1.
2.
3.
4.
5.
6.
7.
atomic #, mass #, “carbon-14
Review stoichiometry
Molar Mass
Calculating Limiting & Excess reactants
Percent Yield & Percent Composition
Getting Empirical formula
Getting Molecular formula from Empirical
1.
Begin new unit on kinetics
HOMEWORK
- Page 567 # 25 - 29
- Continue reviewing for May’s exam
Chemical Kinetics
12.1 – The Rate of a Reaction
What is kinetics?
Rate
of reactions
Factors that affect rates
reaction rate can be studied by
 observing
is used up
 or
rate at which a reactant
the rate with which a product is
formed
reaction rate can be studied by
 How
to measure amount of reactant
lost/product formed?
 What
units?
◦pH
◦color
◦concentration?
change in amount
change in time
= average rate
Rates
of reaction change
during the progress of
the reaction
Therefore,
when we talk
about the rate of a reaction
we need to be very specific
Initial Rate

rate at time just after mixing, when the reaction
first starts
Instantaneous Rate

The instantaneous rate is the rate at any given time
Average Rate

The average rate is the rate during a given time
interval
Increasing & Reducing Speed of Reaction
Frequency or number of collisions
 Frequency or number of collisions

with enough
Ea
Ea = Activation Energy

Collisions with enough Ea

Favorable orientation results in reaction
Factors affecting Reaction Rates
Surface
Area
Concentration
State(s) of reactant
Temperature
Catalyst
Factors affecting Reaction Rates



Surface Area: The greater the surface area of a
substance, the greater the rate of reaction.
Larger surface area increases the frequency of
collisions
Example: Burning a log vs. burning wood
splints
Factors affecting Reaction Rates



Concentration: Increasing
concentration increases the rate of a
reaction (aqueous solutions and
gases only)
Increase in concentration increases
the frequency of collisions
Example: fanning a fire.
Factors affecting Reaction Rates
 In
general, rate of
reactants fastest for
(aq) > (g) > (l) > (s)
 Homogeneous
vs.
Heterogeneous
Factors affecting Reaction Rates
 Temperature:
◦Increasing temperature
generally increases the
rate of a reaction.
 WHY?
 Activation Energy
Factors affecting Reaction Rates
 Catalysts:
increases the rate of
reactions by lowering the activation
energy. The catalyst itself is not
used up in the reaction.
 Example:
driving through a tunnel
vs. driving over a mountain.
Enzymes
Ea = Activation Energy

NO reaction due to orientation

Favorable orientation results in reaction
Reaction Rates
1. Does every collision between reaction particles
lead to products? What other factors are
involved?
2. How does each of the following affect the rate
of a chemical reaction?




3.
Temperature
Concentration
Surface Area
Catalyst
Food is stored in refrigerators have a slower
rate of decay. Why?
Using numbers/math to describe how fast a reaction is
2 N2O5(g)  4 NO2(g) + O2(g)
Rate = [O2]/t
Rate = - [N2O5]/t
the negative sign indicates the
reduction in concentration of reactant
with time
[O2]/t = -1/2 [N2O5]/t
Br2(aq) + HCOOH(aq)  2Br-(aq) + 2H+(aq) + CO2(g)
Rate = [CO2]/2t
Rate = - [Br2]/t
Br2(aq) + HCOOH(aq)  2Br-(aq) + 2H+(aq) + CO2(g)
Rate is proportional to
concentration of Br2
rate = k[Br2]
k = rate constant, a constant between the reaction
rate and the concentration of reactant
 Like
the “m” in y = mx +b
Write the rate expressions for the following reactions
I (aq)
+
4NH3(g)
OCl (aq)

Cl (aq)
+
OI (aq)
+ 5O (g)  4NO(g) + 6H O(g)
2
2
4NO2(g) + O2(g)  2N2O5(g)
Suppose that, at a particular moment during the
reaction, molecular oxygen is reacting at the rate
of 0.024 M/s.
 At
what rate is N2O5 being formed?
 At what rate is NO2 reacting?
4PH3(g)  P4(g) + 6H2(g)
Suppose that, at a particular moment during the
reaction, molecular hydrogen is being formed at
the rate of 0.078 M/s.
At
what rate is P4 being formed?
At what rate is PH3 reacting?
Rate Law helps to
determine rate regardless
of starting concentration
rate = k
[reactant]x[reactant]y…
rate law
k = rate constant
12.2 Reaction rate & Reactant Concentration
rate = k
[reactant]x[reactant]y…
F2(g)+ 2ClO2(g)  2FClO2(g)
rate law is: k [F2][ClO2]
rate of reaction is:
x
y
rate = k[F2] [ClO2]
rate of reaction is:
x
y
rate = k[F2] [ClO2]
rate = k[F2]1
rate = k[F2]
rate of reaction is:
1
y
rate = k[F2] [ClO2]
rate = k[reactant]0
What does it mean for
of
x
to have a value
0, 1, or 2?
0 = Zero Order
reaction rate does not change with
concentration.
rate = k[reactant]1
1 = First Order
reaction rate is DIRECTLY proportional
to concentration
example: previous problem
double F2, rate doubles
increase ClO2 by 4, rate increases by 4
rate = k[reactant]2
2 = Second Order
reaction rate is proportional to the
SQUARE of concentration
example:
change concentration from 1.0 M to 2.0
M, rate increases by 4
increase concentration to 3.0 M, rate
increases by 9
Identify the order for the following
The units for rate is
M/s or mol/L*s
k has different units for different reaction orders
Zero Order: k = M/s or mol/L*s
First Order: k = s-1
Second Order: k = M-1*s-1 or L/mol*s
MEMORIZE?
not necessarily
k will have whatever units needed
to get rate as M/s or mol/L*
2NO(g) + 2H2(g)  N2(g) + 2H2O(g)
 1.
determine the rate law
 2. Determine the rate constant with
proper units
 3. Determine when rate when [NO] =
12.0 x10-3M and [H2] = 6.0 x 10-3M
Summary & Review
1.
2.
3.
4.
5.
What is the function of the rate law?
How do you write the rate law?
How to use initial rates to determine
reaction orders, k, and rate for given
concentrations.
What the reaction orders mean?
Recognizing graphs for each order.
12.3 Reaction rate vs Time
Why another law?
Rate Laws
show
relationship
between
rate
&
concentrations.
Why another law?
Integrated rate law
concentration
and time
How much used?
How much left?
First Order
ln[A]t = -kt +ln[A]0
Decomposition of cyclopropane to propene
is 1st order
CH2CH2CH2  CH3CHCH2
At 500 C, k = 6.7 x 10-4 s-1
Initial concentration of cyclopropane 0.25M
1.
2.
3.
What is the concentration of cyclop after 8.8
minutes?
How long (in minutes) for concentration of
cyclop to decrease from 0.25 M to 0.15 M?
(c) How long (in minutes) will it
Decomposition of cyclopropane to propene
is 1st order
CH2CH2CH2  CH3CHCH2
At 500 C, k = 6.7 x 10-4 s-1
Initial concentration of cyclopropane 0.25M
1.
How long (in minutes) to convert 74% of
starting material to propene
First Order & Half-Life
Half-Life
time it takes to decrease
concentration
by half
First Order & Half-Life
ln[A]t = -kt +ln[A]0
thalf not effected by concentration!
thalf
0.693
=
𝑘
First Order & Half-Life
C2H6(g)  2CH3(g)
@ 700 C, rate constant is 5.36 x 10-4 s-1
1.
Calculate the half-life of the reaction in
minutes.
First-Order
k [A]
Rate Law:
Integrated:
ln[A]t = -kt +ln[A]0
y
Half-Life:
=
mx +
b
Second Order
Integrated Rate Law:
1
1
=kt +
[A]t
[A]0
Second Order
Integrated Rate Law:
Half-Life:
1
1
= kt +
[A]t
[A]0
thalf
1
=
k [A]0
Second Order
Integrated Rate Law:
1
1
= kt +
[A]t
[A]0
I(g) + I(g)  I2(g)
rate constant @ 23 C is 7.0 x 109 M s-1
1.
What is concentration of I after 2 minutes
if initial is 0.086M?
Second Order
I(g) + I(g)  I2(g)
rate constant @ 23 C is 7.0 x 109 M s-1
1.
How long for initial concentration of 0.60
M to reduce by half?
thalf
1
=
k [A]0
Second Order
I(g) + I(g)  I2(g)
rate constant @ 23 C is 7.0 x 109 M s-1
1.
How long for initial concentration of 0.42
M to reduce by half?
thalf
1
=
k [A]0
Second Order
A(g) + A(g)  A2(g)
rate constant @ 25 C is 51.0 M
1.
2.
-1min-1
How long for initial concentration of
0.0092 M to reduce to 3.7 x 10-3 M?
What is the half-life of the reaction?
Second Order
Integrated Rate Law:
1
1
= kt +
[A]t
[A]0
Half-Life:
thalf
1
=
k [A]0
Zero Order
Integrated Rate Law:
[A]t = -kt +[A]0
Zero Order
Integrated Rate Law:
Half-Life:
[A]t = -kt +[A]0
thalf =
[A]0
2k
Zero Order
Integrated Rate Law:
[A]t = -kt +[A]0
Half-Life:
thalf =
[A]0
2k
12.4 View of Reactions at Atomic Level
Ea = Activation Energy

NO reaction due to orientation

Favorable orientation results in reaction
2NO + O2  2NO2

Mechanism 1

Mechannism 2
2NO + O2  2NO2
N2O2 = Intermediate
(product first, then reactant)
 Rate
law for the overall reaction
determined by the slowest
elementary step
 Rate Determining Step (RDS)
Mechanism 1: N2O  N2 + O
Mechanism 2: N2O + O  N2 + O2
What is the overall reaction?
Identify intermediates
Write the rate law for overall reaction
slow
fast
Mechanism 1: 2NO2  NO + NO3
slow
Mechanism 2: NO3 + CO  NO + CO2 fast
What is the overall reaction?
Identify intermediates
Write the rate law for overall reaction
H2(g) + I2(g)  2HI(g)
Mechanism 1: I2  2I
Mechanism 2: H2 + 2I  2HI
Overall:
H2(g) + I2(g)  2HI(g)
fast
slow
I is intermediate
rate law = k2[H2][I]2
𝑘1𝑘2
rate law = ′ [H2][I2]
𝑘1