Kinetics
Topic 6
• Rate of reactions
– can be fast or slow
– defined as the change in [concentration]
of the reactants or products per time
• concentration of reactants decreases over
time
– [reactants]/time
• concentration of products increase over
time
– [products]/time
• normally measure in Molarity per sec (M/s)
– will continue until reaches equilibrium or
one of the reactants is used up
concentration of reactants decrease over time
concentration of products increase over time
(reactants)
Experiments for measuring rate of reactions
– over time, could measure the change of:
• for gaseous reactions
– mass would go down as gas escapes
– volume would increase at constant temp. and
pressure
– pressure would increase at constant temp. and
volume
• change in pH if acids and bases are involved
• change in temperature
• change in electrical conductivity
– if produces ions in solution, conductivity will increase
• using a spectrometer to detect color changes
Analyzing data from rate experiments
(CONCENTRATION, VOLUME, and MASS)
• usually involves a graph of properties over
time
• usually a curve, and the reaction rate is
determined from the slope of the line at a
time (also known as a tangent)
• reaction rates tend to slow with time as
reactants are converted to products
Example
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
• reaction slows
down with time
because the
CONCENTRATION
of the reactants
decreases
• “rise over run”
– .040M/200s
= .0002M/s
– .025M/400s
= .000063M/s
• the change in concentration of a reactant or
product per unit of time
• [ ] refer to the concentration of the
reactants or products
[ A] at time t 2  [ A] at time t1
Rate 
t 2  t1
[ A]
Rate 
t
2NO2(g)  2NO(g) + O2(g)
product
[NO2]
[ NO2 ]
 constant
t
t
[NO2]
t
product
reactant
reaction is creating gas so
gas is being released in the
VOLUME increases over
reaction so MASS
decreases over time
time
Kinetic theory
– the average energy of particles is
proportional to the temperature (Kelvin or
Celsius)
• all particles have same energy if the same
temperature
– ex. 4He 20Ne 40Ar
»would all have the same temp. and energy
– however, lighter particles (4He) would have
greatest speed than larger particles given the
same energy
– 𝐾𝐸 = 1/2𝑚𝑣2
Activation Energy Ea
• a minimum amount of energy required for reaction
to occur
– bonds need to be broken first
• the molecules must posses sufficient energy to
get over the activation energy barrier.
Collision theory
• in order for particles to react
1. particles must collide
2. must collide in the correct orientation/angle
3. must collide with enough kinetic energy to
overcome the activation energy (E > Ea)
• if the previous conditions are “enough”,
particles can overcome the activation energy
and reaction will occur
– meaning the bonds holding the reactants together
will be broken and the reaction can take place
http://phet.colorado.edu/en/simulation/reac
tions-and-rates
Factors That Affect Reaction
Rates
• any factor that increases the frequency of
collisions or increases the energy with which
particles collide will make the reaction go
faster:
1.
2.
3.
4.
5.
temperature
pressure (gas only)
surface area
concentration
catalysts
1. Temperature
• increase temp
– increases number of collisions per unit time
• reaction rate approximately doubles for each
10oC or K rise in temperature
– increases energy of the collisions
• greater chance E > Ea
– will not “help with” geometry
2. Pressure
• only for gasses
– reducing volume while keeping temp
constant increases the reaction rate
• forcing them together will increase number of
collisions per unit time
– will not “help with” geometry and energy
3. Surface area
• smaller particles have more surface area and
will increase the reaction rate
– more collisions per unit time
– will not “help with” geometry and energy
4. Concentration
• increasing concentration will increase the
reaction rate by increasing more collisions per
unit time
• will not “help with” geometry and energy
5. Catalysts
• lowers the activation energy (Ea) for the
reaction
• provide an alternate reaction (rxn)
pathway
• increase the rate of a reaction
• are not used up or chemically changed in
the reaction
Maxwell–Boltzman energy distribution curve
• another way to look at how many particles can react
• area under the curve shows the number of gas
particles
– the # of particles remains constant
• not all gas particles have the same energy
– only some gas particles (blue area) have enough
energy to react
The affect of temperature
• the area under the curve remains the same
because the number of particles doesn't change
• higher temps. shifts the curve to the right
(therefore, the peek must be lower) resulting in
an increase in collision frequency and thus more
successful collisions
The affect of a catalyst
Never move the new activation energy to the left of the peak.
Catalysts don’t help out that much!