Section - 3 REACTION MECHANISMS Collision Theory of ... · LOCUS 1 Chemistry / Chemical Kinetics...
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LOCUS LOCUS LOCUS LOCUS LOCUS 1 Chemistry / Chemical Kinetics Collision Theory of Reaction Rates The fundamental notion of the collision theory of reaction rates is that for reaction to occur, molecules, atoms, or ions must first collide. Increased concentrations of reacting species result in greater number of collisions per unit time. However, not all collisions result in reaction; i.e., not all collisions are effective collisions. For a collision to be effective, the reacting species must (1) possess at least a certain minimum energy necessary to rearrange outer electrons in breaking bonds and forming new ones and (2) have the proper orientations toward each other at the time of collision. Collisions must occur in order for a chemical reaction to proceed, but they do not guarantee that a reaction will occur. A collision between atoms, molecules, or ions is not like one between two hard billiard balls. Whether or not chemical species “collide” depends on the distance at which they can interact with one another. For instance, the gas phase ion-molecule reaction 4 4 5 3 CH CH CH CH can occur with a fairly long-range contact. This is because the interactions between ions and induced dipoles are effective over a relatively long distance. By contrast, the reacting species in the gas reaction 3 3 2 6 CH CH CH are both neutral. They interact appreciably only through very short-range forces between induced dipoles, so they must approach one another very closely before we could say that they “collide.” Recall that the average kinetic energy of a collection of molecules proportional to the absolute temperature. At higher temperatures, more of the molecules possess sufficient energy to react. If colliding molecules have improper orientations, they do not react even though they may possess sufficient energy. Figure 8 depicts collisions between molecules of NO and N 2 O. We assume that each possesses sufficient energy to react according to 2 2 2 NO NO NO N (a) O N N O N O N N O N O N O N N Product Effective orientation of collision (b) O N O N O N N O N O N N No reaction Ineffective orientation of collision N O N Section - 3 REACTION MECHANISMS
Transcript of Section - 3 REACTION MECHANISMS Collision Theory of ... · LOCUS 1 Chemistry / Chemical Kinetics...
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Chemistry / Chemical Kinetics
Collision Theory of Reaction RatesThe fundamental notion of the collision theory of reaction rates is that for reaction to occur, molecules, atoms,or ions must first collide. Increased concentrations of reacting species result in greater number of collisions per unittime. However, not all collisions result in reaction; i.e., not all collisions are effective collisions. For a collision tobe effective, the reacting species must (1) possess at least a certain minimum energy necessary to rearrange outerelectrons in breaking bonds and forming new ones and (2) have the proper orientations toward each other at thetime of collision.
Collisions must occur in order for a chemical reaction to proceed,but they do not guarantee that a reaction will occur.
A collision between atoms, molecules, or ions is not like one between two hard billiard balls. Whether or notchemical species “collide” depends on the distance at which they can interact with one another. For instance, thegas phase ion-molecule reaction 4 4 5 3CH CH CH CH can occur with a fairly long-range contact. Thisis because the interactions between ions and induced dipoles are effective over a relatively long distance. Bycontrast, the reacting species in the gas reaction 3 3 2 6CH CH C H are both neutral. They interact appreciablyonly through very short-range forces between induced dipoles, so they must approach one another very closelybefore we could say that they “collide.”Recall that the average kinetic energy of a collection of molecules proportional to the absolute temperature. Athigher temperatures, more of the molecules possess sufficient energy to react.If colliding molecules have improper orientations, they do not react even though they may possess sufficient energy.Figure 8 depicts collisions between molecules of NO and N2O. We assume that each possesses sufficient energyto react according to
2 2 2NO N O NO N
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ON
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Effectiveorientationof collision
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Section - 3 REACTION MECHANISMS
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Chemistry / Chemical Kinetics
(c)
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Figure 8: Some possible collisions between N2O and NO molecules in the gas phase. (a) A collision that couldbe effective in producing the reaction. (b, c) Collision would be ineffective. The molecules must have theproper orientations relative to each other and have sufficient energy to react.
Transition State TheoryChemical reactions involve the making and breaking of chemical bonds. The energy associated with a chemicalbond is a form of potential energy. Reactions are accompanied by changes in potential energy. Consider thefollowing hypothetical, one-step exothermic reaction at a certain temperature.
2A B AB B heatFigure 9 shows a plot of potential energy versus reaction coordinate. In Figure 9(a) the ground state energy of thereactants, A and B2, is higher than the ground state energy of the products, AB and B. The energy released in thereaction is the difference between these two energies, E. It is related to the change in enthalpy or heat contentthat you have studies in thermodynamics.Quite often, for reaction to occur, some covalent bonds must be broken so that others can be formed. This canoccur only if the molecules collide with enough kinetic energy to overcome the potential energy stabilisation ofthe bonds. According to the transition state theory, the reactants pass through a short-lived, high-energyintermediate state, called a transition state, before the product are formed.
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transition statereactants productsABA +B AB + B
A B — B A B B A — B B
A B B
Eforwardreaction
a Ereversereaction
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ProductsAB + B
Ereaction
ReactantsA + B2
Reaction coordinate
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A B B
Eforwardreaction
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Ereversereaction
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ProductsAB + B
EreactionReactantsA + B2
Reaction coordinate
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Endothermic reactionExothermic reaction
The "reaction coordinate" representsthe leading from reactants to products.This coordinate is sometimes labeled"progress of reaction."
progress along the pathwayFigure 9: A potential energy diagram. (a) A reaction that releasesenergy (exothermic). An example of an exothermic gas-phase reacion is H + I HI + I(b) A reaction that absorbs energy (endothermic). An example of anendothermic gas-phase reaction is I + H HI + I
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