Elementary reactions and molecularity are key concepts in understanding reaction mechanisms. They help us break down complex reactions into simpler steps, making it easier to analyze and predict reaction rates.
Molecularity tells us how many molecules are involved in a single reaction step. This concept is crucial for determining rate laws and understanding the likelihood of different types of reactions occurring in chemical systems.
Elementary vs Complex Reactions
Defining Elementary and Complex Reactions
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Elementary reactions occur in a single step and involve the direct conversion of reactants to products without any intermediate species
Reactant molecules collide and transform into product molecules in a single event (A + B → C)
No are formed during the process
Complex reactions involve multiple elementary steps and the formation of one or more reaction intermediates before yielding the final products
Reaction proceeds through a series of elementary steps, each with its own
Intermediate species are formed and consumed during the reaction (A + B → X, X + C → D)
Overall reaction may have a different than individual elementary steps
Relationship Between Elementary Steps and Overall Reaction
The overall reaction equation for a complex reaction is the sum of all the elementary steps in the reaction mechanism
Elementary steps are added together to give the balanced equation for the overall reaction
Reaction intermediates cancel out when elementary steps are summed
Rate of the overall reaction depends on the rates of individual elementary steps
Molecularity of Elementary Reactions
Defining Molecularity
Molecularity is defined as the number of reactant molecules that must collide simultaneously for an elementary reaction to occur
Represents the number of molecules involved in the rate-determining step
Determines the kinetic order of the reaction with respect to each reactant
The molecularity of an elementary reaction can be determined by counting the number of reactant molecules in the balanced chemical equation for that specific reaction step
Reactant molecules are counted on the left side of the equation arrow
Coefficients in the balanced equation do not affect molecularity
Molecularity is always a positive integer (1, 2, or 3) and should not be confused with the stoichiometric coefficients in the balanced equation
Stoichiometric coefficients represent the relative amounts of reactants and products
Molecularity refers to the number of molecules involved in the rate-determining step
Classifying Elementary Reactions
Unimolecular Reactions
Unimolecular reactions have a molecularity of one, involving a single reactant molecule undergoing a reaction, such as isomerization or decomposition
Reaction involves the rearrangement or fragmentation of a single molecule (A → B)
Examples include radioactive decay, thermal decomposition, and cis-trans isomerization
Rate law for a : Rate = k[A]
Rate is directly proportional to the concentration of the reactant molecule
Bimolecular Reactions
Bimolecular reactions have a molecularity of two, involving the collision and reaction of two reactant molecules, which can be the same or different species
Two molecules must collide with sufficient energy and proper orientation for the reaction to occur (A + B → C)
Examples include Diels-Alder reactions, SN2 reactions, and the formation of hydrogen iodide (H2 + I2 → 2HI)
Rate law for a : Rate = k[A][B] or Rate = k[A]^2 (if A and B are the same species)
Rate depends on the concentrations of both reactant molecules
Termolecular Reactions
Termolecular reactions have a molecularity of three, involving the simultaneous collision and reaction of three reactant molecules
Three molecules must collide simultaneously with sufficient energy and proper orientation (A + B + C → D)
Examples include the formation of ozone (O + O2 + M → O3 + M) and the recombination of nitric oxide with oxygen (2NO + O2 → 2NO2)
Rate law for a termolecular reaction: Rate = k[A][B][C]
Rate depends on the concentrations of all three reactant molecules
Rate Law and Molecularity
Relationship Between Rate Law and Molecularity
For an elementary reaction, the rate law can be determined directly from the balanced chemical equation
Molecularity of the reaction determines the form of the rate law
Order of the reaction with respect to each reactant is equal to its stoichiometric coefficient in the balanced equation
The rate law for an elementary reaction is proportional to the product of the concentrations of the reactants, with each concentration raised to a power equal to its molecularity
Unimolecular reaction: Rate = k[A]
Bimolecular reaction: Rate = k[A][B] or Rate = k[A]^2 (if A and B are the same species)
Termolecular reaction: Rate = k[A][B][C]
The overall order of an elementary reaction is equal to its molecularity
Sum of the exponents in the rate law equals the molecularity of the reaction
Unimolecular reactions are first-order, bimolecular reactions are second-order, and termolecular reactions are third-order
Determining Rate Laws from Elementary Steps
Rate law for a complex reaction cannot be determined directly from the overall balanced equation
Overall rate law depends on the slowest elementary step (rate-determining step)
Reaction intermediates do not appear in the overall rate law
To determine the rate law for a complex reaction, identify the slowest elementary step and use its rate law as the overall rate law
Slowest step is the bottleneck of the reaction and determines the overall rate
Other elementary steps are assumed to be at equilibrium and do not affect the overall rate
Molecularity: Unimolecular vs Bimolecular vs Termolecular
Prevalence of Different Molecularities
Unimolecular and bimolecular reactions are the most common types of elementary reactions
Majority of elementary reactions involve one or two reactant molecules
Unimolecular reactions are common in isomerization, decomposition, and radioactive decay
Bimolecular reactions are prevalent in substitution, addition, and elimination reactions
The probability of three molecules colliding simultaneously with sufficient energy and proper orientation is extremely low, making termolecular reactions rare
Collision of three molecules is statistically improbable
Termolecular reactions have very low reaction rates and are kinetically unfavorable
Most apparent termolecular reactions are actually a sequence of bimolecular and/or unimolecular elementary steps
Overall reaction may appear to be termolecular but proceeds through a series of lower-molecularity steps
Reaction intermediates are formed and consumed in the process
Examples of Rare Termolecular Reactions
Formation of ozone (O + O2 + M → O3 + M)
Involves the collision of an oxygen atom, an oxygen molecule, and a third molecule (M) that absorbs excess energy
Third molecule (M) is necessary to stabilize the ozone molecule and prevent its rapid decomposition
Recombination of nitric oxide with oxygen (2NO + O2 → 2NO2)
Involves the simultaneous collision of two nitric oxide molecules and an oxygen molecule
Reaction is important in the formation of nitrogen dioxide in the atmosphere
Termolecular reactions are rare exceptions to the general trend of unimolecular and bimolecular elementary steps
Require highly specific conditions and molecular orientations
Play a minor role in the overall kinetics of most chemical systems