# Kinetics

Kinetics/Equilibrium
A unit dedicated to studying the
rate of chemical reactions (how
quickly reactions take place)
Spontaneous processes
• A spontaneous process is a chemical
reaction in which a system releases free
energy (most often as heat) and moves to
a lower, more thermodynamically stable,
energy state.
– Examples:
• Burning a log after igniting the wood.
• Rusting
• A ball falling through the air.
What ultimately determines if a
reaction is spontaneous or not?
• ΔG = ΔH - TΔS
ΔG = Change in Gibbs Free Energy
ΔH = Enthalpy change
T = temperature of the system (Kelvin)
ΔS = entropy change
If ΔG is negative, the reaction is spontaneous.
Factors affecting spontaneity?
• When ΔS is positive and ΔH is negative, a process is spontaneous
• When ΔS is positive and ΔH is positive, a process is spontaneous at
high temperatures, where exothermicity plays a small role in the
balance.
• When ΔS is negative and ΔH is negative, a process is spontaneous
at low temperatures, where exothermicity is important.
• When ΔS is negative and ΔH is positive, a process is not
spontaneous at any temperature, but the reverse process is
spontaneous.
Examples of entropy changes –
you decide??
• Phases changes from solids  liquids 
gases?
• Chemical reactions in which the moles of
gases produced is greater than the moles
of gases used? Example: C(s) + O2(g)  2CO(g)
• Dissolving a solid into a liquid solvent to
produce ions?
• Dissolving a gas into a liquid solvent?
• Molecular complexity? NaCl  CaCl2  AlCl3
Homework:
• Read Chapter 8 – section 1
• Answer questions on pages 262-263,
questions 2,4,6,8,9
Reaction rates
• The rate of a reaction is the same as the
rate of any other process – it is a measure
of the change in something over time.
– For a moving object, rate (speed) is the
change in the position of the object over time.
– For reactions, it is a change in the
concentration of something over time.
What does the plot of reaction rate vs. time look
like?
NOTE: What happens to the rate of a reaction as the
concentration of the reactants decreases?
NOTE: What happens to the concentration of the reactants
as the reaction proceeds?
To find Instantaneous rate at any given point of the
reaction, simply…
• Find the tangent
of the curve at
any given point on
the plot of
reactant
concentration vs.
time.
Factors that effect the rate of a
reaction:
• 1) The initial concentration of the reactants (the
greater the concentration, the faster the rate of the
reaction.)
• 2) The presence of a catalyst (A catalyst is something
that lowers the activation energy of the reaction, and
therefore speeds up the rate of the reaction.)
• 3) The temperature of the reactants (at higher
temperatures, the reactants are moving faster and are at
a higher energy state, so there is a greater rate which
collisions result in products.
• 4) The surface area of the reactants (The greater the
surface area, the faster the reaction.)
What theory explains these
factors?
• &quot;collision theory&quot; - which states that for
molecules to react, they must:
– collide
– have the right energy
– have the right geometry
To increase the rate, you must make the above more likely to occur. This is
possible by changing other factors such as:
1) increasing the surface area (of solids)-this allows for more collisions and
gives more molecules the right geometry
2) increasing the temperature-this gives more molecules the right energy
(also called the activation energy, Ea)
3) increasing the concentration (of gases and solutions)-this allows for
more collisions and more correct geometry
4) using a catalyst-helps molecules achieve the correct geometry by
providing a different way to react
How do catalysts speed up
reactions???
• Catalysts move the energy barrier to the left as noted earlier – or,
more significantly, they lower the activation energy of the
reaction.
Rate Laws?
• for the general reaction of A
and B to give products E and
F:
• a A + b B --&gt; e E + f F
(where a, b, e, and f are
stoichiometric coefficients) the
rate is defined as
• R = -d[A]/a dt = -d[B]/b dt = +
d[E]/e dt = + d[F]/f dt
• Thus for the general
reaction of A and B
considered to the left, we
might find
• R = -= k [A]a[B]b
where the exponents a
and b need not be
integers or half-integers
and are not necessarily
equal to the
stoichiometric
coefficients a and b.
Can you arrive at the correct rate
law expression?
• Example Problem: Find the rate law and rate
constant of A + B --&gt; C using the following data
Another example!!
•
•
•
•
•
How do we determine the rate law from initial rates?
[A]
0.100
0.100
0.200
[B]
0.100
0.200
0.100
Initial rate
4 x 10 -5 M/s
4 x 10 -5 M/s
1.6 x 10 -4 M/s
The rate law for this data is: Rate = k [A]2[B]0
• The general rate law for any reaction is:
Rate = k [reactant 1]m[reactant 2]n
With m &amp; n the reaction orders for each reactant with
their sum the overall reaction order.
What is meant by the order of a
reaction?
• Reaction order is a topic which comes with reaction rates. If you
have a reaction in which A, B, and C are possible reactants, then we
can describe the order of the reaction following this chart.
Homework???
• Complete the worksheet given in class.
What is meant by overall reaction
molecularity?
consider the following reaction:
NO(g) + O3(g)  NO2(g) + O2(g)
• If NO molecules collide with ozone
molecules, we will produce nitrogen
dioxide and oxygen gas. (Note: This
equation shows the decay of ozone from
nitrous oxide – therefore NO is an ozone
depleting gas!!)
The previous reaction is bimolecular – (a result of a collision of 2
reactant molecules)
New term – molecularity – the number of molecules that participates as
reactants in the elementary step of a reaction.
Example: The decomposition of a peroxide into oxygen and water would
be unimolecular.
Example: reactions involving the simultaneous collisions of three
molecules would be called termolecular (very rare – it is hard to
get three molecules to collide with the right orientation)
Summarize:
A products
A + A  products
A + B  products
A + A + A  products
A + A + B  products
A + B + C  products
unimolecular
bimolecular
bimolecular
termolecular
termolecular
termolecular
The previous reaction is called an elementary process (because it
occurs in one elementary step) – Not all reactions are elementary
processes.
• Consider: NO2(g) + CO(g)  NO(g) +
CO2(g)
The actual reaction is the following:
NO2 + NO2  NO3 + NO
NO3 + CO  NO2 + CO2
• Note: If we add these two elementary steps up, we get the above
net reaction.
• NO3 would be called an intermediate – it is used up as soon as it is
formed.
• Also, it is important to note that each of the elementary steps has
there own rate constant and there own rate law. The slower step
would be the rate determining step.
How do we know a reaction
mechanism is needed to explain
the reaction???
• Any time the rate law expression
generated from the balanced equation
does NOT agree with what is observed
through experimentation, you need to
consider a reaction mechanism!!
Homework???
• Complete the worksheet given in class!!
A cool on line lab!!
• Lab #1 – persulfate ion and iodide ion
reaction