Chapter 14 – Kinetics
I. Introduction
Chemical Kinetics: The area the chemistry concerned with:
 _______________________________
 _______________________________
 _______________________________
Thermodynamics taught us if a reaction will occur or not.
However, even if it is spontaneous, it might not happen at a measurable ______.
Rate =
Rate Example:
II. Reaction Rates
A. Definition of Reaction Rates
[A] = __________________
The reaction rate is measured based on:

Rate of increase in
 OR Rate of decrease in
 Usually expressed in units of moles per liter / unit of time
(i.e., M/min or M/sec)
Concentration changes:
 Reactant concentrations go __________ over time.
 Eventually they come to a steady state.
o This concentration is not necessarily zero.
 Product concentrations go __________ over time.
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Example Reaction
2 HCl(g)  H2(g) + Cl2(g)
Time [HCl],
(hours) M
0.00 0.200
2.00
0.130
4.00
0.086
∆
[HCl]
 [H2],
 [Cl2]
Current
[H2] OR
[Cl2]
0.000
To get the ∆[ ] take [ ]f – [ ]i
For time interval 0-2mis:
∆[HCl] =
∆[ H2] =
[H2] =
To get the slope, divide ∆[ ] by ∆time
For time interval 0-2mis:
HCl slope =
H2 slope =
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HCl
slope
H2 OR Cl2
Slope
 The [HCl] decreases with time, as predicted.
 The concentration of the products H2 and Cl2 increase with time.
Rate = A
time
slope = concentration
time
THE RATE IS RELATED TO THE ______________ OF THE GRAPH.
Notice that:
 The rate of decrease of the reactant (HCl) is twice that of the production of the
product (H2 or Br2)
 Rates of production of products are _________________. (conc is +)
 Rates of use of reactants are _________________ . (conc is -)
o The negative may be given in words (concentration goes down by …)
B. Determining Reaction Rates
Avg. Rxn. Rate = a positive quantity that depends on time, but not on the species
(product or reactant) and not on the coefficients of the equation.
Avg. Rxn. Rate =
Avg. Rxn. Rate =




Use a (_-_) in calculating the rate of disappearance of a reactant
 Always refers to final - initial
You should always calculate a positive answer
1/coeff normalizes the results so that you get the same result regardless of which
reactant or product you pick.
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The rate changes as the reaction proceeds.
1. Specify the time.
2. Reaction rates __________________ as the reaction mixture runs out of
reactants.
If we know the rate of appearance of any product or rate of disappearance of any
reactant, we can calculate the other rates based on the stoichiometry of the balanced
reaction.
LP# 1.: For the hypothetical reaction:
A + 3B  2C,
the average reaction rate is 12M/min
a) What is the rate of disappearance of A?
A is disappearing at _________________
b) What is the rate of disappearance of B?
This can be solved either mathematically:
Or it can be solved logically: Since stoichiometrically 3 mole B react with 1 mole
B, it gets used up 3 times faster! = ______________
b) What is the rate of appearance of C?
Since 2 moles of C are produced for every 1 mole of A used up
___________________________
OR Since 2 moles of C are produced for every 3 mole of B used up
___________________________
If the rate of appearance or disappearance of any component is known, we can
determine the rate of the others using stoichiometric ratios.
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Problem:
H2O2 can be used as a disinfectant; it decomposes as:
2 H2O2  2 H2O + O2
If the rate of appearance of O2 is 0.0014 Ms–1, (another way of writing M/s)
what is the rate of disappearance of H2O2?
a) -0.0014 Ms–1
b) -0.00070 Ms–1
c) -0.0028 Ms–1
d) None of the above are correct
Relationship between Rates and Equilibrium
What is Equilibrium?
 All reactions that were studied in the first semester of general chemistry were
assumed to go to completion: 100% reactants turned into 100% products.
 This semester, we will be a little bit more sophisticated.
 Many reactions are ______________________.
 After the forward reaction (reactants  products) gets partway complete,
the reverse reaction (products  reactants) starts up.
 The reaction get stuck somewhere between reactants and products
Equilibrium occurs when concentrations stop changing.
Equilibrium is defined as the point where the
__________________________
__________________________
(We will study this in much greater detail in the
following chapter!)
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C. Ways to Express Rates of Reactions
All data for [reactants] vs time basically resemble the following graph.
Average Rate
This is the type of rate calculation that we started with = _______________
This can be somewhat erroneous for large ∆t values.
Instantaneous Rate
= ______________________________________.
Initial Rate
= ______________________________________.
Why are we interested in initial rates?
 Many reactions go in both directions (forward & reverse).
 Only at t=0 (all reactants, no products) does the reaction proceed only in the
forward direction. (NO reverse reaction.)
 Using initial rates allows us to study the kinetics of ________________________
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D. Factors Affecting Reaction Rates
Reactants must be in contact with each other (collide) in order to react.
This requires effective collisions.
 These have _______________________
More to come on
this later!
 These have _______________________
Concentration
 Higher conc. means ___________particles in the _________ volume
 The # collisions .
 Usually, as [reactants] ____________ the rate ____________
Why do Hospitals make sure that no flames are allowed near patients receiving oxygen?
Flammable materials burn faster in pure oxygen than in air because the concentration of
O2 is greater.
 Occasionally, changes in [reactants] do not affect the rate of reaction.
Temperature
 Usually, as temp. ____________ the rate ____________
 The rate goes up ___________________ as the temperature rises!
 Rule of Thumb: reaction rate doubles for every ______ rise in temperature.
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Surface Area (for solids and liquids)
Reality Check: Small pieces of wood burn more quickly than large pieces. That is why
we use them to start fires.
 Reactions occur at the surface of a solid or a liquid-liquid or a liquid-gas interface.
Simplified reaction for the burning of wood:
C(s) + O2(g)  CO2(g) + heat
 For a reaction to occur there must be a
collision between ___________ .
 Oxygen can’t collide with C when it’s in
the middle of the large piece.
 Higher surface area leads to more collisions.
 More collisions lead to higher reaction rates.
 More surface area leads to _________________ Rxn Rates
For Liquids
Spray nozzles create small droplets with greater surface area.
(aerosols vs bulk liquids)
Example: burning bulk vs sprayed hair spray.
Catalysts
Definition: A substance that speeds up a chemical reaction but itself is unchanged in the
process.
 Addition of a catalyst will _________________ the rate of a reaction.
 The greater the conc. of the catalyst, the ________________ the effect.
They:
 Provide an easier _________________ for the reaction to occur.
 Lower the ____________________ energy(to be defined later)
 Still make the same products
 ______________________ are biological catalysts.
 The opposite of a catalyst is _____________________________.
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The catalyst generally does not appear in the overall balanced chemical equation,
although its presence is often indicated by writing its formula over the arrow.
Demo: Alka Seltzer tablet in a film canister.
a) ¼ tablet in ½ closed canister of water. Time as a baseline measurement.
b) 2 x ¼ tablet at the once.
this shows the effect of __________________________
c) ¼ tablet in ice cold water.
this shows the effect of __________________________
d) ¼ tablet broken in pieces:
this shows the effect of __________________________
III. Reaction Rate Laws
Rate law - states the dependence of the reaction rate on _____________________
A. General Reaction
c
aA + bB  xX + yY
(c is the catalyst if there is one)
General Rate Law
Rate = _____________________________
k = _________________________ (a proportionality constant)
(a function only of T) (like the equilibrium constant K)
The exponents (m, n, p) ______________________________
the stoichiometric coefficients.
The exponents (m, n, p) must be ___________________ determined.
 k has units. (They may be bizarre and depend on the form of the rate law.)
 k changes with _____________________.
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B. Reaction Order
Values of m and n and p indicate the reaction order with respect to A, B, and C.
 exponent = 0; _____ order with respect to that reactant.
 exponent = 1; _____ order with respect to that reactant.
 exponent = 2; _____ order with respect to that reactant.
 exponent = 3; _____ order (possible but very uncommon)
Overall reaction order = _________ of exponents (e.g.= m + n + p)
Usually small positive integers but can be negative, zero, or even fractions.
2 N2O5(g)  4 NO2(g) + O2(g)
Example Reaction #1
Experimentally determined rate law:
rate = k [N2O5]1
(the one is usually not written)
The reaction is _____ order with respect to N2O5
Overall Reaction Order = _____ order
Example Reaction #2
2 NO(g) + Cl2(g)  2 NOCl(g)
Experimentally determined rate law:
2
Rate = k [NO] [Cl2]
(exponents only match coeff. by accident.)
Reaction order: ______ with respect to NO
______ with respect to Cl2
______ overall
Relationship between Rate and Reaction Orders
Change in reaction rate when the concentration of a reactant A is doubled or tripled for
different values of the exponent m in the rate law, rate = k[A]m[B]n
Value of m
0
1
2
3
Reaction Rate if
[] is doubled
20 = 1
(Unchanged)
21 = ___ faster
22 = ___ faster
23 = ___ faster
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Reaction Rate if
[] is tripled
30 = 1
(Unchanged)
31 = ___ faster
32 = ___ faster
33 = ___ faster
LP# 2.:
The reaction: 2 NO(g) + 2 H2(g)  N2(g) + 2 H2O(g)
is first order in H2 and second order in NO.
a) Write the rate law.
________________________________________
b) What is the overall order of the reaction?
____________________
c) How does the reaction rate change if the concentration of H2 is doubled and the
concentration of NO is tripled?
change in reaction rate = __________________
d) How does the reaction rate change if the concentration of NO is cut in half and
the concentration of H2 is held constant?
the rate will ___________________________
Now It’s Your Turn!
Problem:
What is the order of the reaction with respect to OH– for
2 ClO2(aq) + 2 OH–(aq) → ClO3–(aq) + ClO2–(aq) + H2O(l)
Based on the following data:
[OH–]
0.030
0.090
a)
b)
c)
d)
e)
Initial Rate
(M/s)
0.0027
0.0083
0
1
2
3
more information is needed.
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Let’s Stop and Review Logs Before Proceeding
A base 10 log tells us the exponent of a power of 10 or how many times 10 is multiplied
by itself.
If
y=10x
Example:
y = 104
If
Then log y = log 104 = _____
Then log y = _____
Similarly for natural logs
If
y=ex
Example:
y = e4
If
where e ~ 2.718
Then ln e4 = _____
Then ln y = _____
Note: Log and 10x cancel out. Also ln and ex cancel out.
Multiplication Rule
log (x)(y) = ____________________________
Example:
log (103)(102) = log 103 + log 102 = 3 + 2 = 5
log (105)
=5
Subtraction Rule
log
x
= ______________________________
y
Example:
log
105
= log 105 – log 102 = 5 – 2 = 3
2
10
log 103
=3
Powers Rule
log (ym) = _____________________________
Example:
log (102)3 = 3 log 102 = 3 x 2 = 6
log (1006) = 6
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