Topic 21 Topic 21 Topic 21: Reaction Rates Table of Contents Basic Concepts Additional Concepts...

Topic 21Topic 21

Topic 21: Reaction Rates

Table of ContentsTable of Contents

Basic Concepts

Additional Concepts

Topic 21Topic 21

• An exothermic reaction releases heat, and an endothermic reaction absorbs heat.

Exothermic and Endothermic Reactions

Reaction Rates: Basic ConceptsReaction Rates: Basic Concepts Reaction Rates: Basic ConceptsReaction Rates: Basic Concepts

• The exothermic reaction gives off heat because the products are at a lower energy level than the reactants.

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Reaction Rates: Basic ConceptsReaction Rates: Basic Concepts Reaction Rates: Basic ConceptsReaction Rates: Basic Concepts Topic 21Topic 21

• The endothermic reaction absorbs heatbecause the products are at a higher energy level than the reactants.

• Scientists have observed that the energy released in the formation of a compound from its elements is always identical to the energy required to decompose that compound into its elements.




• This observation is an illustration of an important scientific principle known as the law of conservation of energy. That law states that energy is neither created nor destroyed in a chemical change, but is simply changed from one form to another.



• In an exothermic reaction, the heat released comes from the change from reactants at higher energy to products at lower energy.

• In an endothermic reaction, the heat absorbed comes from the opposite change.


• As you know, some chemical reactions are fast and others are slow; however, fast and slow are inexact, relative terms.

Expressing Reaction Rates

• Chemists, engineers, medical researchers, and others often need to be more specific.


• Recall from your study of math that the Greek letter delta (∆) before a quantity indicates a change in the quantity.


• In equation form, average rate or speed is written as

• We generally define the average rate of an action or process to be the change in a given quantity during a specific period of time.


• For chemical reactions, this equation defines the average rate at which reactants produce products, which is the amount of change of a reactant in a given period of time.


• Most often, chemists are concerned with changes in the molar concentration (mol/L, M) of a reactant or product during a reaction.


• Therefore, the reaction rate of a chemical reaction is stated as the change in concentration of a reactant or product per unit time, expressed as mol/(L∙s).


• Brackets around the formula for a substance denote the molar concentration. For example, [NO2] represents the molar concentration of NO2.


• It’s important to understand that reaction rates are determined experimentally by measuring the concentrations of reactants and/or products in an actual chemical reaction.


• Reaction rates cannot be calculated from balanced equations as stoichiometric amounts can.


• Reaction rates must always be positive.


• When the rate is measured by the consumption of a reactant, scientists commonly apply a negative sign to the calculation to get a positive reaction rate.

• Therefore, the following form of the average rate equation is used to calculate the rate of consumption:


• Reaction data for the reaction between butyl chloride (C4H9Cl) and water (H2O) is given in the table.

Calculating Average Reaction Rates


• Calculate the average reaction rate over thistime period expressed as moles of C4H9Cl consumed per liter per second.

• You are given the initial and final concentrations of C4H9Cl and the initial and final times.


• You can calculate the average reaction rate of the chemical reaction using the change in concentration of butyl chloride in four seconds.

• In this problem, the reactant butyl chloride is consumed.


• Write the equation for the average reaction rate, insert the known quantities, and perform the calculation.



• According to the collision theory, atoms, ions, and molecules must collide with each other in order to react.

Collision Theory

• The following three statements summarize the collision theory.

1. Particles must collide in order to react.

2. The particles must collide with the correct

orientation.


3. The particles must collide with enough energy to form an unstable activated

complex, also called a transition state, which is an intermediate particle made up of the joined reactants.

Collision Theory


• The minimum amount of energy that colliding particles must have in order to form an activated complex is called the activation energy of the reaction.

Collision Theory


• Particles that collide with less than the activation energy cannot form an activated complex.

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Collision Theory


• In an exothermic reaction, molecules collide with enough energy to overcome the activation energy barrier, form an activatedcomplex, then release energy and form products at a lower energy level.

Collision Theory


• In the reverse endothermic reaction, the reactant molecules lying at a low energy level

must absorb energy to overcome the activation energy barrier and form high-energy products.

• Reaction spontaneity is related to change in free energy, ∆G.

Collision Theory

• If ∆G is negative, the reaction is spontaneous under the conditions specified.

• If ∆G is positive, the reaction is not spontaneous.

• Let’s now consider whether spontaneity has any effect on reaction rates.


• To investigate the relationship between spontaneity and reaction rate, consider the following gas-phase reaction between hydrogen and oxygen.

Collision Theory

• Here, ∆G = –458 kJ at 298 K (25°C) and 1 atm pressure.

• Because ∆G is negative, the reaction is spontaneous.


• For the same reaction, ∆G = 2484 kJ, which means that the reaction is highly exothermic.

Collision Theory

• You can examine the speed of this reaction by filling a tape-wrapped soda bottle with stoichiometric quantities of the two gases—two volumes hydrogen and one volume oxygen.


• A thermometer in the stopper allows you to monitor the temperature inside the bottle.

Collision Theory

• As you watch for evidence of a reaction, the temperature remains constant for hours.

• Have the gases escaped? Or have they simply failed to react?


• If you remove the stopper and hold a burning splint to the mouth of the bottle, a reaction occurs explosively.

Collision Theory

• Clearly, the hydrogen and oxygen gases have not escaped from the bottle.

• Yet they did not react noticeably until you supplied additional energy in the form of a lighted splint.


• As these examples show, reaction spontaneity in the form of ∆G implies absolutely nothing about the speed of the reaction; ∆G merely indicates the natural tendency for a reaction or process to proceed.

Collision Theory

• Factors other than spontaneity, however, do affect the rate of a chemical reaction.

• These factors are keys in controlling and using the power of chemistry.


• The reaction rate for almost any chemical reaction can be modified by varying the conditions of the reaction.

Factors Affecting Reaction Rates


• An important factor that affects the rate of a chemical reaction is the reactive nature of the reactants. As you know, some substances react more readily than others.

The Nature of Reactants

• The tendency of a substance to react influences the rate of a reaction involving the substance.

• The more reactive a substance is, the faster the reaction rate.


• Reactions speed up when the concentrations of reacting particles are increased.

Concentration

• One of the fundamental principles of the collision theory is that particles must collide to react.

• The number of particles in a reaction makes a difference in the rate at which the reaction takes place.


• Think about a reaction where reactant A combines with reactant B.

Concentration

• At a given concentration of A and B, the molecules of A collide with B to produce AB at a particular rate.


• What happens if the amount of B is increased?

Concentration

• Increasing the concentration of B makes more molecules available with which A can collide.

• Reactant A “finds” reactant B more easily because there are more B molecules in the area, which increases probability of collision, and ultimately increases the rate of reaction between A and B.


Surface Area


• Increasing the surface area of reactants provides more opportunity for collisions withother reactants, thereby increasing the reaction rate.

• Generally, increasing the temperature at which a reaction occurs increases the reaction rate.

Temperature


Temperature


• As you can see, increasing the temperature of the reactants increases the reaction ratebecause raising the kinetic energy of the reacting particles raises both the collision frequency and the collision energy.

Temperature


• When the temperature is increased, more of the collisions at high temperatures haveenough energy to overcome the activation energy barrier and react.

• You’ve seen that increasing the temperature and/or the concentration of reactants can dramatically increase the rate of a reaction.

Catalysts

• However, an increase in temperature is not always the best (or most practical) thing to do.


Catalysts

• For example, suppose that you want to increase the rate of a reaction such as the decomposition of glucose in a living cell.

• Increasing the temperature and/or the concentration of reactants is not a viable alternative because it might harm or kill the cell.


Catalysts

• It is a fact that many chemical reactions in living organisms would not occur quickly enough to sustain life at normal living temperatures if it were not for the presence of enzymes.

• An enzyme is a type of catalyst, a substance that increases the rate of a chemical reaction without itself being consumed in the reaction.


Catalysts

• Although catalysts are important substances in a chemical reaction, a catalyst does not yield more product and is not included in the product(s) of the reaction.

• In fact, catalysts are not included in the chemical equation.


Catalysts


• This energy diagram shows how the activation energy of the catalyzed reaction is lower andtherefore the reaction produces the products at a faster rate than the uncatalyzed reaction does.

Catalysts

• Note that the activation energy for the catalyzed reaction is much lower than for the uncatalyzed reaction.


• The lower activation energy for the catalyzed reaction means that more collisions have sufficient energy to initiate reaction.

Catalysts

• Another type of substance that affects reaction rates is called an inhibitor.

• Unlike a catalyst, which speeds up reaction rates, an inhibitor is a substance that slows down, or inhibits, reaction rates.

• Some inhibitors, in fact, actually prevent a reaction from happening at all.


Heterogeneous and homogeneous catalysts

• A heterogeneous catalyst exists in a physical state different than that of the reaction it catalyzes.

• Because the catalysts in a catalytic converter are solids and the reactions they catalyze are gaseous, the catalysts are called heterogeneous catalysts.


Heterogeneous and homogeneous catalysts

• A catalyst that exists in the same physical state as the reaction it catalyzes is called a homogeneous catalyst.

• For example, if both an enzyme and the reaction it catalyzes are in aqueous solution, the enzyme is a homogeneous catalyst.


Basic Assessment QuestionsBasic Assessment Questions

Question 1

In aqueous solution, bromine reacts with formic acid (HCOOH) to produce hydrogen bromide and carbon dioxide gas. Calculate the average reaction rate over 150 s expressed in terms of the disappearance of Br2.

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AnswerTopic 21Topic 21


Question 2

Use the factors discussed in this section to account for the following observation.

Topic 21Topic 21

When two gases react, compressing the gases generally increases the rate of reaction.


Question 3

The rate of gaseous reactions can also be increased by pumping more gas into the reaction container.

Topic 21Topic 21

Use the factors discussed in this section to account for the following observation.



Adding gas increases concentration by adding particles to a given space.

Reaction Rates: Additional ConceptsReaction Rates: Additional Concepts Reaction Rates: Additional ConceptsReaction Rates: Additional Concepts Topic 21Topic 21

Additional Concepts

Reaction Rates: Additional ConceptsReaction Rates: Additional Concepts Reaction Rates: Additional ConceptsReaction Rates: Additional Concepts

Reaction Rate Laws

• The equation that expresses the mathematical relationship between the rate of a chemical reaction and the concentration of reactants is called a rate law.

• For example, the reaction A → B, which is a one-step reaction, has only one activated complex between reactants and products.

Topic 21Topic 21

Reaction Rate Laws

• The rate law for this reaction is expressed as, Rate = k [A] where k is the specific rate constant, or a numerical value that relates reaction rate and concentration of reactants at a given temperature.


Reaction Rate Laws• Units for the rate constant include L/(mol∙s),

L2/(mol2∙s), and s–1. • Depending on the reaction conditions,

especially temperature, k is unique for every reaction.

• The rate law means that the reaction rate is directly proportional to the molar concentration of A.

• Thus, doubling the concentration of A will double the reaction rate.


Reaction Rate Laws• Increasing the concentration of A by a factor

of 5 will increase the reaction rate by a factor of 5.

• The specific rate constant, k, does not change with concentration; however, k does change with temperature.

• A large value of k means that A reacts rapidly to form B.

• What does a small value of k mean?


Reaction order

• In the expression Rate = k[A], it is understood that the notation [A] means the same as [A]1.

• In other words, for reactant A, the understood exponent 1 is called the reaction order.


Reaction order

• The reaction order for a reactant defines how the rate is affected by the concentration of that reactant.

• For example, the rate law for the decomposition of H2O2 is expressed by the following equation.


Reaction order

• Because the reaction rate is directly proportional to the concentration of H2O2 raised to the first power, [H2O2]1, the decomposition of H2O2 is said to be first order in H2O2.

• Because the reaction is first order in H2O2, the reaction rate changes in the same proportion that the concentration of H2O2 changes.


Reaction order

• So if the H2O2 concentration decreases to one-half its original value, the reaction rate is halved as well.


Reaction order

• Because reaction order is based on reaction rates, it follows that reaction order also is determined experimentally.

• Finally because the rate constant describes the reaction rate, k, too, must be determined experimentally.


Determining Rate Laws

• Consider the following reaction and its experimentally determined rate law.

• This rate equation shows that the rate depends on the concentrations of both NO2 and F2, each to the first power.



• In other words, the reaction is first order in NO2 and first order in F2.

• So, if [NO2] is doubled while [F2] remains the same, the rate doubles.

• Also, if [F2] is doubled while [NO2] remains the same, the reaction rate doubles.



• Now, examine the following reaction and its rate law.

• Because the rate depends on the square of the concentration of NO, doubling [NO] while leaving [O2] the same will increase the reaction rate by a factor of 22, or 4.



• Use the data in the table below to determine the form of the rate law for the following reaction.



• The general rate law for this type of reaction is as follows.

• To start compare the data from trial 1 and 2.

• Notice that [NO] in trial 1 is 4.0 x 10–3 , mo1/L and [NO] in trial 2 is 8.0 x 10–3

mo1/L, double that in trial 1.



• Now see how the rate changes from trial 1 to trial 2.

• The rate in trial 2, 4.8 x 10–5 mo1/L ∙s, is four times the rate in trial 1, which is 1.2 x 10–5 mo1/L∙s.



• When [NO] doubles, the initial reaction rate quadruples.

• Therefore, it is likely that the reaction rate depends on the square of the concentration of NO.



• Next, determine how the rate depends on the change in [H2].

• When [H2] doubles, the initial rate doubles.

• This result indicates that the rate is directly proportional to the concentration of H2.



• Now you can write the rate law based on your comparisons.

• The rate law means that the reaction is second order in [NO] and first order in [H2].


• In a complex reaction, a certain series of steps must occur with correct fit and in an exact sequence in order to yield the reaction products. This series of steps is called a reaction mechanism.

Instantaneous Reaction Rates and Reaction Mechanisms

• Many chemical reactions are complex.



• As an example, consider that hydrogen peroxide (H2O2) decomposes when iodide ions (I–) are present.

• Experiments have shown that the reaction takes place in two steps as follows.



• The IO– ion is called an intermediate in the reaction.

• An intermediate is an atom, an ion, or a molecule produced in one step of a reaction and consumed in a later step.



• In a complex reaction, one step is always slower than the others.

• This step is called the rate-determining step because it will determine how fast the reaction forms products, no matter how fast the other steps are.


• In the reaction, the first step is the slower step, therefore, it is the rate-determining step for the reaction

Additional Assessment QuestionsAdditional Assessment Questions

Question 1 Topic 21Topic 21

The rate law for a reaction is rate = k[A][B]. What happens to the rate if the concentration of B is tripled while [A] remains the same? If the concentration of A is halved while [B] remains the same?

Answer

rate triples; rate halves

Additional Assessment QuestionsAdditional Assessment QuestionsTopic 21Topic 21

Question 2


The rate law for a reaction is rate = k[X][Y]2. What happens to the rate if the concentration of X is doubled while [Y] remains the same?

Answer

rate doubles


Question 3

Discuss two possible ways to determine the average rate of the following reaction in the laboratory. What quantities would you measure in each case?


Answer

There are three possibilities. Measure the reduction in concentration of H2, the reduction

in concentration of I2, or the increase in

concentration of HI. In all three cases, time is the second quantity that must be measured.


Question 4

The collision theory says that particles must collide in order to react. What other two requirements must be met in order for particles to react?


Answer

The particles must collide in the correct orientation and with sufficient energy to form an activated complex.


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