Chemical Kinetics

Chapter 14

Chemical KineticsChemical Kinetics

ThermodynamicsThermodynamics – does a reaction take place? – does a reaction take place?

KineticsKinetics – how fast does a reaction proceed? – how fast does a reaction proceed?

Factors That Affect Reaction Rates

Physical State of the Reactants

Concentration of Reactants

Temperature

Presence of a Catalyst

Fig 14.2

Reaction rateReaction rate - the change in the concentration of a - the change in the concentration of a reactant or a product with time (reactant or a product with time (MM/s)/s)

A BA B

rate = −rate = −[A][A]

tt

rate = rate = [B][B]

tt

[A] = change in concentration of A over[A] = change in concentration of A over time period time period tt

[B] = change in concentration of B over[B] = change in concentration of B over time period time period tt

Because [A] decreases with time, Because [A] decreases with time, [A] is negative[A] is negative

rate = −[A]t

rate = [B]t

Fig 14.3 Progress of a hypothetical reaction A → B

• Average rate decreases as reaction proceeds• As the reaction goes forward, there are fewer collisions

between reactant molecules

Change of Rate with Time

C4H9Cl (aq) + H2O (l) → C4H9OH (aq) + HCl (aq)

Fig 14.4 Concentration of butylchloride as a function of time

• Instantaneous rate ≡ slope of line tangent to the curve at any point

Initial rate ≡ rate at t = 0

rate = −d[CH4]

dt

Reaction Rates and Stoichiometry

In this reaction, the ratio of C4H9Cl to C4H9OH is 1:1

Rate of consumption of C4H9Cl = rate of formation of C4H9OH

C4H9Cl (aq) + H2O (l) → C4H9OH (aq) + HCl (aq)

Rate =−[C4H9Cl]

t=

[C4H9OH]t

Reaction Rates and Stoichiometry

What if the ratio is not 1:1?

2 HI (g) → H2 (g)

• In such a case:

Rate = − 12

[HI]t

=[I2]t

aA + bB cC + dD

rate = −Δ[A]Δt

1a

= −Δ[B]Δt

1b

=Δ[C]Δt

1c

=Δ[D]Δt

1d

Reaction Rates and Stoichiometry

Eqn [14.4]

Fig 14.5 Basic components of a spectrophotometer

Br2 (aq) + HCOOH (aq) 2Br − (aq) + 2H+ (aq) + CO2 (g)

time

Br2 (aq) Br − (aq)

Br2 (aq) + HCOOH (aq) 2Br − (aq) + 2H+ (aq) + CO2 (g)

time

[Br2] Absorption3

93 n

m

Br2 (aq)

Br2 (aq) Br − (aq)

Beer’s Law:

A = abc

The Rate Law

Rate law - expresses the relationship of the rate of a reaction to the rate constant and the concentrations of the reactants raised to some powers

aA + bB cC + dD

Rate = k [A]x[B]y

reaction is xth order in A

reaction is yth order in B

reaction is (x +y)th order overall

F2 (g) + 2ClO2 (g) 2FClO2 (g)

rate = k [F2][ClO2]

Rate Laws

• Rate laws always determined experimentally

• Reaction order always defined in terms of reactant (not product) concentrations

• Order of a reactant is not related to the stoichiometric coefficient of the reactant in the balanced chemical equation

1

rate [Br2]

rate = k [Br2]

k = rate[Br2]

= rate constant

= 3.50 x 10-3 s-1

y = mx + b

Plot of rate vs [Br2]

Br2 (aq) + HCOOH (aq) 2Br − (aq) + 2H+ (aq) + CO2 (g)