Chemical Kinetics
From CK12 - Flexbooks
| Author: Therese Forsythe, | (ThereseF) |
| Chapter Info | |
|---|---|
| Title: | Chemical Kinetics |
| Part of: | Chemistry |
| License: | CC BY SA |
| Subject: | |
| Grade level: | 12, 11, 10, 9 |
| Difficulty: | Intermediate |
| About: | This chapter covers reaction rates and the factors that affect reaction rates. |
| Keywords: | reaction, reaction rates |
| Readiness: | CR-2 |
| Warnings: |
|
| | |
| Nature of Science | |
|---|---|
| Science Inquiry | |
| Science and Technology | |
| Science and Society | |
| Science Measurement and Data | |
| Semantic Properties List | |
| Category - Earth Science | |
|---|---|
| Earth Systems | |
| Weather and Climate | |
| Astronomy | |
| Semantic Properties List | ,, |
| Category - Life Science | |
|---|---|
| Living Systems | |
| Plant Biology | |
| Animal Biology | |
| Human Biology | |
| Genetics | |
| Reproduction and Development | |
| Evolution | |
| Semantic Properties List | , , , , , , |
| Category - Physical Science | |
|---|---|
| Motion and Forces | |
| Energy and Heat | Thermal Energy and Changes of State, Heat Transfer, Temperature — Thermal Energy — Heat, Transformations of Energy, Forms of Energy, Conservation of Energy, Kinetic Energy, Potential Energy |
| Wave-Sound-Light | |
| Electricity and Magnetism | States of Matters — Changes of State, Elements — Molecules — Compounds, Characteristics of Matter, Properties of Matter, Behavior of Gases, Behavior of Liquids, Behavior of Solids, Solutions — Mixtures |
| Elements and the Periodic Table | |
| Chemical Formulas and Reactions | Stoichiometric Relationships, Kinetics — Equilibrium — Reaction Rates |
| Acids-Bases-Solutions | |
| Organic Chemistry | |
| Nuclear Chemistry | |
| Semantic Properties List | ,Thermal Energy and Changes of State
, Heat Transfer , Temperature — Thermal Energy — Heat , Transformations of Energy , Forms of Energy , Conservation of Energy , Kinetic Energy , Potential Energy ,,States of Matters — Changes of State , Elements — Molecules — Compounds , Characteristics of Matter , Properties of Matter , Behavior of Gases , Behavior of Liquids , Behavior of Solids , Solutions — Mixtures ,,Stoichiometric Relationships , Kinetics — Equilibrium — Reaction Rates ,, , |
| Category - Mathematics | |
|---|---|
| Number Sense | |
| Number Operations | |
| Measurement | |
| Consumer Applications | |
| Basic Algebra | |
| Advanced Algebra | |
| Basic Geometry | |
| Advanced Geometry | |
| Data Displays | |
| Statistics | |
| Probability | |
| Trigonometry | |
| Mathematical Analysis | |
| Special Mathematical Topics | |
| Functions | |
| Instructional Technology | |
| Semantic Properties List | |
Rate of Reactions
Lesson Objectives
- Define chemical kinetics and rates of reactions.
- Write the rate expression and the units for the rate expression.
- Define instantaneous rate.
- Calculate instantaneous rate using a tangent line.
Introduction
In this chapter the study of chemical kinetics will be explored. Chemical kinetics is the study of rates of chemical reactions and what factors affect the rate of reactions. These factors include concentration, temperature, pressure, surface area, and the effect of a catalyst. Many chemical reactions are required to be sped up or slowed down depending on the nature of the reaction. For example, when food is placed in the refrigerator the temperature is decreased to keep the food from decomposing. The rate of reaction is slowed down by decreasing the temperature. Chemical kinetics plays an important role both in industry and in our daily lives. As we work through this chapter gaining insight into the background of this topic, we will learn more about how the science and the math of this wonderful branch of chemistry work in our world. To begin, we will introduce some of the basic concepts of chemical kinetics.
Change in Concentration Over Time
The term rate of reaction is used to denote the measure of the rate at which the products are formed in a time interval or the rate at which the reactants are consumed over a time interval. A reaction rate measures how fast or how slow a reaction is. The rusting of a piece of metal, like a car part, has a slow reaction rate because the iron oxidizes in the air in a relatively long time period. A forest fire has a fast reaction rate since it consumes trees in its path in a very short time interval. Reaction rates can be measured in change in mass per unit time (grams/second) or in charge in molarity per unit time (mols/liter/second).
Symbolically, the reaction rate is given the letter 
in expressions. Therefore to write the expression for the reaction rate, you can write the following.
Remember that the symbol 
means the change in.
Sample question:
For the reaction 
, under certain conditions the 
at 
and 
at 
. What is the rate of production of 
? Note: Remember that the brackets indicate concentration
Solution:
Therefore, the rate of production of is 
.
Units for Rate of Reaction
Notice in the sample question in the previous section that the units to measure the reaction rate are in 
. Therefore, the units are measuring the concentration/time or the M/time. This expression of the units for reaction rate is consistent to allow for comparison of rates. In other words, if all reaction rates were to use the same units, we can compare one rate to the other. For example, with the HI reaction under a different set of conditions, we found that the reaction rate was found to be 
for these conditions. Then, we could predict that the new set of conditions are favorable for this reaction since the reaction rate was faster for the production of HI in the same time interval.
Graphing Instantaneous Rate
Instantaneous rate is defined as the rate of change at a particular moment. For example, a police officer stops a car for speeding. The radar gun on a police cruiser is set to measure the speed of a motorist as the motorist comes close to the cruiser. The driver of the vehicle is stopped doing 
in a 
zone. The cruiser measured the rate of speed at that instant in time when the driver passed the police officer. This is instantaneous rate. If we were to take all of the measures of instantaneous rate and graph them, we would obtain a curve of the overall speed (or the average speed) of the vehicle. The same is true for reactions. For reactions, the instantaneous rate is the rate of the reaction at a specific time in the reaction sequence. If you were to graph the rate of the reactant being consumed versus time, the graph would look like Figure below. As the reaction proceeds, the concentration of the reactants decreases over time.
| Todo 10 | Replace reference to Figure 1 with variable Che-2501-01 |
|---|
The initial rate of the reaction is found when 
, or when the reaction is just beginning. It is at this point when the maximum amount of the reactant is present. To find the instantaneous rate, or the rate at any given time in the reaction, a tangent line is drawn to this curve. Then the slope of the tangent line is found for a point on the curve. For example, we wanted to know the instantaneous rate at 
. After drawing the tangent line (see Figure below), we can calculate the slope of the tangent line to find the instantaneous rate at .
| Todo 10 | Replace reference to Figure 2 with variable Che-2501-02 |
|---|
Lesson Summary
- Chemical kinetics is the study of rates of chemical reactions and how factors affect rates of reactions. The term rate of reaction is used to denote the measure at which the products are formed over a time interval or the rate at which the reactants are consumed over a time interval.
- The units to measure the reaction rate are in
. Instantaneous rate is defined as the rate of change at a particular moment.
Review Questions
- What is the rate of production of
in the system below given that the concentration of is 
at 
and 
at 
.
- For the graph below, draw a tangent line at
, and calculate the instantaneous rate.
- Which expression represents the rate for the product formation of the following reaction?
-
-
-
- All of these are accurate representations of the rate.
-
- The speed of a car is
.
- Half the product is produced.
- A family consumes
of milk.
- I ran for
- The higher the rate, the lesser the slope of a line on a concentration-time graph.
- The instantaneous rate is the slope of the tangent to a line on a concentration-time graph.
- The instantaneous rate is the slope of the cosine to a line on a concentration-time graph.
- All of these statements are correct.
gas if the concentration decreases from 
at 
and 
at 
?
for the concentration of a reactant to decrease from 
to 
What is the rate of reaction in 
Vocabulary
- chemical kinetics
- The study of rates of chemical reactions and how factors affect rates of reactions.
- rate of reaction
- The measure at which the products are formed over a time interval or the rate at which the reactants are consumed over a time interval.
- instantaneous rate
- The rate of change at a particular time interval.
Collision Theory
Lesson Objectives
- Define the collision theory.
- Describe the conditions for successful collisions.
- Explain how the kinetic molecular theory applies to the collision theory.
- Describe the rate in terms of the conditions of successful collisions.
Introduction
Consider the chemical reaction 
. In the reactants, the carbon atoms are bonded to hydrogen atoms and the oxygen atoms are bonded to oxygen atoms. Each atom in the reactants is bonded to its full capacity and therefore, cannot form any more bonds. In the products, the carbon atoms are bonded to oxygen atoms and the hydrogen atoms are bonded to oxygen atoms. The bonds that are present in the products cannot form unless the bonds in the reactants are first broken. Breaking bonds requires an input of energy. If the atoms in the reactants were in the form of carbon atoms, hydrogen atoms, and oxygen atoms, they would have a much higher potential energy that they have in the form of 
and 
molecules. Therefore, in order to get the reactant molecules into a form that will allow them to form new bonds, the old bonds must be broken and this requires that their potential energy be raised considerably.
The energy to break the old bonds comes from the kinetic energy of the reactant particles. The reactant particles are moving around in their random motion and their average kinetic energy is related to their temperature. If a reaction is to occur, the kinetic energy of the reactants must be high enough that when the reactant particles collide, the collision is forceful enough to break the old bonds. Once the old bonds are broken, the atoms in the reactants would be available to form new bonds. At that point, the new bonds of the products could be formed. When the new bonds are formed, potential energy is given off. The potential energy that is given off becomes kinetic energy and is absorbed by the surroundings (primarily the products, the solvent solution if there is one, and the reaction vessel).
Chemists have chosen to give a name to the group of particles that exist for a split second just after the reactant bonds have been broken and before the product bonds form. This group of un-bonded particles is called the activated complex. The activated part comes from the fact that these atoms are ready to form bonds and the complex part comes from the fact that the group of particles is a jumble of particles from all the reactant molecules. A successful collision would proceed as follows:
This collision is successful and results in reaction. (Source: Richard Parsons. CC-BY-SA)
| Todo 10 | Replace reference to Figure 3 with variable Chem-1701-AC |
|---|
All three groups of particles, the reactants, the activated complex, and the products have a precise amount of potential energy in their bonds. The potential energy of the activated complex is called the threshold energy. This threshold energy is the minimum potential energy that must be reached in order for a reaction to occur. The input of energy that is necessary to raise the potential energy of the reactants to this threshold energy is called the activation energy. The activation energy must be provided from the kinetic energy of the reactant particles during the collision. In those cases where the reactants do not collide with enough energy to break the old bonds (inadequate activation energy), the reactant particles will simply bounce off each other and remain reactant particles.
How Reactions Occur
We know that a chemical system can be made up of atoms 
, ions 
), or molecules 
. We also know that in a chemical system, these particles are moving around in a random motion. The collision theory explains why reactions occur at this particle level between these atoms, ions, and/or molecules. The collision theory provides us with the ability to predict what conditions are necessary for a successful reaction to take place. These conditions include:
- The particles must collide with each other.
- The particles must have proper orientation.
- The particles must collide with sufficient energy to break the old bonds.
A chemical reaction involves breaking bonds in the reactants, re-arranging the atoms into new groupings (the products), and the formation of new bonds in the products. Therefore, not only must a collision occur between reactant particles but the collision has to have sufficient energy to break all the reactant bonds that need to be broken in order to form the products. Some collision geometries need less collision energy than others. The optimal collision geometry requires the smallest amount of particle kinetic energy for the reaction to occur. The amount of kinetic energy the reactant particles must have in order to break the old bonds is called the activation energy. The activation energy is the minimum amount of kinetic energy the reactant particles must have in order for a successful collision to occur, assuming optimal geometry. If the reactant particles collide with less than the activation energy, the particles will rebound (bounce off each other), and no reaction will occur.
The reaction rate, discussed in the previous section is proportional to the number of successful collisions per second.
Collision Theory and the Kinetic Molecular Theory
The kinetic molecular theory provides the foundation for the collision theory. The kinetic molecular theory tells us about the particles involved in the collisions. Part of the kinetic molecular theory maintains that the collision between particles are "perfectly elastic". The term "perfectly elastic" is a term from physics which means that kinetic energy in conserved in the collision. That is, if no bonds are broken so no 
is converted to 
, the colliding particles simply rebound and total after the collision is exactly the same as the total before the collision. The kinetic molecular theory states that gas molecules consist of particles that are moving in random motion. This random motion is always in a straight line and the particles deviate only when there is a collision with the walls of a container or with another particle. The only collisions of any consequence are those between other particles. The kinetic energy, defined as the energy of motion, is dependent on the temperature of the system. As the temperature is increased, the kinetic energy of the particles increases and the particles move at a faster pace. With the particles moving faster, they are more likely to collide and will collide with more energy.
In the chapter on kinetic-molecular theory, it was discussed that the particles in a sample of material are not all at exactly the same temperature. The temperature of a substance is the average kinetic energy of all the particles. Some of the particles have more than the average kinetic energy and some have less. The particles of the substance actually have a distribution of kinetic energies and the temperature of the substance is an expression of the average kinetic energy. Therefore, it is not only possible but likely, that in the mass of reactants for a reaction, some of the reactant particles will have sufficient to react and some of the reactant particles will not.
In a slow reaction, the majority of molecules do not have the minimum amount of energy necessary for a reaction to take place. In Figure below, the graph illustrates the number of molecules in the system versus the kinetic energy of these molecules. The area under the curve represents the total number of particles. The area shaded in red shows the number of molecules that DO have sufficient energy for a successful collision. The minimum amount of energy necessary for a reaction is known as the activation energy.
| Todo 10 | Replace reference to Figure 4 with variable Che-2502-01 |
|---|
If the temperature is increased, the average kinetic energy of the particles increases and the number of molecules with sufficient kinetic energy for a successful collision would also increase. Figure below shows changes for the increased temperature.
| Todo 10 | Replace reference to Figure 5 with variable Che-2502-02 |
|---|
At the higher temperature 
, the number of molecules with energy greater than the activation energy increases. Therefore the number of molecules with kinetic energy great enough to have successful collisions increases with increasing average kinetic energy.
Reactions May Occur When Particles Collide
Looking back at the three conditions introduced in the first section, consider the following reaction.
If there is not enough energy, the particles will simply rebound off each other and bonds will not be broken. The original reactants will remain.
An unsuccessful collision results in rebound. (Source: Therese Forsythe. CC-BY-SA)
| Todo 10 | Replace reference to Figure 6 with variable Che-2502-03 |
|---|
In order to have a successful collision, the particles must collide with enough energy and with the correct geometry to break the 
and 
bonds and form the 
and 
products. The would then further react with another element as it is not normally found unreacted (or just as alone).
An successful collision results in reaction. (Source: Therese Forsythe. CC-BY-SA)
| Todo 10 | Replace reference to Figure 7 with variable Che-2502-04 |
|---|
Rate of Reaction Dependent On Various Factors
As stated at the beginning of the lesson, there are three conditions that must occur in order for a successful collision to occur. First of all, the reactant particles must collide. The total number of collisions per second is known as the collision frequency, whether these collisions are successful or not. The collision frequency depends on the concentration of the particles in the container, the temperature of the reaction and the size of the particles themselves. Second, the particles must collide with the proper orientation. And third, the particles must collide with sufficient energy. Since then, we have learned that the kinetic energy is related to the force of the particles. Therefore the particles have to achieve an energy greater than the threshold energy to have successful collisions. From this knowledge we can conclude that the rate of the reaction depends on the fraction of molecules that have enough energy and that collide with the proper orientation, and the rate depends on the collision frequency itself.
- Putting this all together we get the following:
Where the collision frequency is the first factor of a successful collision, the collision geometric orientation is the second factor of a successful collision, and the collision energy is the third factor of a successful collision.
Lesson Summary
- The collision theory explains why reactions occur between atoms, ions, and/or molecules and allows us to predict what conditions are necessary for a successful reaction to take place.
- The kinetic molecular theory provides the foundation for the collision theory on the molecular level. The minimum amount of energy necessary for a reaction to take place is known as the threshold energy.
- With increasing temperature, the kinetic energy of the particles increase and the number of particles with energy greater than the activation energy increases.
- The total number of collisions per second is known as the collision frequency, whether these collisions are successful or not.
Review Questions
- According to the collision theory, it is not enough for particles to collide in order to have a successful reaction to produce products. Explain
- Due to the number of requirements for a successful collision, according to the collision theory, the percentage of successful collisions is extremely small. Yet, chemical reactions are still observed at room temperature and some at very reasonable rates. Explain.
- What is a basic assumption of the kinetic molecular theory?
- all particles will lose energy as the velocity increases
- all particles will lose energy as the temperature increases
- all particles will increase velocity as the temperature decreases
- all particles are in random motion
- i, ii
- i, iii
- ii, iii
- i, ii, iii
- particles must collide
- particles must have proper geometric orientation
- particles must have collisions with enough energy
- the particles would rebound and there would be no reaction
- the particles would keep bouncing off each other until they eventually react, therefore the rate would be slow
- the particles would still collide but the byproducts would form
- the temperature of the reaction vessel would increase
Further Reading / Supplemental Links
Vocabulary
- collision theory
- Explains why reactions occur at this particle level between atoms, ions, and/or molecules; even more important from the collision theory is the ability to predict what conditions are necessary for a successful reaction to take place.
- kinetic molecular theory
- Provides the foundation for the collision theory on the atomic level; the collisions between particles are considered to elastic in nature.
- threshold energy
- The minimum amount of energy necessary for a reaction to take place.
- collision frequency
- The total number of collisions per second.
Potential Energy Diagrams
Lesson Objectives
- Define enthalpy, activation energy, activated complex.
- Describe and draw the difference between endothermic and exothermic potential energy diagrams.
- Draw and label the parts of a potential energy diagram.
Introduction
In the previous lesson the kinetic molecular theory was discussed in relation to the collision theory. The kinetic molecular theory was used to help understand the effect of increasing the temperature on the number of effective collisions in a reaction.
In this lesson, we go beyond the theoretical perspectives of the collision theory to consider how particle collisions can be represented in energy diagrams. Potential energy diagrams in the study of kinetics show how the potential energy changes during collisions from reactants and products. In this first lesson, we will examine the features of such diagrams.
Internal Energy of Reactants and Products, ΔH
As stated in the introduction, potential energy diagrams illustrate the potential energy of the reactants and products for a chemical reaction and how this energy changes during a reaction. Take a look at the potential energy of the reaction drawn in Figure below.
| Todo 10 | Replace reference to Figure 8 with variable Che-2503-01 |
|---|
The 
-axis represents the potential energy. The potential energy measures the energy stored within the bonds and phases of the reactants and products. This potential energy represents the internal energy of the molecules and, in chemistry, is often called enthalpy. The 
-axis represents the reaction progress. Chemical reactions proceed (or are read) from left to right. Therefore, looking at the potential energy diagram, the reactants are usually found to the left on the diagram and the products on the right.
The enthalpy of a substance is sometimes called heat content. The potential energy stored in the bonds of the substance was thought of as heat stored as potential energy. When a reaction occurs, the enthalpy or heat content of the reactants changes into the enthalpy or heat content of the products. The enthalpy of the reactants and products is almost never the same. Therefore, when a reaction occurs, there is a change in the amount of potential energy stored in the bonds between the reactants and the products. If the bonds of the products store more energy than the bonds of the reactants, then energy must be taken in during the reaction. If the bonds of the products store less potential energy than the bonds of the reactants, then excess potential energy will be left over and will come out of the reaction as kinetic energy. The difference in the enthalpy or heat content of the reactants and that of the products is expressed as 
, or the change in enthalpy. Since this energy is either taken in or given off during the reaction, it is also called the heat of reaction. Notice that in Figure above, the enthalpy change of the reaction is noted by the symbol . The change in enthalpy, , is the difference between the potential energy of the reactants and the potential energy of the products.
Exothermic and Endothermic Potential Energy Diagrams
There are two types of potential energy diagrams. These two types center on the difference between the energies of the reactants and products. Consider the figure below.
Endothermic Reaction (Left); Exothermic Reaction (Right). (Source: Therese Forsythe. CC-BY-SA)
| Todo 10 | Replace reference to Figure 9 with variable Chemistry-2503-02A |
|---|
The definition of is the heat content (enthalpy) of the products minus the heat content (enthalpy) of the reactants, 
. When the enthalpy of the reactants is greater than the enthalpy of the products, heat will be given off by the reaction and because of the way is defined, the will be negative. The opposite is true when the enthalpy of the products is greater than the enthalpy of the reactants.
If the difference between the potential energy of the products and the reactants is positive, or 
, the reaction is considered to be endothermic (kinetic energy is absorbed and becomes potential energy in the bonds) and is represented by Left Figure above. If the difference between the potential energy of the products and the reactants is negative, or 
, the reaction is considered to be exothermic (excess potential energy from the bonds is left over and comes out into the surroundings as kinetic energy) and is represented by Right Figure above.
Activation Energy Barrier
The activation energy in a potential energy diagram represents the minimum amount of energy required to overcome the energy barrier. This energy must be supplied from the collision energy of the reactant molecules. If the molecules do not have sufficient collision energy to provide the activation energy, then the reaction must be heated to increase the kinetic energy of the reactants in order for the reaction to occur. For instance, hydrogen gas and oxygen gas can be kept in the same container at room temperature for long periods of time without reacting. Even though the molecules are colliding, they do not react since there is insufficient activation energy.
In potential energy diagrams, the symbol for activation energy is often designated as 
. Look at the two exothermic reactions whose diagrams are represented in Figures A and B below and notice the activation energy marked in each.
| Todo 10 | Replace reference to Figure 10 with variable Che-2503-03A |
|---|
When a reaction has a low activation energy, like A above, most of the reactant molecules have sufficient kinetic energy to react and the reaction will most likely be rapid (a high reaction rate). When a reaction has a high activation energy, like B above, most of the reactant molecules will NOT have enough energy to react and the reaction will most likely be very slow.
Activated Complex
The activated complex is a transitional state between the reactants and products. Consider what is happening in the reaction. The reactant bonds are breaking and the product bonds are forming during successful collisions. In the intermediate stages, a transitional complex is formed that creates a short-lived, very unstable species that is the intermediate between the reactants and products. The activated complex contains the highest amount of energy of all of the species in the reaction. The position therefore is at the top of the activation energy barrier as is shown in Figure below.
| Todo 10 | Replace reference to Figure 11 with variable Che-2503-05 |
|---|
Sample Question 1: Consider the reaction between 
and 
.
Under certain circumstances, the enthalpy of the reactants is 
, the activation energy is 
, and the enthalpy of reaction is 
. Draw a potential energy diagram properly labeling the following.
(a) the axes
(b) the activation energy
(c) the change in enthalpy
(d) the activated complex
Solution:
(Source: Therese Forsythe. CC-BY-SA)
| Todo 10 | Replace reference to Figure 12 with variable Che-2503-06 |
|---|
Sample Question 2:
From the potential energy diagram above, list the values for
(a) the enthalpy of the reactants
(b) the enthalpy of the products
(c) the threshold energy
(d) the activation energy
(e) the change in enthalpy
Solution:
(a)
(b) 
(c) 
(d) 
(e) 
(Source: Therese Forsythe. CC-BY-SA)
| Todo 10 | Replace reference to Figure 13 with variable Che-2503-FI |
|---|
Lesson Summary
- Potential energy diagrams in the study of kinetics show how the potential energy changes during reactions from reactants and products. The potential energy measures the energy stored within the bonds (both intermolecular and intramolecular) of the reactants and products, therefore is the internal energy.
- Exothermic reactions have a potential energy difference between the products and reactants that is negative. Endothermic reactions have a potential energy difference between the products and reactants that is positive.
- The threshold energy in a potential energy diagram represents the minimum amount of potential energy necessary for a successful collision to take place. The activation energy is the difference between the energy of the reactants and the threshold energy. The activated complex is a transitional state between the reactants and products.
Review Questions
- Define and explain the importance of the activation energy.
Use the diagram below to answer questions 2 through 6.
- Which letter represents the activation energy barrier?
- a
- b
- c
- d
- The reaction is exothermic in the forward reaction.
- The reaction is endothermic in the forward reaction.
- The reaction is exothermic in the reverse reaction.
- The reaction is exothermic only at high temperatures.
- b
- c
- d
- e
- a
- b
- c
- d
- a transitional species that can eventually be isolated
- a transitional species of that must be made before the products can be formed
- a reactant molecule breaking into a product molecule
- part of the activation energy barrier
.
Draw a potential energy diagram properly labeling the following.
- The axes
- The reactants and products
- The activation energy
- The enthalpy
Vocabulary
- potential energy diagrams
- Potential energy diagrams in the study of kinetics show how the potential energy changes during reactions from reactants to products.
- potential energy
- The potential energy measures the energy stored within the bonds of the reactants and products and therefore is the internal energy.
- exothermic reactions
- Reactions that have a potential energy difference between the products and reactants that is negative.
- endothermic reactions
- Reactions that have a potential energy difference between the products and reactants that is positive.
- activation energy
- The minimum amount of energy that needs to be supplied to the system so that a reaction can occur.
- activated complex
- A high energy transitional state between the reactants and products.
Factors That Affect Reaction Rates
Lesson Objectives
- State how the rate of reaction changes as a function of temperature.
- Explain how increased temperature increases the number of particles that can overcome the energy barrier.
- Provide examples of the temperature dependence on the rate in society.
- Describe the effect of increasing the concentration on the rate of a reaction.
- Indicate which reactants in a multi-step process can affect the rate of a reaction.
- Calculate, using experimental data, the relationship between the ratio of the change in concentration of reactants and ratio of the change in rate.
- Describe the surface area to volume ratio.
- Describe the effect of surface area on reaction rate.
- Describe how the change in the surface area affects the collision frequency.
- Describe real world examples of the effect of surface area on reaction rate.
- Define a catalyst.
- Identify a catalyst in a single equation.
- Identify a catalyst in a multi-step process.
- Describe how a catalyst affects the potential energy diagram.
- Explain how a catalyst affects the rate of the reaction.
- Explain how a catalyst affects our everyday lives, particularly with vitamins.
Introduction
Chemists use reactions to generate a product for which they have use. In some cases, the desired product is simply the energy that is released in an exothermic reaction. The energy released is used to heat homes, generate electric power, and so forth. In other reactions, the desired product is some specific compound for which the chemist has some use. For the most part, the reactions that produce some desired compound are only useful if the reaction occurs at a reasonable rate. For example, using a reaction to produce brake fluid would not be useful if the reaction required 8,000 years complete the product. Such a reaction would also not be useful if the reaction was so fast that it was explosive. For these reasons, chemists wish to be able to control reaction rates. In some cases, chemists wish to speed up reactions that are too slow and slow down reactions that are too fast. In order to gain any control over reaction rates, we must know the factors that affect reaction rates.
Chemists have identified many factors that affect the rate of a reaction. Some of these factors can be altered or controlled by chemists and some cannot. The factors that affect reaction rates have been placed in five categories. The first category contains all those factors that cannot be altered by chemists. The other four factors will be listed and discussed in detail in this section.
Factors Affecting Reaction Rates
- The Nature of the Reactants
- The Temperature of the Reactants
- The Concentration of the Reactants
- Surface Area
- The Presence of a Catalyst
The Nature of the Reactants
There are a number of factors that affect the rate of reaction that cannot be altered. That is, once the reactants are chosen, these factors are permanently determined and cannot be changed. Since chemists can't do anything about these factors, they are lumped together in a single category. We will discuss them here and then not mention them in the remainder of the section. The focus of the remainder of this section will be on those factors affecting reaction rates that chemists can alter and thereby exert some control over reaction rates.
So, what are some of the factors affecting reaction rates that cannot be altered. You may remember from an earlier chapter, the discussion of bond strength. When a bond forms, energy is given off. If a large amount of energy is given off during bond formation, then the same large amount of energy will be required to break the bond. Bonds that give off great amounts of energy when the bond is formed are called strong bonds. Bonds that give off only a small amount of energy when they are formed are called weak bonds and are easily broken - requiring only the same small input of energy to break the bond. If the molecules in the reactants have strong bonds, then a large activation energy will be required under room conditions and the reaction rate will most likely be very slow. Since the bond strengths in the reactants cannot be changed, this factor affecting reaction rates is not alterable and therefore, it is placed in the "Nature of the Reactants" category.
Consider the number of bonds that need to be broken in the reactants. If a reaction involves combining silver ions and chloride ions in solution, you might note that there are zero bonds to be broken. In a solution, silver ions and chloride ions have no bonds to break. All that has to occur for these ions to react is that the ions have to come in contact. A reaction between such ions in solution will be almost instantaneous. Compare this to a reaction involving a reactant molecule that has a hundred bonds that have to be broken. Look at all the bonds that have to broken if the reactants for a reaction are 
. These reactants have so many bonds to break that there is no way a single collision could break them all. There are so many bonds to break that this reaction will have to occur with a long series of successive collisions. Such a reaction will always be slow under normal conditions. Once again, the number of bonds to be broken is dependent on the reactants involved in the reaction and cannot be changed. This factor, also, must go into the "Nature of the Reactants" category.
Let's consider one more factor affecting reaction rates that cannot be altered. Suppose you have two reactants that are essentially spheres such as copper atoms and silver ions. These two reactants can collide on any side from any direction and produce a successful collision. Compare that to a reactant whose molecule is shaped like a 
and the only successful collision must be inside the opening of the . For such a reactant, a very special collision geometry is required and very few of the collisions will have this particular geometry. In such cases, again, the reaction will be slow and there is no way to change the shape of the reactant. Reaction rates controlled by special collision geometries also go into the category "The Nature of the Reactants".
Effect of Temperature on Rate of Reaction
Increased Temperature
The rate of reaction was discussed in terms of three factors: collision frequency, the collision energy, and the geometric orientation. Remember that the collision frequency is the number of collisions per second. The collision frequency is dependent, among other factors, on the temperature of the reaction.
When the temperature is increased, the average velocity of the particles is increased. The average kinetic energy of these particles is also increased. The result is that the particles will collide more frequently because the particles move around faster and will encounter more reactant particles but this is only a minor part of the reason why the rate is increased. Just because the particles are colliding more frequently does not mean that the reaction will definitely occur.
The major effect of increasing the temperature is that more of the particles that collide will have the amount of energy needed to have an effective collision. In other words, more particles will have the activation energy needed to overcome the activation energy barrier and form the activated complex. The effect of raising the temperature, therefore, is to produce more activated complexes and, with the greater number of activated complexes that are formed, the faster the rate of reaction.
At room temperature, the hydrogen and oxygen in the atmosphere do not have sufficient energy to attain the activation energy needed to produce water.
At any one moment in the atmosphere, there are many collisions occurring between these two reactants. And, when this reaction does occur it is exothermic which tends to mean that the reaction should be happening. But what we find is that water is not formed from the oxygen and hydrogen molecules colliding in the atmosphere because the activation energy barrier is just too high and all the collisions are resulting in rebound. When the necessary energy is supplied to the molecules, the molecules overcome the activation energy barrier, the activated complex is formed, and water is produced:
Decreased Temperature
There are times when the rate of a reaction needs to be slowed down. Using the factors as specified previously, one of ways to accomplish this would be to keep the reactants in separate containers so that there can be no collisions between the particles. At times that might not be practical so lowering the temperature could also be used to decrease the number of collisions that would occur and lowering the temperature would also reduce the kinetic energy available for activation energy. If the particles have insufficient activation energy, the collisions will result in rebound rather than reaction. Using this idea, when the rate of a reaction needs to be lower, keeping the particles from having sufficient activation energy will definitely keep the reaction at a lower rate.
A Generalization for Increased Temperature
The rate of most reactions can be dramatically increased with increased temperature. For reactions that normally occur at room temperature, a general “rule of thumb” is that for every increase of 
, the rate will be doubled. And if the temperature for these reactions is increased by 
, the rate will be increased by a factor of 
; increasing the temperature by 
the rate will be increased by a factor to 
. However, for any specific reaction, the actual rate increase will have to be determined by experimentation.
Examples of Temperature on Reaction Rate
Society uses the effect of temperature on rate every day. Food storage is a prime example of how the temperature effect on reaction rate is utilized by society. Consumers store food in freezers and refrigerators to slow down the processes that cause it to spoil. The decrease in temperature decreases the rate at which the food will break down or be broken down by bacteria.
In the early years of the 
century, explorers were fascinated with trying to be the first one to reach the South Pole. In order to attempt such a difficult task at a time without most of the technology we take for granted today, they devised a variety of ways of surviving. One method was to store their food in the snow to be used later during their advances to the pole. On some explorations, they buried so much food, that they didn’t need to use all of it and it was left. Many years later, when this food was located and thawed, it was found to still be edible.
When milk, for instance, is stored in the refrigerator, the molecules in the milk have less energy. This means that while molecules will still collide with other molecules, few of them will react (which means in this case “spoil”) because the molecules do not have sufficient energy to overcome the activation energy barrier. The molecules do have energy and are colliding, however, and so, over time, over time, even in the refrigerator, the milk will spoil. Eventually the higher energy molecules will gain the energy needed to make the activated complex and when enough of these reactions occur, the milk becomes “soured”.
However, if that same carton of milk was at room temperature, the milk would react (in other words “spoil’) much more quickly. Now most of the molecules will have sufficient energy to overcome the energy barrier and at room temperature many more collisions will be occurring. This allows for the milk to spoil in a fairly short amount of time. This is also the reason why most fruits and vegetables ripen in the summer when the temperature is much warmer. You may have experienced this first hand if you have ever bitten into an unripe banana – it was probably sour tasting and might even have felt like biting into a piece of wood! When a banana ripens, numerous reactions occur that produce all the compounds that we expect to taste in a banana. But this can only happen if the temperature is high enough to allow these reactions to make those products.
Lesson Summary
- With an increase in temperature, there is an increase in the amount of kinetic energy that can be converted into activation energy in a collision and therefore there will be an increase in the reaction rate. A decrease in temperature would have the opposite effect.
- With an increase in temperature there is an increase in the number of collisions which is a minor factor in increasing the rate of reaction. The major reason reaction rate increases with increasing temperature is an increase in temperature increases the number of particles that have sufficient activation energy to overcome the activation barrier and form the activated complex.
- A rule of thumb used for the effect of temperature on the rate is that if the temperature is increased by , the rate is doubled.
Review Questions
- Why does an increase in temperature increase the rate of the reaction?
- Why does higher temperature increase the reaction rate?
- more of the reacting molecules will have higher kinetic energy
- increasing the temperature causes the reactant molecules to heat up
- the activation energy will decrease
- increasing the temperature causes the potential energy to decrease
- number of collisions
- activation energy requirement
- number of successful collisions
- all of the above change
. The rate of reaction for a hypothetical reaction was found to be 
at 
. What would be the rate at 
?Vocabulary
- effective collision
- A collision that results in a reaction.
Effect of Concentration
Increasing Concentration
If you had an enclosed space, like a classroom, and there was one red ball and one green ball flying around the room with random motion and undergoing perfectly elastic collisions with the walls and with each other, in a given amount of time, the balls would collide with each other a certain number of times determined by probability. If you now put two red balls and one green ball in the room under the same conditions, the probability of a collision between a red ball and the green ball would exactly double. The green ball would have twice the chance of encountering a red ball in the same amount of time. In terms of chemical reactions, a similar situation exists. Particles of two gaseous reactants or two reactants in solution have a certain probability of undergoing collisions with each other in a reaction vessel. If you double the concentration of either reactant, the probability of a collision doubles. The rate of reaction is proportional to the number of collisions per unit time. If one concentration of one of the reactants is doubled, the number of collisions will also double. Assuming that the percent of collision that are successful does not change, then having twice as many collisions will result in twice as many successful collisions. The rate of reaction is proportional to the number of collisions per unit time and increasing the concentration of either reactant increases the number of collisions and therefore, increases the number of successful collisions and the reaction rate.
Some reactions occur by a single collision between two reactant molecules while other reactions occur by a series of collisions between multiple reactant particles. We will consider the case of the single step now and the case of the multiple step reactions later in the unit.
The rate of a single collision chemical reaction at a given temperature can be expressed as a product of the concentrations of the reactants. For the reaction, 
, the reaction rate can be expressed as:
, where 
is a constant for the reaction called the reaction constant and 
and 
are the molarities of the reactants.
If only the concentration of 
is doubled, the equation would become,
and the rate would obviously be double the previous rate.
If only the concentration of 
is doubled, the equation would become,
and once again the rate would obviously be double the previous rate.
If the concentrations of both and are doubled, the equation becomes
and now the rate would be 
the original rate.
Experimental Determination of Reaction Rate
When the reaction involves a series of collisions, the relationship between the reaction rate and the concentration of any single reaction can only be determined by a laboratory procedure. Consider the reaction below between and 
. This reaction does not occur by a single collision but rather in a two step process.
The effect of the concentration of a reactant on the rate of this multiple step reaction can only be known through experimentation. Let’s look at one experiment where it can be determined how the concentration of the reactants affects the rate of the reaction.
Sample question: For the hypothetical reaction 
, the following data was collected in an experiment to attempt to determine the effect of increasing the concentration of the reactants on the rate.
| Trial | 
| 
| 
|
|---|---|---|---|
| 1 | 
|
| 
|
| 2 | 
|
|
|
| 3 |
|
| 
|
| Todo 10 | Table need caption. Table needs ID. Table needs title. References to table should be updated with ID. |
|---|
Determine the effect of increasing the and increasing on the rate of the reaction.
Solution:
Step 1: Try to find two trials where the concentration of is changing while the concentration of remains the same.
In this case, in trials 
and 
, changes while remains constant. Since is changing and is staying the same, any change in the reaction rate is due to the change in .
Step 2: Determine the effect of changing the concentration of on the reaction rate.
From trial 1 to trial 2, the concentration of has doubled. The reaction rate in these two trials, however, did not change.
Therefore, the concentration of has no effect on the rate for this reaction.
Step 3: Try to find two trials in which changes while stays the same.
In trials 1 and 3, the concentration of changes while stays the same. From trial 1 to trial 3, the concentration of doubles. The reaction rate also doubled between these two trials. When the concentration of doubled, the reaction rate also doubled. Therefore, we can conclude that the reaction rate is directly proportional to the concentration of .
Therefore, when the concentration of is increased, there is no effect on the rate of the reaction.
Therefore, when the concentration of is doubled, the rate is doubled. In other words, increasing the concentration of increases the rate.
Example of the Effect of Concentration on Reaction Rate
The chemical test used to identify a gas as oxygen or not relies on the fact that increasing the concentration of a reactant increases reaction rate. The reaction we call combustion refers to a reaction in which a flammable substance reacts with oxygen. If we light a wooden splint (a thin splinter of wood) on fire and then blow the fire out, the splint will continue to glow in air for a period of time. If we insert that glowing splint into any gas that does not contain oxygen, the splint will immediately cease to glow - that is the reaction stops. Oxygen is the only gas that will support combustion. Air is approximately, 
oxygen gas. If we take that glowing splint and insert it into pure oxygen gas, the reaction will increase its rate by a factor of five - since pure oxygen has 
the concentration of oxygen that is in air. When the reaction occurring on the glowing splint increases its rate by a factor of five, the glowing splint will suddenly burst back into full flame. This test, of thrusting a glowing splint into a gas, is used to identify the gas as oxygen. Only a greater concentration of oxygen than that found in air will cause the glowing splint to burst into flame.
Lesson Summary
- Increasing the concentration of a reactant increases the frequency of collisions between reactants and will, therefore, increase the reaction rate.
Review Questions
- Explain how concentration affects reaction rate using the collision theory. You may want to include a diagram to help illustrate your explanation.
- Why is the increase in concentration directly proportional to the rate of the reaction?
- The kinetic energy increases.
- The activation energy increases.
- The number of successful collisions increases.
- All of the above.
is doubled, the rate of reaction stays the same. If the concentration of 
doubles, the rate of the reaction quadruples.
What is the explanation for this observation?
- The reaction is nearing completion and all is used up.
- The reaction occurs in more than one step.
- Excess has been added.
- Not enough information is given.
Which of the following would have the greatest effect on the rate of reaction?
- Increase
- Increase
- Increase
- Increase and
- A, B, C
- A
- B
- C
in 
at 
.
Determine the effect of decreasing the 
on the rate of the reaction.
| Trial | 
| 
|
|---|---|---|
| 1 | 
| 
|
| 2 | 
| 
|
| 3 | 
| 
|
| 4 | 
| 
|
| 5 | 
| 
|
| Todo 10 | Table need caption. Table needs ID. Table needs title. References to table should be updated with ID. |
|---|
Further Reading / Supplemental Links
Vocabulary
- multi-step process
- Reactions that take more than one step in order to make the products.
Effect of Surface Area
The Relationship Between Surface and Reaction Rate
The very first requirement for a reaction to occur between reactant particles is that the particles must collide with each other. The previous section pointed out how increasing the concentration of the reactants increases reaction rate because it increased the frequency of collisions between reactant particles. It can be shown that the number of collisions that occur between reactant particles is also dependent on the surface area of reactants. Consider a reaction between reactant RED and reactant BLUE in which reactant BLUE is in the form of a single lump (Figure A below). Then compare this to the same reaction where reactant BLUE has been broken up into many smaller pieces (Figure B below).
Figure When the large piece of blue molecules is broken into smaller pieces, red molecules can collide with many more blue molecules.[1]
| Todo 10 | Replace reference to Figure 14 with variable Chem-21-04-SA |
|---|
In Figure A above, only the BLUE particles on the outside surface of the lump are available for collision with reactant RED. The BLUE particles on the interior of the lump are protected by the BLUE particles on the surface. In Figure A above, if you count the number of BLUE particles available for collision, you will find that only 20 BLUE particles could be struck by a particle of reactant RED. In Figure A above, there are a number of BLUE particles on the interior of the lump that cannot be struck. In Figure B above, however, the lump has been broken up into smaller pieces and all the interior BLUE particles are now on a surface and available for collision. In diagram B, more collisions between BLUE and RED will occur, and therefore, the reaction in Figure B above will occur at faster rate than the same reaction in Figure A above. Increasing the surface area of a reactant increases the frequency of collisions and increases the reaction rate.
The more surface area that is available for particles to react the faster the reaction will occur. You can see an example of this in everyday life if you have ever tried to start a fire in the fireplace. If you hold a match up against a large log in an attempt to start the log burning, you will find it to be an unsuccessful effort. Flammable material like wood requires a significant input of activation energy for the reaction to occur. The reaction between wood and oxygen is an exothermic reaction and once the fire has been started, the heat released by the first reactions to occur will provide the activation energy for the succeeding reactions. However, holding a match against a large log will not cause enough reactions to occur in order to keep the fire going by providing sufficient activation energy for further reactions. In order to start a wood fire, it is common to break a log up into many small, thin sticks called kindling. These thinner sticks of wood provide many times the surface area of a single log. The match will successfully cause enough reactions in the kindling so that sufficient heat is given off to provide activation energy for further reactions.
There have been, unfortunately, cases where serious accidents were caused by the failure to understand the relationship between surface area and reaction rate. One such example occurred in flour mills. A grain of wheat is not very flammable. It takes a significant effort to get a grain of wheat to burn. If the grain of wheat, however, is pulverized and scattered through the air, only a spark is necessary to cause an explosion. When the wheat is ground to make flour, it is pulverized into a fine powder and some of the powder gets scattered around in the air. A small spark then, is sufficient to start a very rapid reaction which can destroy the entire flour mill. In a 10-year period from 1988 to 1998, there were 
grain dust explosions in mills in the United States. Efforts are now made in flour mills to have huge fans circulate the air in the mill through filters to remove the majority of the flour dust particles. Another example is in the operation of coal mines. Coal, of course, will burn but it takes an effort to get the coal started and once it is burning, it burns slowly because only the surface particles are available to collide with oxygen particles. The interior particles of coal have to wait until the outer surface of the coal lump burns off before they can collide with oxygen. In coal mines, huge blocks of coal must be broken up before the coal can be brought out of the mine. In the process of breaking up the huge blocks of coal, drills are used to drill into the walls of coal. This drilling produces fine coal dust that mixes into the air and then a spark from a tool can cause a massive explosion in the mine. There are explosions in coal mines for other reasons but coal dust explosions contributed to the death of many miners. In modern coal mines, lawn sprinklers are used to spray water through the air in the mine and this reduces the coal dust in the air and eliminates coal dust explosions.
Increasing Surface Area Increases Frequency of Collision
An increase in the surface area allows for an increase in the frequency of collisions between reacting molecules which will increase the rate of reaction. Look at the diagram of a hypothetical tablet below that is going to be used to react with hydrochloric acid (Right Figure below). When the tablet is crushed before the reaction begins, look at the difference in the number of reaction sites at which the hydrochloric acid can react (Left Figure below).
| Todo 10 | Replace reference to Figure 15 with variable Chemistry-2506-01A |
|---|
An increase in the surface area increases the frequency in collisions and therefore the rate of the reaction is increased.
Examples
A way to study this factor is the following: take two solids, put them together and observe the reaction; then, put one of these solids into solution, add the other solid and observe the reaction. For example: If you were to take a few grams of copper (II) chloride and place them into a beaker along with a piece of aluminum foil, it would take a numbers days if not weeks before you would observe more than just a few changes. However if you were to make a solution of the copper (II) chloride and then add the aluminum foil, you would observe an almost immediate reaction. In this case, the surface area of only one of the reactants was changed, but that change would dramatically affect the rare of reaction because the copper (II) chloride ions would now be totally separated and could individually interact with the atoms present in the aluminum foil.
Lesson Summary
- Increasing the surface area of a reactant increases the number of particles available for collision and will increase the number of collisions between reactants per unit time.
- Increasing the frequency of collisions increases the reaction rate.
Review Questions
- Why, using the collision theory, do reactions with higher surface are have a faster reaction rates?
- When does an increase in surface area not increase the rate of reaction?
- The rate will not be increased if there is insufficient activation energy present.
- The rate will not increase if there is not an increase in collisions.
- The rate will not increase if the concentration doesn’t change.
- The rate will not increase if does not increase.
- a block of ice or crushed ice
- sugar cubes or sugar crystals
- a piece of wood or wood shavings
-
or 
-
or 
- the purpose of the lighter fluid; and,
- the purpose of the coals.
Vocabulary
- surface area to volume ratio
- The comparison of the volume inside a solid to the area exposed on the surface.
Effect of a Catalyst
Introduction
The final factor that affects the rate of the reaction is the effect of the catalyst. A catalyst is a substance that speeds up the rate of the reaction without itself being consumed by the reaction.
There are a number of catalysts that we recognize. There are surface catalysts which merely provide a surface for intermediate products to adhere to and there are catalysts that are used at the beginning of a reaction but are completely reproduced at the end. The substances called enzymes in biology are catalysts that help carry out numerous chemical reactions in the body. Many commercial preparations of chemicals for industry rely on catalysts to prepare their products more efficiently and cost effectively. For example in the production of sulfuric acid, iron oxide or vanadium oxide is used in combination with platinum as surface catalysts to produce the desired 
.
In the reaction of potassium chlorate breaking down to potassium chloride and oxygen, a catalyst is available to make this reaction occur much faster than it would occur by itself under room conditions. The catalyst is manganese dioxide and its presence causes the reaction shown above to run many times faster than it occurs without the catalyst. When the reaction has reached completion, the 
can be removed from the reaction vessel and its condition is exactly the same as it was before the reaction. This is part of the definition of a catalyst . . . that it is not consumed by the reaction.
We already know that we can increase the rate of a reaction by adding one of the reactants to increase the concentration of that reactant and thus increase the reaction rate. But, that added reactant will be consumed by the reaction. A catalyst is a substance that can be added and will increase the reaction rate but will not be consumed by the reaction.
In this lesson, the effect of the catalyst on the reaction rate will be considered as well as its effect on the potential energy diagrams. Let’s begin.
A Catalyst is Not a Reactant
As stated in the introduction, a catalyst is defined as a substance that speeds up the rate of the reaction but is itself not consumed by the reaction. In other words, the catalyst is not seen in the reaction as either a reactant or a product. Consider Equation 1 again.
Equation 1 is a typical laboratory experiment where potassium chlorate is heated to produce potassium chloride and oxygen gas. The reaction is very slow unless you add manganese dioxide as a catalyst. Manganese dioxide is a black powder; potassium chlorate is a white powder. After heating the potassium chlorate and obtaining the oxygen gas at the end of the reaction, all of the black can be recovered. You should note that the catalyst is not written into the equation as a reactant or product but is noted above the yields arrow. This is standard notation for the use of a catalyst.
Look at the following three-step process below:
In this three-step process, all of the reactions are added together and all substances that appear on both sides of the equation are eliminated before writing the final overall equation. Notice how 
is consumed in the first equation of the sequence and then produced in the final equation of the sequence. Since is consumed and then produced, it is a catalyst. The presence of the water molecule causes this reaction to occur at a higher rate than it will occur without the presence of water. Therefore, the water is introduced into the reaction and is used in the reaction but it is reproduced at the end of the reaction and so the total amount of it is available at the end. This is the behavior of a catalyst.
Catalysts Provide a Different Path with Lower Activation Energy
Some reactions occur very slowly without the presence of a catalyst. In other words the activation energy for these reactions is very high. When the catalyst is added, the activation energy is lowered because the catalyst provides a new reaction pathway with lower activation energy.
| Todo 10 | Replace reference to Figure 16 with variable che-2507-03A |
|---|
Remember that the catalyst does not get consumed itself in the reaction so the reactants and product positions are not affected by the addition of the catalyst. In Left Figure above, the endothermic reaction shows the catalyst reaction in red with the lower activation energy, designated 
. The new reaction pathway has lower activation energy but has no effect on the energy of the reactants, the products, or the value of .
The same is true for the exothermic reaction in Right Figure above. The activation energy of the catalyzed reaction (again designated by ) is lower than that of the uncatalyzed reaction. The new reaction pathway provided by the catalyst for the exothermic reaction in Right Figure above affects the energy required for reactant bonds to break and product bonds to form.
Surface Catalysts and Enzymes
An example of a surface catalyst is the behavior of a platinum catalyst in the reaction in which hydrogen and oxygen form water.
This reaction, as you are probably aware is so slow under room conditions that it essentially doesn't occur. That's why we have hydrogen gas and oxygen gas in our atmosphere without this reaction occurring. Even if hydrogen and oxygen gases are mixed in a reaction vessel, no reaction occurs. This reaction has a very high activation energy requirement and room conditions simply do not provide sufficient energy for the reaction to occur.
(Source: Therese Forsythe. CC-BY-SA)
The uncatalyzed reaction between hydrogen gas and oxygen gas requires all the reactant particles to collide in a single collision with enough activation energy to break the bond in both the hydrogen molecules AND break the double bond in the oxygen molecule. Under room conditions, these particles are not nearly energetic enough for a collision to provide that much activation energy and therefore, no reaction occurs.
(Source: Richard Parsons. CC-BY-SA)
If a platinum surface is available, the oxygen molecules can strike the platinum surface and break the bond holding the oxygen atoms together. The oxygen atoms then adhere to the surface of the platinum. This collision requires less energy than the giant collision necessary when platinum is not present. Later, a hydrogen molecule can collide with one of the oxygen atoms adhered to the platinum surface and this collision breaks the bond in the hydrogen molecule and then the hydrogen and oxygen can combine and leave the surface of the platinum. Eventually, another hydrogen molecule repeats this process with the final oxygen atom. When all the oxygen atoms have left the surface of the platinum, the platinum is exactly the same as it was before the reaction. In this way, the reaction has occurred with several smaller collisions rather than one large one. Room conditions are sufficient to provide the activation energy necessary for the small collisions but not the large one. Therefore, in the presence of a platinum catalyst this reaction will occur at room temperature. The reaction rate has been significantly increased. The fact that the platinum increases the reaction rate but is not permanently consumed, qualifies it as a catalyst in this reaction. As evidence of the success of the platinum as a catalyst for this reaction, if a piece of platinum is dropped into a small container of hydrogen and oxygen gas at room conditions, a small explosion occurs as this reaction goes to completion almost immediately.
The potential energy diagram for the catalyzed reaction has some parts that have changed and some parts that have not changed. The reactants and products are exactly the same as before because it is the same reaction. Since the products and reactants are the same, they have the same enthalpy stored in their bonds and therefore, the will be exactly the same for both reactions. The area that has changed is the energy barrier. Each of collisions has a lower activation energy requirement so the energy barrier is lower in the catalyzed reaction than it was in the uncatalyzed reaction. The reaction mechanism for catalyzed reaction is different - the reaction does not occur by the same process. The catalyst provides a different reaction path for the same reaction and the new path has a lower activation energy requirement. The lower activation energy allows for a much faster reaction rate.
While many reactions in the laboratory can be increased by increasing the temperature, that is not possible for all the reactions that occur in our bodies throughout our entire lives. In fact, the body needs to be maintained at a very specific temperature: 
or 
. Of course there are times, for instance, when the body is fighting an infection, when the body temperature may be increased. But generally, in a healthy person, the temperature is quite consistent. However, many of the reactions that a healthy body depends on could never occur at body temperature. The answer to this dilemma is catalysts or what are also referred to as enzymes. Many of these enzymes are made in your cells since your DNA carries the directions to make them. However, there are a some enzymes that your body must have but are not made in your cells. These catalysts must be supplied to your body in the food you eat and are called vitamins.
One example of these catalysts is the water-soluble vitamins. These vitamins include 
(Thiamin), 
(Riboflavin), 
(Niacin), 
(Pyridoxine), 
(Cyanocobalamin), Biotin, Folacin, Pantothenic Acid, and Vitamin . Table below shows the structure, the use of the different vitamins and where they are found commonly in foods. The role of the vitamin in the body is as a catalyst in metabolism. Remember that catalysts are substances that are not consumed by a reaction. They are used in order to speed up the reaction. The vitamin, itself, does not affect how much of the particular amino acid, protein, or other product that is produced in the reaction to which it is involved. It is important to remember that while these catalysts are not used up in a reaction, a person nonetheless needs to continue to take in vitamins since the body only uses them for a limited amount of time and then they are discarded.
| Name/Found in | Structure | Important Functions |
|---|---|---|
| (Thiamin)
Found in: Peanuts, pork, bran, egg yolk, beans are some examples. | 
|
|
| (Riboflavin)
| 
|
Found in: Organ meats, nuts, cheese, eggs, milk leafy green vegetables, fish, whole grains, and yogurt are some examples. |
| (Niacin)
Found in: Liver, poultry, fish, nuts, cereals, asparagus, seeds, milk, and leafy green vegetables, are some examples. | 
|
|
| (Pyridoxine)
Found in: Brewer's yeast, eggs, chicken, carrots, fish, liver, peas, walnuts are some examples. | 
|
|
| (Cyanocobalamin)
Found in: | 
|
|
Biotin – Vitamin 
Found in: Cheese, beef liver, cauliflower, eggs, mushrooms, chicken breasts, salmon, spinach, brewer's yeast, and nuts are some examples. Vitamin | 
|
|
Folacin (Folic Acid) – 
Found in: Fresh green vegetables (broccoli, spinach), fruit, beans, whole grains and liver are some examples. | 
|
|
Pantothenic Acid – 
Found in: Beef, brewer’s yeast, eggs, fresh vegetables, kidney, legumes, liver, mushrooms, nuts, pork, saltwater fish, and whole wheat are some examples. | 
|
|
| Vitamin C
Found in: Leafy green vegetables, berries, citrus fruits, tomatoes, melons, papayas are good examples. | 
|
|
You can see from Table above, these water-soluble vitamins act on a large number of the systems in our bodies. One example is Vitamin . It is only needed to be added to the diet of humans and several kinds of animals.
It is interesting how some of the vitamins work in the body. For some vitamins, their role is to carry the chemical groups between enzymes; therefore they are given the term coenzymes and act as the precursor for enzymes. Their function is essential to the development and/or maintenance of systems within our bodies.
If you look at folic acid as an example, it is responsible for assisting in DNA synthesis. Folic acid has a structure shown below.
An example of Folic Acid.
Methyl-tetrahydrofolate is the chemical name for folic acid. Folic acid has the job of supplying thymidine triphosphate (see Figure below).
In the synthesis of DNA, there are four nucleotide bases required. Thymidine triphosphate is one of these bases and is so important that if it is not present, the DNA synthesis stops. Folic acid (or folacin) supplies this nucleotide base. Neural tube defects such as, spina bifida, are the direct result of a lack of folic acid, specifically thymidine triphosphate. A low amount of this nucleotide base causes the neural tube in the brain to improperly close or not close at all causing serious defects. The lower amount of the base, the more serious the defect.
Observing molecules during chemical reactions helps explain the role of catalysts. Dynamic equilibrium is also demonstrated. Molecules in Action (http://www.learner.org/vod/vod_window.html?pid=806)
Surface science examines how surfaces react with each other at the molecular level. On the Surface (http://www.learner.org/vod/vod_window.html?pid=812)
Lesson Summary
- The catalyst is a substance that speeds up the rate of the reaction without itself being consumed by the reaction. When the catalyst is added, the activation energy is lowered because the catalyst provides a new reaction pathway with lower activation energy.
- The new reaction pathway has lower activation energy but has no effect on the energy of the reactants, the products, or the value of ΔH. Water-soluble vitamins are a common example of a catalyst acting as coenzymes.
Review Questions
- Draw a potential energy diagram for an exothermic reaction labeling the following.
- The activation energy of
- The enthalpy of
- The reactants and product
- The axes
- The activation energy for the catalyzed reaction.
- The activation energy of
- provide an alternate reaction pathway
- change the kinetic energy of the reacting particles
- eliminate the slow step
- add another reactant
- the activation energy of the forward reaction is lowered
- the activation energy of the reverse reaction is lowered
- the activation energy in general is lowered
- the enthalpy of the reaction is lowered
Further Reading / Supplemental Links
The learner.org website allows users to view streaming videos of the Annenberg series of chemistry videos. You are required to register before you can watch the videos but there is no charge. The website has one video that relates to this lesson called Molecules in Action.
Vocabulary
- catalyst
- A substance that speeds up the rate of the reaction without itself being consumed by the reaction.
Reaction Mechanism
Lesson Objectives
- Define reaction mechanisms.
- Identify the rate-determining step.
- Draw a potential energy diagram for a multi-step process.
Introduction
In the last section, the reaction between an uncatalyzed reaction between hydrogen and oxygen was compared to a catalyzed reaction between the same two reactants. It was pointed out that the uncatalyzed reaction required three particles to collide at the same time with sufficient energy to break all the bonds in all the molecules. That uncatalyzed reaction had a very high activation energy and therefore did not occur under room conditions. The catalyzed reaction could occur because it allowed for a series of collisions so the bonds could be broken one at a time, thus requiring less activation energy.
Consider a reaction in which a molecule containing 
is going to reacted with 
oxygen molecules. What do you think is the possibility of 
randomly moving particles to all arrive at exactly the same spot in space so they can have one big collision? You are right, that is not going to happen. And, if it did happen, what are the chances that the collision would be hard enough to break all the bonds in the large reactant and all of the double bonds in the oxygen molecules? Once again, the answer is that the chances are approximately ZERO. Complicated reactions involving many bonds and many molecules do not occur in single collisions. These reactions, instead, occur in a series of collisions. Each collision in the series produces an intermediate product that undergoes further collisions until finally, after many collisions, the final products are produced. This long series of collisions producing intermediate products is called a reaction mechanism.
The reaction below is an example of a multi-step reaction. Nitrogen dioxide and carbon monoxide reactant is a two-step process to form nitrogen monoxide and carbon dioxide. The two steps are shown below.
In this lesson, we will discuss these multi-step processes, called reaction mechanisms, as well as the individual reactions in the multi-step process.
Most Reactions Have Multi-Steps
As stated in the introduction, when there is a reaction involving more than two reactant particles the likelihood of a successful single collision reaction is small. This is due to the fact that the particles have to have a successful collision with enough energy and the proper orientation. What will happen in these cases is that the overall reaction will take place in a series of single steps, often called elementary steps. An elementary step is a single, simple step in a multi-step process. An elementary step almost always involves only two particles. The series of elementary steps outline the process of the reaction. Most reactions do not take place in one step but rather occur as a combination of two or more elementary steps. This series of steps is referred to as a reaction mechanism. Even some two-particle collisions require a reaction mechanism.
In another lesson, it was discussed that the concentration of some reactants can affect the rate of the reaction. The rate of the reaction is dependant on the reactants in the slowest step of the multi-step process. If we look at the reaction below from this same lesson, this is only a two particle collision, and 
, and yet it is known that the rate is only affected by the concentration of . This indicates that the reaction proceeds by way of a reaction mechanism.
Sample question: Which of the following reactions would most likely involve a reaction mechanism? Explain.
(a) 
(b) 
Solution:
(a) reaction mechanism; three reactant particles are present
(b) elementary step; only one reactant particle is present
Each Step Has Its Own Activated Complex
When there is a single step reaction mechanism, we can draw potential energy diagrams as we have seen earlier in this chapter. However, when there is a multi-step process where two or more elementary steps combine to form one reaction mechanism, the potential energy diagram looks quite different. Look at the reaction mechanism below. This mechanism is involved in the depletion of the ozone layer.
The overall reaction is 
If we were to draw the potential energy diagram for this two-step process, it could be represented in Figure below. Notice that for each reaction in the reaction mechanism, there is an activation energy barrier. Therefore 
is the activation energy associated with reaction 1 and 
is the activation energy associated with reaction 2. The slow step has an activation energy barrier that is higher than that of the faster reaction.
| Todo 10 | Replace reference to Figure 30 with variable 2508-01 |
|---|
Each reaction also has its own activated complex. Remember that at the top of the activation energy barrier is the activated complex, the transition state between reactants and products that has the most potential energy. 
is the complex created in the first reaction. 
is the activated complex created in the second reaction. Thus, for this two-step process, there are two activated complexes.
Sample question: Draw the potential energy diagram for the following reaction mechanism 
. Properly label the diagram.
Solution:
Potential Energy Diagram for a Multi-step Process. (Source: Therese Forsythe. CC-BY-SA)
| Todo 10 | Replace reference to Figure 31 with variable Che-2508-02 |
|---|
Rate of Reaction is Determined by Slowest Step
In a series of reactions that make up a reaction mechanism, each individual reaction step has its own reaction rate that is determined by the five factors that have been discussed in this chapter. The overall reaction rate for the overall reaction (the sum of all the individual steps) can be determined from the rates of the individual steps. The relationship between the overall rate and the individual rates, however, is not what you might expect. Beginning students often think that to get the overall rate, you would add up the individual rates or average the individual rates but neither of these is correct. In fact, the overall rate for the reaction is exactly the same as the rate of the slowest step. Let's look at an example from life to see how this occurs.
Suppose you and two of your friends organize a car wash. You set up an assembly line operation where the cars in position 1 are wetted with a hose, then the cars get in line for position 2 where they are washed with soapy water and rinsed, and then they get in line for position 3 where they are towel dried.
| Todo 10 | Replace reference to Figure 32 with variable Chem-21-carwash |
|---|
Since you are the leader, you get the job of wetting the cars which takes 
to accomplish. The job of washing at station takes 
and the drying job requires 
When each job is finished, the car gets in line for the next station. In order to evaluate the efficiency of your assembly line, you count the minutes between the finished cars coming off the end of the line. The time lapse between completed cars is your reaction rate.
Regardless of Figure above, you should realize that there will be no cars in the line for station 
. Each car requires in station 2 and when the cars move to station 3, another car immediately goes to station 2. The drying of the first car and the washing of the second car begin at exactly the same time. Since it takes to wash and 
to dry, the car at station 3 is always finished and gone 
before the car at station 2. Therefore, the worker at station 3 always stands around and waits for before he gets his next car to dry. The time between cars coming off the line will be the the station 3 worker waits plus the required to dry . . . so the time between cars will be 
The reaction rate for this arrangement will be You should note that this overall reaction rate is exactly the same as the slowest step in the process, namely the wash step at station 2.
Suppose that you now bring in another person to work in your car wash and you assign that person to station 2 so that you have two people washing cars. The time to wash a car now becomes only 
since you have doubled the work force. Now the wash process will finish before the drying process and cars will back up in line for station 3. Your new arrangement now has a new slowest step - the drying stage is now the slowest step. Therefore, when the worker at station 3 finishes drying a car and the cars leaves the line, another car immediately enters station 3 and the worker immediately begins drying the car. After the worker at station 3 sends a car out, he will take to finish the next car and send it out. With this new organization, the time lapse between cars coming out will now be The overall reaction rate is faster. The important point is that the slowest step became and the overall rate became
Suppose once again that you bring another worker and put them at the drying station. The drying process now requires only 
with two people working. What is the slowest step in the process now? The slowest step is now the washing step again because it takes 
Now when the drying station finishes a car, there will be no cars in line because the washing station takes 2 more minutes to wash that it takes to dry. Therefore, once again, the drying station will have to wait for 
after drying a car before another car comes to station 3. With the 
wait and the 
drying time, the lapse between cars has now become The overall rate is and the slowest step is Regardless of how what the process is or how you set up the organization, the overall rate will always be exactly the same as the slowest step.
Let's look at another important point about this concept. We will go back to the original set up with a 
wetting (station 1), an 
washing (station 2), and a 
drying (station 3). Suppose when the first extra worker was brought in, you assigned that person to help with the drying so that the drying time became with the other times remaining the same. What would the overall rate be? The overall rate is exactly the same as the slowest step and the slowest step is the washing time. Therefore, you have increased the rate of the drying step but you have NOT affected the overall reaction rate. The only way you can alter the overall reaction rate is to increase the rate of the slowest step. Increasing the rate of steps other than the slowest step does nothing to the overall rate.
In chemical reactions, the speed of the other steps is so much faster than the slow step that the slow step is referred to as the rate-determining step. It is the speed of this slowest step that determines the rate of the overall reaction and changing the concentrations of the reactants in this step will change the rate.
Sample question:
In the reaction mechanism below, identify the rate-determining step and write the overall reaction.
Overall reaction: 
The overall reaction is found by adding the two elementary steps together and cancelling identical species that appear on both sides of the chemical equation, that is, 
and 
.
Solution:
The slow step (reaction 2) is the rate-determining step. It is the one marked slow. Whatever the reaction rate is for reaction 2, the overall rate will be exactly the same.
Lesson Summary
- A reaction mechanism is a multi-step process that is a combination of two or more elementary steps. An elementary step is a single, simple step in a multi-step process involving one or two particles.
- The rate-determining step is the slowest step in the reaction mechanism and the overall reaction rate will be exactly the same as the rate of the slowest step.
Review Questions
- Why do most reactions take place in more than one step?
- The overall rate of a reaction depends on
- the temperature
- the surface area
- the pressure
- the slowest step
Which step in the mechanism is the rate determining step?
or 
? Explain
Explain why this equation is not likely to represent the reaction mechanism.Vocabulary
- elementary step
- A single, simple step in a multi-step process involving one or two particles.
- reaction mechanism
- Most reactions do not take place in one step but rather occur as a combination of two or more elementary steps.
- rate-determining step
- The slowest step in the reaction mechanism.
