Gas-phase reactions are all about molecules colliding. Collision theory explains how these collisions lead to reactions, considering factors like energy, orientation, and frequency. The rate of a reaction depends on these collision factors and the concentrations of reactants.
Temperature and pressure play crucial roles in gas-phase reactions. Higher temperatures increase molecular energy and collision frequency , while higher pressures increase concentrations. Both effects typically speed up reactions. For complex reactions, the slowest step often determines the overall rate.
Collision Theory and Elementary Gas-Phase Reactions
Rate expressions in collision theory
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Collision theory explains reactions occur when reactant molecules collide with enough energy (activation energy ) and proper orientation
Elementary bimolecular gas-phase reaction (H2 + I2 -> 2HI):
Rate expression : rate = k [ \ce H 2 ] [ \ce I 2 ] \text{rate} = k[\ce{H2}][\ce{I2}] rate = k [ \ce H 2 ] [ \ce I 2 ] , k k k is rate constant , [ \ce H 2 ] [\ce{H2}] [ \ce H 2 ] and [ \ce I 2 ] [\ce{I2}] [ \ce I 2 ] are concentrations
k k k depends on collision frequency (Z A B Z_{AB} Z A B ) and steric factor (P P P ): k = P Z A B e − E a / R T k = PZ_{AB}e^{-E_a/RT} k = P Z A B e − E a / RT
Z A B Z_{AB} Z A B involves collision cross-section (σ A B \sigma_{AB} σ A B ), temperature (T T T ), and reduced mass (μ \mu μ ): Z A B = σ A B 8 π k B T μ Z_{AB} = \sigma_{AB}\sqrt{\frac{8\pi k_BT}{\mu}} Z A B = σ A B μ 8 π k B T
P P P considers proper orientation of colliding molecules
E a E_a E a is activation energy, R R R is gas constant, T T T is temperature
Elementary unimolecular gas-phase reaction (N2O5 -> 2NO2 + 1/2O2):
Rate expression: rate = k [ \ce N 2 O 5 ] \text{rate} = k[\ce{N2O5}] rate = k [ \ce N 2 O 5 ] , k k k is rate constant, [ \ce N 2 O 5 ] [\ce{N2O5}] [ \ce N 2 O 5 ] is concentration
Factors Affecting Gas-Phase Reaction Rates
Temperature and pressure effects on reactions
Temperature effects:
Higher temperature increases average kinetic energy of molecules, causing more collisions with enough energy to overcome activation energy barrier
Rate constant k k k increases exponentially with temperature according to Arrhenius equation: k = A e − E a / R T k = Ae^{-E_a/RT} k = A e − E a / RT
A A A is pre-exponential factor , E a E_a E a is activation energy, R R R is gas constant, T T T is temperature
Pressure effects:
Higher pressure increases reactant concentrations, leading to more frequent collisions and faster reaction rate
For gas-phase reaction \ce a A + b B − > p r o d u c t s \ce{aA + bB -> products} \ce a A + b B − > p ro d u c t s , rate is proportional to [ \ce A ] a [ \ce B ] b [\ce{A}]^a[\ce{B}]^b [ \ce A ] a [ \ce B ] b
Doubling pressure doubles concentrations, increasing rate by factor of 2 a + b 2^{a+b} 2 a + b
Multi-Step Gas-Phase Reaction Mechanisms
Rate-determining steps in reaction mechanisms
In multi-step reaction mechanism , slowest step is rate-determining step (RDS)
Overall reaction rate controlled by rate of RDS
Determining RDS:
Write rate expression for each elementary step
Compare rates of each step; slowest step is RDS
Overall rate law consistent with rate law of RDS
Reactants not in RDS do not appear in overall rate law
Steady-state approximation derives overall rate law from elementary steps
Assumes concentrations of reaction intermediates remain constant (formation rate = consumption rate)
Experimental Determination of Reaction Order
Experimental determination of reaction order
Reaction order with respect to reactant is power its concentration is raised to in rate law
For gas-phase reaction \ce a A + b B − > p r o d u c t s \ce{aA + bB -> products} \ce a A + b B − > p ro d u c t s with rate law rate = k [ \ce A ] m [ \ce B ] n \text{rate} = k[\ce{A}]^m[\ce{B}]^n rate = k [ \ce A ] m [ \ce B ] n , m m m and n n n are orders for \ce A \ce{A} \ce A and \ce B \ce{B} \ce B
Methods for determining reaction order:
Initial rates method : Compare initial rates of experiments varying concentration of one reactant while holding others constant
For reactant \ce A \ce{A} \ce A : rate 2 rate 1 = ( [ \ce A ] 2 [ \ce A ] 1 ) m \frac{\text{rate}_2}{\text{rate}_1} = \left(\frac{[\ce{A}]_2}{[\ce{A}]_1}\right)^m rate 1 rate 2 = ( [ \ce A ] 1 [ \ce A ] 2 ) m , solve for m m m
Graphical methods : Plot log of initial rate vs log of initial concentration of reactant; slope is order for that reactant
Integrated rate laws : Plot concentration vs time data using integrated rate law equations for different orders; best linear fit indicates correct order