have determined the order of the reaction, we can go back and
plug in one set of our initial values and solve for k. We find that:
rate = k[N2O5]1 = k[N2O5]
Substituting in our first set of values from our table, we have
1.4×10-3 M/s = k[0.020 M]
0.020 M
0.020 M
k = 0.07 /s
Example 2: Second-Order Reaction
Consider the following set of data:
Trial
Reactant A (M)
Reactant B (M)
1
0.200
0.200
2
0.300
0.200
3
0.200
0.400
Initial Rate (M/s)
5.46
12.28
5.42
Determining Reaction Order in A
In order to determine the reaction order for A, we need to choose two
experimental trials in which the initial concentration of A changes, but the initial
concentration of B is constant, so that the concentration of B cancels. We can set
up our first equation by picking Trials 1 and 2 as follows:
r1 = k[A]x1[B]y1
r2 k[A]x2[B]y2
5.46 = k[0.200]x [0.200]y
12.28 k[0.300]x [0.200]y
5.46 = k[0.200]x
12.28 k[0.300]x
0.444 = (²∕3)x
log (0.444) = log (²∕3)x
log (0.444) = x log (²∕3)
log (0.444) = x log (0.67)
log (0.67) log (0.67)
x≈2
Note: log Ax = x log A
Use scientific calculator to
get the log value.
Therefore, the reaction is second-order in A. Rate Law is r = k[A]2[B]y
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Determining Reaction Order in B
To determine the reaction order for B we have to pick two trials in which
the concentration of B changes, but the concentration of A does not. Trials 1 and
3 will do this for us, and we set up our ratios as follows:
r1 = k[A]x1[B]y1
r3 k[A]x3[B]y3
5.46 = k[0.200]x [0.200]y
5.42
k[0.200]x [0.400]y
5.46 = [0.200]y
5.42
[0.400]y
1 = (½)y
Any non-zero number to the
If y = 0 then
zero power equals one.
1=1
Therefore, the reaction is zero-order in B. Rate law is r = k[A]2[B]0
Overall Reaction Order
• The overall order of a reaction is the sum of all the exponents of the
concentration terms in the rate equation.
• We have determined that the reaction is second-order in A, and zero-order
in B. Therefore, the overall order for the reaction is second-order (2+0 =
2) , and the rate law will be: rate = k[A]2
The Collision Theory
Collision theory states that for a chemical reaction to occur, the reacting
particles must collide with one another.
• In a chemical reaction, products are formed due to the collision between
the reactant molecules.
• The conditions for the collisions to form products are:
Effective Collisions
Collision
Theory
Right orientation
Sufficient energy
A. Collisions should be effective.
✓ The reactant particles must collide with each other.
Collisions between reactant molecules convert the reactants into the products of
the reaction.
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But remember…
NOT EVERY COLLISION LEADS TO REACTION!!!!!
In order for collisions to be effective, there must be considerable FORCE
in the collisions.
The slower moving molecules do not have enough kinetic energy to react
when they collide...they bounce off one another and retain their identity.
Only those molecules moving at high speed have enough energy for
collisions to result in a reaction.
Head-on collision:
More energetic – reaction takes place
Glancing blow:
Less energetic – no reaction
Consider the reaction: NO3(g) + CO(g) ----->
NO2(g) + CO2(g)
O
O
N
O
O
N
O
O
C
O
C
C
O
B. The right orientation of reactant molecules towards each other.
Consider another reaction between ClNO2 and NO.
ClNO2(g) + NO(g) 9------: NO2(g) + ClNO(g)
Reaction won't occur if
the oxygen end of the
NO molecule collides
with the chlorine atom
on ClNO2
Nor will it occur if
one of the oxygen
atoms on ClNO2
collides with the
nitrogen atom on
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In the course of this
reaction, a chlorine
atom is transferred
from one nitrogen atom
to another. In order for
the reaction to occur,
the nitrogen atom in
NO must collide with
the chlorine atom in
ClNO2.
Based on Collision Theory, reactions occur when two reactant molecules
effectively collide, each having minimum energy and correct orientation.
O
O
N
O
O
N
O
C
O
C
O
NO3(g) + CO(g) 9-----> NO2(g) + CO2(g)
c. All molecules should possess a minimum amount of energy to form
product molecules.
➢ The minimum amount of energy required to initiate a chemical reaction
is called activation energy, Ea (expressed in kJ).
(Activation energy will be further discuss in the next module.)
Explore
Determining Reactant Orders from Actual Data
1. Given the following data, determine the
(a) order of reaction of reactant A,
(b) order of reaction of reactant B, and
(c) total reaction order for the equation.
Experiment
[NO] (M)
[Cl2] (M)
1
0.0300
0.0100
2
0.0150
0.0100
3
0.0150
0.0400
Rate (M/s)
3.4 × 1034
8.5 × 1035
3.4 × 1034
2. The initial rate data for the reaction 2N2O5(g) → 4NO2(g) + O2(g) is shown in the
following table.
Experiment
[N2O5](M)
Rate (M/s)
1
1.28 × 102
22.5
2
2
2.56 x 10
45.0
a. Determine the order of reaction of the given reactant.
b. Determine the value of the rate constant for this reaction.
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Deepen
Use the equation to answer the questions.
C6H12O6 (s) + 6 O2(g) → 6 H2O
(g)
+ 6 CO2 (g)
1. What happens to the concentrations of:
a. C6H12O6 and O2 as the reaction proceeds →?
_____________________________________________________________________________
b. H2O + CO2 as the reaction proceeds →?
______________________________________________________________________________
2. According to the collision theory, what 3 circumstances are needed for
C6H12O6 & O2 to react?
______________________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
3. Complete the table below.
Does this
increase or
Explain why using
Change in condition:
decrease the rate
Collision Theory.
of reaction?
a. Increasing the temperature
Ex: Increases (speeds
up)
b. Increasing the
concentration of C6H12O6
c. Decreasing the
concentration of O2
d. Increase the surface area by
chewing up food in your mouth
e. Decreasing the temperature
f. Increasing the pressure in
the container
g. Decreasing the
concentration of H2O
h. Increasing the volume of the
container the reaction occurs in
i. Increasing the
concentration of CO2
j. Using a catalyst (like salivary
amylase)
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0
Ex: Molecules move faster
& collide more
= Increased reaction rate
Gauge
Answer the following. Use a separate sheet for your answer.
1. For a reaction 2A + B → 2C, with the rate equation: Rate = k[A]2[B]. What is
the correct reaction order?
A. the order with respect to A is 1 and the order overall is 1
B. the order with respect to A is 2 and the order overall is 2
C. the order with respect to A is 2 and the order overall is 3
D. the order with respect to B is 2 and the order overall is 2
2. Which one of the following is NOT a key concept of the collision theory?
A. particles must collide in order to react
B. particles must move slowly when they collide, otherwise they simply
<bounce off= one another
C. particles must collide with the proper orientation
D. particles must collide with sufficient energy to reach the activated
complex in order to react
3. Which of the following statements is correct?
A. The rate of a reaction decreases with passage of time as the
concentration of reactants decreases.
B. The rate of a reaction is same at any time during the reaction.
C. The rate of a reaction is independent of temperature change.
D. The rate of a reaction decreases with increase in concentration of
reactant(s).
4. What will you do if you wanted to create a more space for particles to collide?
A. Raise the pressure
B. Raise the temperature
C. Increase the concentration
D. Increase the surface area
5. If the rate law for the reaction 2A + 3B → products is first order in A and
second order in B, then what is the correct rate law equation?
A. k[A]2[B]3
B. k[A]2[B]2
C. k[A]2[B]
D. k[A][B]2
6. Which is a factor that would not affect the rate of a reaction?
A. Temperature
B. Adding a catalyst
C. Concentration
D. The weather
7. The rate law for a reaction is k[A][B]2. Which one of the following statements
is false?
A. The reaction is first order in A.
B. The reaction is second order in B.
C. The reaction is second order overall. D. k is the reaction rate constant
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8. As the temperature of a reaction is increased, the rate of the reaction
increases. Why?
A. activation energy is lowered
B. reactant molecules collide less frequently
C. reactant molecules collide more frequently and with greater energy per
collision
D. reactant molecules collide less frequently and with greater energy per
Collision
9. A reaction was found to be zero order in A. If we increase the concentration
of A by a factor of 3, what will happen to the reaction rate?
A. remain constant
B. increase by a factor of 9
C. increase by a factor of 27
D. triple
10. The rate law of the overall reaction is k[A][B]0. Which of the following will
not increase the rate of the reaction?
A. increasing the concentration of reactant A
B. increasing the concentration of reactant B
C. increasing the temperature of the reaction
D. adding a catalyst for the reaction
11. Which of the following would NOT increase the rate of reaction?
A. raising the temperature
B. adding catalyst
C. increasing the volume of the container
D. increasing the concentration of the reactants
12. What is the correct expression of the rate of a chemical reaction?
A. grams per mole
B. energy consumed per mole
C. molarity per second
D. volume of gas per unit time
13. Why an increase in temperature increases a reaction rate?
A. only because there are more collisions
B. only because the particles collide with greater energy
C. by providing an alternative pathway for the reaction to follow
D. because the particles collide more frequently and with greater energy
14. Which condition will increase the rate of a chemical reaction?
A. Increased temperature and decreased concentration of reactants
B. Increased temperature and increased concentration of reactants
C. Decreased temperature and decreased concentration of reactants
D. Decreased temperature and increased concentration of reactants
15. In terms of collision theory, When is a chemical reaction occurs?
A. chemical bonds are broken by a catalyst.
B. two reactant molecules collide with each other.
C. reactants collide with sufficient energy to form new bonds.
D. reactants collide with sufficient energy to break bonds in the reactants
0
0
0
1. C 2. B 3. A 4. D 5. D 6. D 7. C 8. C 9. A 10. B 11. C 12. C 13. D 14. B 15. D
Change in condition:
a. Increasing the
temperature
b. Increasing the
concentration of C6H12O6
c. Decreasing the
concentration of O2
d. Increase the surface
area by chewing up food in
your mouth
e. Decreasing the
temperature
f. Increasing the pressure
in the container
g. Decreasing the
concentration of H2O
h. Increasing the volume
of the container the
reaction occurs in
i. Increasing the
concentration of CO2
j. Using a catalyst (like
salivary amylase)
Does this
increase or
decrease the rate
of reaction?
Ex: Increases
(speeds up)
Increases
Decreases
Increases
Decreases
Increases
Decreases
Decreases
Increases
increases
GAUGE:
0
Explain why using Collision Theory.
Ex: Molecules move faster & collide more
= Increased reaction rate
With more concentration, the number of
C6H12O6 particles is greater, so the number of
collisions is greater and the rate of the reaction
is higher (faster).
Less particles, less chances of collision.
Increasing the surface area increases the
number of collisions per second and therefore
increases the reaction rate.
Molecules move slower decreasing the chance
of effective collision.
When you increase the pressure, the molecules
have less space in which they can move. That
greater density of molecules increases the
number of collisions.
No effect, only reactants are involved in the
reaction.
Increasing the volume increases the space.
Wider or larger space decreases the chance of
collision.
No effect, only reactants are involved in the
reaction.
A catalyst speeds up the reaction because it
supplies the energy needed in the reaction.
3.
a. 2nd-order
b. 1st-order
c. 2 + 1 = 3
EXPLORE
1. a. decreases
b. increases
1.
c. sufficient energy
2.
2. a. effective collision b. correct orientation
a. 1st-order
b. k = 0.18
DEEPEN
Answer Key
References:
Chang, Raymond. (2010). Chemistry.-10th edition. USA. McGraw-Hill. pp. 556 3
562
Tro, Nivaldo J. (2011). Chemistry: a molecular approach. USA. Pearson Prentice
Hall. pp. 562 - 569.
<Factors that Affect Reaction Rate=, accessed December 23, 2020,
https://courses.lumenlearning.com/introchem/chapter/factors-that-affectreaction
<Factors Affecting Rate of Reaction=,
https://www.bcsoh.org/cms/lib3
accessed
December
26,
2020,
<Rates of reaction=, accessed December 31, 2020, https://whs.rocklinusd.org
<Rate
of
Reaction=,
accessed
https://byjus.com/chemistry/rate-of-reaction
December
25,
2020,
<The
Collision
Theory=,
accessed
December
28,
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https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry
_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry
)/Kinetics/_Modeling_Reaction_Kinetics/_Collision_Theory/_The_Collision_The
ory
0
0