Transcript Kinetics. Topic 6

Kinetics
Topic 6
http://www.wwnorton.com/colle
ge/chemistry/gilbert2/chemtours.
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• Rate of reactions
– can be fast- explosions
– can be slow- rusting
– as a reaction takes place
• concentration of reactants decreases
(negative) over time
• concentration of products increase
(positive) over time
– normally measure in Molarity per sec (M/s)
– will continue until reaches equilibrium or
one of the reactants is used up
Measuring rate of reactions
– can measure:
• mass or volume change 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 electrical conductivity
– if produces ions in solution, conductivity will increase
• using a spectrometer to detect color changes
Determining Rate of Reaction from
reactions (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
Rate 
[ A ] at time t 2  [ A ] at time t1
t 2  t1
Rate 
[ A]
t
2NO2(g)  2NO(g) + O2(g)
product
[NO2]
t
[ N O2 ]
t
product
reactant
 constant
reaction is creating gas so
gas is being released in the
VOLUME increases over
reaction so MASS
decreases over time
time
• Kinetic theory (6.2.1)
– energy of particles is proportional to the
temperature (Kelvin or Celsius)
• all particles have same energy if the same
temperature
• lighter particles would have greater 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 (Topic 6.2)
• in order for particles to react
– particles must collide
– must collide in the correct orientation/angle
– must collide with enough kinetic energy to
overcome the activation energy (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
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
surface area
concentration
catalysts
1. Temperature
• increase temp
– increases number of collision per unit time
• reaction rate approximately doubles for each
10oC or K rise in temperature
– increases energy of the collisions
2. Pressure
• only for gasses
– reducing volume while keeping temp
constant
• forcing them together will increase number of
collisions
3. Surface area
• smaller particles have more surface area
– only the particles on the surface can come in contact
with a reactant
– more collisions per unit time
4. Concentration
• increasing concentration will increase more
collisions per unit time
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 what particles can react
• area under the curve shows the number of gas
particles
• 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!