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Chapter 15
Chemical Kinetics
 
 
 
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 Ozone cycle in the stratosphere
Notes
 Stratospheric ozone. In the stratosphere radiation from the sun causes O2 to dissociate into O atoms. When the O atoms collide with O2 molecules, O3 may be formed. This O3 molecule is a high energy species which either transfers its excess energy to a molecule it collides with or dissociates to form O2 and an O atom. A stable O3 molecule can absorb ultraviolet light from the sun and decompose to form O2 and an O atom.
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 Catalytic decomposition of ozone by chlorine atoms from CFCs.
Notes
 CFC's and Stratospheric ozone. CFCs such as freon are very stable molecules in the lower atmosphere. When they diffuse into the stratosphere ultraviolet light from the sun causes a C-Cl bond to break, producing Cl atoms which react with O3 to produce O2 and ClO. The ClO collides with an O atom to produce O2 and a Cl atom. The Cl atom is a catalyst which breaks down numerous O3 molecules, causing depletion of the O3 in the stratosphere.
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 NO acts as a catalyst to break down ozone in the stratosphere
Notes
 Catalytic destruction of Stratospheric ozone. NO is present in trace amounts in the stratosphere. When NO collides with O3 the products are NO2 and O2. When the resulting NO2 atom collides with an O atom it produces O2 and NO. Since NO is unchanged in the reaction, it catalyzes the conversion of O3 and an O atom to two O2 molecules. Since NO acts as a catalyst, one NO molecule can cause the breakdown of many O3 molecules.
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 The reaction between Al and Br2 to form AlBr3 is demonstrated.
Notes
 Formation of Aluminum Bromide. When Al is placed on the surface of Br2 an exothermic reaction occurs. The Al is oxidized to Al3+ by the Br2, which is reduced to Br- ions. The ionic product, AlBr3, can be observed on the watch glass after the reaction.
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 Copper oxide reacts with carbon to form copper and carbon dioxide
Notes
 Reactions with CuO. When black carbon and black copper oxide are heated together the Cu2+ ions are reduced to metallic Cu and a gas is evolved. When the gas is collected in Ca(OH)2 a white precipitate of CaCO3 is formed. The reaction which occurs involves the reduction of Cu2+ ions by carbon which is oxidized to CO2.
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 A flame from a candle causes hydrogen gas in a balloon to react explosively with oxygen in the air.
Notes
 Formation of water. An explosive reaction occurs when a candle touches a balloon filled with hydrogen gas. The hydrogen reacts with oxygen in the air to form water in an exothermic reaction.
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 The decomposition of water to form hydrogen and oxygen.
Notes
 electrolysis of water. When a direct current is passed through water it decomposes to form oxygen and hydrogen. The volume of hydrogen gas produced at the negative electrode is twice the volume of the oxygen gas formed at the positive electrode. This indicates that water contains twice as many hydrogen atoms as oxygen atoms, which is an illustration of the law of constant composition.
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 Illustration of the half life of a first order process
Notes
 First order process. The rate of a first order process is proportional to the concentration of a single substance. The half-life of a reaction is the time required for the concentration of a reactant to decrease to half of its original value. For a first order reaction the half-life is independent of the concentration of the substance, so all of the half-lives have the same value.
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 The explosive decomposition of nitrogen triiodide is demonstrated.
Notes
 Nitrogen triiodide. Nitrogen triiodide is extremely unstable when it is dry. Touching it with a feather causes it to decompose explosively. The explosion occurs as chemical energy is released by the decomposition of nitrogen triiodide to N2 and I2. Violet iodine vapor can be observed after the explosion.
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 The exothermic reaction between Fe2O3 and aluminum is demonstrated.
Notes
 Thermite reactions. In the thermite reaction, Al reduces Fe2O3 to Fe in an extremely exothermic reaction in which Al is oxidized to Al2O3. The reaction produces enough heat to melt the iron. Because of the extreme heat produced in the thermite reaction, it is used industrially to weld iron.
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 Mechanism for the reaction of methyl chloride with a hydroxide ion
Notes
 Bimolecular reaction. A reaction mechanism is the process by which the reaction occurs. In the reaction of CH3Cl with OH-, the OH- approaches the CH3Cl molecule on the side opposite the Cl atom. In the transition state a bond begins to form between the OH- and C atoms, the H atoms assume a planar arrangement, and the C-Cl bond gets longer. As the C-OH bond forms and Cl- leaves, the arrangement of the H atoms is inverted.
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 Ozone cycle in the stratosphere
Notes
 Stratospheric ozone. In the stratosphere radiation from the sun causes O2 to dissociate into O atoms. When the O atoms collide with O2 molecules, O3 may be formed. This O3 molecule is a high energy species which either transfers its excess energy to a molecule it collides with or dissociates to form O2 and an O atom. A stable O3 molecule can absorb ultraviolet light from the sun and decompose to form O2 and an O atom.
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 Catalytic decomposition of ozone by chlorine atoms from CFCs
Notes
 CFC's and Stratospheric ozone. CFCs such as freon are very stable molecules in the lower atmosphere. When they diffuse into the stratosphere ultraviolet light from the sun causes a C-Cl bond to break, producing Cl atoms which react with O3 to produce O2 and ClO. The ClO collides with an O atom to produce O2 and a Cl atom. The Cl atom is a catalyst which breaks down numerous O3 molecules, causing depletion of the O3 in the stratosphere.
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 NO acts as a catalyst to break down ozone in the stratosphere
Notes
 Catalytic destruction of Styatospheric ozone. NO is present in trace amounts in the stratosphere. When NO collides with O3 the products are NO2 and O2. When the resulting NO2 atom collides with an O atom it produces O2 and NO. Since NO is unchanged in the reaction, it catalyzes the conversion of O3 and an O atom to two O2 molecules. Since NO acts as a catalyst, one NO molecule can cause the breakdown of many O3 molecules.
15.1
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 15-1
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 Experimental set-up for determining the rate of decomposition of H2O2. Oxygen gas given off by the reaction mixture is trapped, and its volume is measured in the gas buret. The amount of H2O2 consumed and the remaining concentration of H2O2 can be calculated from the measured volume of O2(g).
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15.2
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 15-2
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 Graphical representation of kinetic data for the reaction: 2 H2O2(aq) n 2 H2O(l) 1 O2(g). This is the usual form in which concentration-time data are plotted. Reaction rates are determined from the slopes of the tangent lines. The blue line has a slope of 21.70 M/2800 s 5 26.1 3 1024 M s21. The slope of the black line and its relation to the initial rate of reaction are described in Example 15-2.
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15.3
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 15-3
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 A zero-order reaction: A n products. The initial concentration of the reactant A is [A]0, that is, [A] 5 [A]0 at t 5 0. [A] decreases at a constant rate until the reaction stops. This occurs at the time, tf, where [A] 5 0. The slope of the line is (0 2 [A]0)/(tf 2 0) 5 2[A]0/tf. The rate constant is the negative of the slope: k 5 2slope 5 [A]0/tf.
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15.4
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 15-4
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 Test for a first-order reaction: decomposition of H2O2(aq). Here we plot ln [H2O2] vs. t. The data are based on Table 15.1 and are listed below. The slope of the line is used in the text. t, s [H2O2], M ln [H2O2] ÊÊÊ0 2.32 0.842 Ê200 2.01 0.698 Ê400 1.72 0.542 Ê600 1.49 0.399 1200 0.98 20.020 1800 0.62 20.48 3000 0.25 21.39
Notes
  
15.5
Title
 15-5
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 Decomposition of di-t-butyl peroxide (DTBP) at 147 ¼C. The decomposition reaction is described through Equation (15.15). In this graph of the partial pressure of DTBP as a function of time, three successive half-life periods of 80 min each are indicated. This constancy of the half-life is proof that the reaction is first order.
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15.6
Title
 15-6
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 A straight-line plot for the second-order reaction A n products. The reciprocal of the concentration, 1/[A], is plotted against time. As the reaction proceeds, [A] decreases and 1/[A] increases in a linear fashion. The slope of the line is the rate constant k.
Notes
  
15.7
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 15-7
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 Testing for the order of a reaction-Example 15-8 illustrated. The straight-line plot is obtained for 1/[A] vs. t, graph (3). The reaction is second order.
Notes
  
15.8
Title
 15-8
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 Distribution of molecular kinetic energies. At both temperatures, the fraction of all molecules having kinetic energies in excess of the value marked by the heavy black arrow is small. (Note the shaded areas on the right.) At the higher temperature T2 (red), however, this fraction is considerably larger than at the lower temperature T1 (blue).
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15.8.1UN
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 15-9
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 Orientation of the colliding molecules is a crucial matter in the reaction of N2O and NO
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15.10.1C
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 15-10
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 A reaction profile for the reaction N2O(g) + NO(g) --> N2(g) + NO2(g)
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15.11
Title
 15-11
Caption
 An analogy to reaction profile and activation energy. A hike (red path) is taken from the valley on the left (reactants) over the ridge to the valley on the right (products). The ridge above the starting point corresponds to the transition state. It is probably the height of this ridge (activation energy) more than anything else that determines how many people are willing to take the hike, regardless of the fact that it is all “downhill” on the other side.
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15.12
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 15-12
Caption
 Temperature dependence of the rate constant k for the reaction 2 N2O5 (in CCl4) n 2 N2O4 (in CCl4) 1 O2(g). Data are plotted as follows, for the representative point in black. t 5 25 ¼C 5 298 K 1/T 5 1/298 5 0.00336 5 3.36 3 1023 K21 k 5 3.46 3 1025 s21; ln k 5 ln 3.46 3 1025 5 210.272 To evaluate Ea, slope of line 5 2Ea/R 5 21.2 3 104 K Ea 5 8.3145 J mol21 K21 3 1.2 3 104 K 5 1.0 3 105 J/mol 5 1.0 3 102 kJ/mol (A more precise plot yields a value of Ea 5 106 kJ/mol. The arrow points to data referred to in Example 15-9.)
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15.13
Title
 15-13
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 The San Ysidro/Tijuana border station: an analogy to a rate-determining step. Note the tie-up of traffic on the Mexican side of the border (top) and the relatively few cars on the United States side (bottom). This station is a “bottleneck” and hence the rate-determining part of the trip by car from Tijuana, Mexico, to San Diego, CA, two cities located only about a dozen miles apart.
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15.14
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 15-14
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Notes
 A reaction profile for a two-step mechanism.
15.15
Title
 15-15
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 An example of homogeneous catalysis. The activation energy is lowered in the presence of H1, a catalyst for the decomposition of HCOOH.
Notes
  
15.16
Title
 15-16
Caption
 Heterogeneous catalysis in the reaction 2 CO 1 2 NO 78nRh 2 CO2 1 N2. (a) Molecules of CO and NO are adsorbed on the rhodium surface. (b) The adsorbed NO molecules dissociate into adsorbed N and O atoms. (c) Adsorbed CO molecules and O atoms combine to form CO2 molecules, which desorb into the gaseous state. Two N atoms combine and are desorbed as an N2 molecule.
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15.17
Title
 15-17
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 A reaction profile for a surface-catalyzed reaction
Notes
  
15.17.1UN
Title
 P 611
Caption
  
Notes
 A computer graphic representationÉ.
15.18a-c
Title
 15-18
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 Lock-and-key model of enzyme action. (a) The substrate attaches itself to an active site on an enzyme molecule. (b) Reaction occurs. (c) Product molecules detach themselves from the site, freeing the enzyme molecule to attach another molecule of substrate. The substrate and enzyme must have complementary structures to produce a complex, hence the term lock-and-key.
Notes
  
15.19
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 15-19
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 Effect of substrate concentration on the rate of an enzyme reaction.
Notes
  
P 602
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 P 602
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Notes
 Fireflies
Table 15.1
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 Table 15.1
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Notes
 Decomposition of H2O2
Table 15.2
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 Table 15.2
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Notes
 Decomposition of H2O2-derived rate data
Table 15.3
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 Table 15.3
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Notes
 Kinetic data for the reaction:***
Table 15.4
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 Table 15.4
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Notes
 Some typical first-order processes
Table 15.8
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 Table 15.8
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Notes
 Kinetic data for example 15.8