Rates of reaction
Rates of Reaction
Chemical kinetics is the study of reaction rates, how reactions change under varying conditions, and what molecular events occur during the overall reaction
By noting how the rate of a reaction is affected by changing conditions,
you can learn what is happening at the molecular level!
Rates of Reaction
Variables affecting reaction rates:
Concentration of reactants
The rate of reaction is often increased when the concentration of a reactant is increased; some reactions only require that some reactant be present!
13.1­13.3
Rates of reaction
Rates of Reaction
Surface Area of a solid reactant
Since the reaction occurs at the surface of the solid, more surface area increases the reaction rate
Rates of Reaction
Temperature
Usually reactions speed up when the temperature increases
13.1­13.3
Rates of reaction
Rates of Reaction
Catalysts
A substance that increases the rate of reaction without being consumed in the overall reaction
MnO2
2 H2O2 (aq) 2 H
2O (l) + O2 (g) Reaction Rate: the increase in molar concentration of product of a reaction per unit time, or decrease in molar concentration of reactant per unit time
mol/L sec
mol
L sec
[ ] = molar concentration!
average reaction rate =
13.1­13.3
[ ]
t
Rates of reaction
Average Reaction Rate
A car travels 84 miles in 2 hours.
What was its average rate of travel?
42 mi/hr
2 N2O5 (g) 4 NO2 (g) + O2 (g)
Time
[N2O5]
600 s
1.24 x 10 M
1200 s
0.93 x 10­2 M
(0.93 ­ 1.24) x 10­2 M
­2
(1200­600) s
­ 0.31 x 10­2 M
600 s
= 5.2 x 10­6 M/s
*rates expressed in reactant terms = negative
reactants decreasing, so makes rate positive!
Experimental Determination of Rate
To obtain the rate of a reaction, you must determine [reactant] or [product] during the course of the reaction
­ slow reaction, take samples and analyze
­ gas reaction, monitor pressure changes
­ color changes can be monitored through absorption with a spectrometer or IR/NMR
13.1­13.3
Rates of reaction
Dependence of Rate on Concentration
2 NO2 (g) + F2 (g) 2 NO2F (g)
Experimentally, it is seen that if you double [NO2] or [F2] the reaction rate doubles
Rate Law:
an equation that relates the rate of a reaction to the [reactants] (or catalyst) raised to various powers
2 NO2 (g) + F2 (g) 2 NO2F (g)
Rate = k [NO2] [F2]
Rate constant: fixed value at any given temperature
varies with temperature!
units: depend on the form of the rate law
k =
13.1­13.3
rate
[NO2] [F2]
=
mol
L s
mol 2
L
=
L
mol s
Rates of reaction
Rate = k [NO2] [F2]
Exponents are typically integers
determined experimentally!
not just from balanced equation!
From the rate law...
when you know k, you can calculate the rate of a reaction for any reactant concentrations
Reaction Order
A reaction can be classified by its orders; the reaction order is the exponent of a species in the rate law, as determined experimentally
The overall order of a reaction is the sum of the orders of the reactants in the rate law
Isomerization of cyclopropane, C3H6
C3H6 (g) CH2=CHCH3 (g)
Rate = k [C3H6]
Reaction is first order with respect to C3H6
Reaction is first order overall
13.1­13.3
Rates of reaction
Reaction Order
2 NO (g) + 2 H2 (g) N2 + 2 H2O (g)
Rate = k [NO]2 [H2]
Reaction is second order with respect to NO
Reaction is first order with respect to H2
Reaction is third order overall
H+
CH3COCH3 (aq) + I2 (aq) CH3COCH3I (aq) + HI (aq)
Rate = k [CH3COCH3] [H+]
Reaction is first order with respect to CH3COCH3
Reaction is first order with respect to H+
*Reaction is zero order with respect to I2
Reaction is second order overall
*[I2]0 = 1
as long as some I2 is present,
reaction rate is not affected!
Orders are typically 1 or 2 (can be 0, 3)
Can also be negative or fractional, but very rare
Determining the Rate Law
Experimentally! find the order with respect to each reactant and any catalyst
Initial­rate method: run a series of experiments with the starting reactant concentrations varied and compare the rates
Helpful Table to Know!
13.1­13.3
Rates of reaction
Initial­rate method:
Rate = k [H2O2]x [I­]y [H+]z
Rate2
=
Rate1
[H2O2]2
[H2O2]1
2.30 x 10­6
=
1.15 x 10­6
x
0.020
0.010
x
2 = 2x
x = 1
Rate = k [H2O2]x [I­]y [H+]z
x =
y =
z = 13.1­13.3
Rate = k [H2O2] [I­]
Rates of reaction
Find k!
Rate = k [H2O2] [I­]
k = Exp. 2
13.1­13.3
k = 2.30 x 10­6
(0.020) (0.010)
Rate [H2O2] [I­]
= 1.2 x 10­2 L/(mol s)