Transcript Reaction Rate - DocLockert.com

Ch 17 Reaction Rates &
Ch 18 Equilibrium
Part 1. How Fast Is the Reaction?
Collision Theory
 In
order to react, molecules and atoms
must touch each other.
 And they must hit each other hard
enough to react.
 Anything that increases these things will
make the reaction faster.
 There is a certain amount of energy
needed to start the reaction. This energy
is called the activation energy.
Collision Theory
 Collision
theory - atoms,
ions, and molecules must
collide in order to react.
 For any reaction to occur,
the particles must come
in contact with each
other.
 Only a small amount of
collisions produce
reactions.
Orientation

Even though particles collide, the particles must
have the correct orientation to react.

When correct orientation does occur a short
lived intermediate substance is formed is called
an activated complex

activated complex - a temporary, unstable
arrangement of atoms that may form products
or may break apart to re-form reactants.
Activation Energy
Even when particles collide and with correct
orientation, the reaction still might not occur
 An activated complex will not form if there is
insufficient energy.
 Activation energy (Ea) - the minimum amount of
energy that reacting particles must have to form
the activated complex and lead to a reaction.

Correct Orientation and
Energy
Energy
Reactants
Products
Reaction coordinate
Energy
Activation Energy Minimum energy to
make the reaction
happen
Reactants
Products
Reaction coordinate
Energy
Activated
Complex or
Transition State
Reactants
Products
Reaction coordinate
Energy
Reactants
Overall energy
change
Products
Reaction coordinate
Reaction Pathway
 Shows
the change in energy during a
chemical reaction
Exothermic Reaction
 reaction
that
releases
energy
 products have
lower PE
than reactants
energy
released
2H2(l) + O2(l)  2H2O(g) + energy
Endothermic Reaction
 reaction
that
absorbs
energy
 reactants have
lower PE
than products
energy
absorbed
2Al2O3 + energy  4Al + 3O2
Factors Affecting Reaction
Rates
Movie
Things that Affect Reaction Rate
Temperature is important.
 Higher temperature means faster particles.
 Faster means more and harder collisions.
 This leads to faster reactions.
 Concentration is important.
 More concentrated the reactants, the closer
together the molecules are.
 Closer molecules means colliding more often.
 Again leading to faster reactions.

Things that Affect Reaction Rate
 Particle
size can be a factor.
 Molecules can only collide at the surface.
 Smaller particles, bigger the surface area.
 Smaller particles means faster reactions.
 Smallest possible molecules or ions
result from dissolving compounds.
Dissolving speeds up reactions.
 Getting two solids to react with each
other is slow.
Things that Affect Rate
 Catalysts-
substances that speed up a
reaction without being used up.(enzyme).
 Speeds up reaction by giving the reaction
a new path.
Things That Affect Rate
 The
new path has a lower activation
energy.
 More molecules have this energy.
 The reaction goes faster.
 Inhibitor- a substance that slows or even
blocks a reaction.
The Nature of Reactants
The reactive nature of reactants can determine
the rate of reactions
 Some substances react more readily than
others such as the alkali metals.

For
example, both
calcium and sodium are
reactive metals,
however when placed in
water, sodium reacts
much more readily and
produces much more
energy than calcium.
Concentration
 Reactions
speed up when the
concentrations of reacting particles are
increased
 The
reaction speeds up because the
amount of collisions between to reacting
particles increases.
Surface Area
 When
the surface area increase the
reaction rate increases
 Increasing
the surface area allows more
particles to collide with one another, thus
increasing the reaction rate.
Temperature

Usually increasing the temperature of a
reaction generally increases the reaction rate.

This occurs because when temperature is
increased, the particles move more quickly thus
more collisions occur.

Also, when temperature is increased the
energy of the collisions is greater.
Energy
Reactants
Products
Reaction coordinate
Catalysts
H H
 Hydrogen
bonds to
surface of metal.
 Break H-H bonds
H
H
Pt surface
H H
H H
Catalysts
H
H
H
C
C
H
H H
H
H
Pt surface
Catalysts
 The
double bond breaks and bonds to the
catalyst.
H
H
H
C
H
C
H
H
Pt surface
H H
Catalysts
 The
hydrogen atoms bond with the carbon
H
H
H
C
H
C
H
H
Pt surface
H H
Catalysts
H
H
H
H
C
C
H
H
H
Pt surface
H
Catalytic Converter
Inhibitors
 Inhibitors
slow down reaction rates and
some stop reactions from occurring all
together.
 The
food industry uses inhibitors to keep
foods fresher, longer
 Example
is the chemical that is put on
apples to keep apples from browning.
Use an Inhibitor to Slow Down
the Reaction
Expressing Reaction Rates
Chemical reaction occur at certain rates.
 The reaction between vinegar and baking soda
occurs relatively fast where as the reactions
that occur in the formation of fossil fuels occurs
much slower
 reaction rate - the change in concentration of a
reactant or a product per unit time, expressed
as M/s.

Calculating Reaction Rates
 Average
reaction rate = Δ quantity / Δ time
 Δ quantity = final molar - initial molar
 Δ time = final time - initial time
 The value is never negative, so take the
absolute value
Sample Question
In a reaction between butyl chloride (C4H9Cl)
and water the initial concentration of butyl
chloride was 0.220M at time 0.00 s and the
concentration at time 4.00 s was 0.100M.
Calculate the average reaction rate over this
time period.
 Known: t1 = 0.00 s
C4H9Cl at t1 = 0.220M
t2 = 4.00 s
C4H9Cl at t2 = 0.100M


Unknown: average reaction rate = ? M/s
Answer
 Average
reaction rate = Δ quantity / Δ
time
 = 0.100 M - 0.220 M / 4.00 s - 0.00 s
 = |-0.120 M/4.00 s| note: never negative
 = 0.0300 M/s
Reaction Rate Laws
 The
equation that expresses the
mathematical relationship between the
rate of a chemical reaction and the
concentration of reactants is called the
rate law.
 Rate = k[A]m[B]n
 Rates are determined experimentally
 k is the specific rate constant and varies
with temperature.
Reaction Rate Laws
m
and n are reaction orders and define how
the rate is affected by the concentration
(molarity) of reactants A and B.
 The overall reaction order is equal to the
sum of the individual reactant reaction
orders.
 If doubling the concentration doubles the
reaction, then the order is first order or one,
ie [A]1
 If doubling the concentration quadruples
the reaction, then the order is second order
or two, [A]2
Overall Reaction Order
 We
can use the method of initial rates to
determine the overall reaction order.
 Examine the experimental data below to
determine the reaction order.
Overall Reaction Order
Recall the general rate law is
 Rate = k[A]m[B]n
 Looking at trial 1 and 2 we see doubling A
doubles the initial rate. Therefore the reaction is
1st order with A concentration (m=1).
 Looking at trials 2 and 3 we see doubling B
quadruples the initial rate. Therefore the reaction
is 2nd order with B (n=2).
 The overall rate is 3rd order which is the sum of m
and n.

Your Turn
 Determine
the general rate law formula
and overall reaction order based on the
experimental data below.
Reaction Mechanism
 Elementary
reaction- a reaction that
happens in a single step.
 Reaction mechanism is a description of
how the reaction really happens.
 It is a series of elementary reactions.
 The product of an elementary reaction is
an intermediate.
Reaction Mechanism
 An
intermediate is a product that
immediately gets used in the next
reaction.
Reaction Mechanisms

Many chemical reactions consist of a sequence
of two or more reactions

Such is evident in the earth’s stratosphere
where 2O3
3O2

This is the overall reaction after three steps
occur which is started when intense UV
radiation from the sun liberates chlorine atoms
from certain compounds.
Reaction
 Elementary
step: Cl + O3
 Elementary
Step: O3
 Elementary
Step: ClO + O
 Complex
reaction: 2O3
O2 + ClO
O2 + O
Cl + O2
3O2
Reaction Mechanisms
 The
destruction of ozone on the previous
slide is a complex reaction
 The complete sequence of elementary
steps that make up a complex reaction is
known as a reaction mechanism.
 In a reaction mechanism some
substances are produce and remain while
others known as intermediates are
produced by one reaction but consumed
in a subsequent reaction.

This reaction takes place in three steps

Ea
First step is fast
Low activation energy

Ea
Second step is slow
High activation energy

Ea
Third step is fast
Low activation energy
Second step is rate determining
Intermediates are present
Activated Complexes or
Transition States
Mechanisms and Rates
 There
is an activation energy for each
elementary step.
 Slowest step (rate determining) must
have the highest activation energy.
Thermodynamics Ch.16.5
Will a reaction happen?
Energy
 Substances
tend to react to achieve the
lowest energy state.
 Most chemical reactions are exothermic.
 Doesn’t work for things like ice melting.
 An ice cube must absorb heat to melt.
Why?
 Thermite Reaction movie
Entropy
 The
degree of randomness or disorder.
 S is the symbol of entropy.
 The first law of thermodynamics states
the energy of the universe is constant.
 The second law of thermodynamics
states the entropy of the universe
increases in any change.
 Drop a box of marbles.
 Watch your room for a week.
Entropy
 Defined
in terms of probability.
 Substances take the arrangement that is
most likely.
 The most likely is the most random.
 Calculate the number of arrangements for
a system.
2
possible
arrangements
 50 % chance of
finding the left
empty
4
possible
arrangements
 25% chance of
finding the left
empty
 50 % chance of
them being
evenly
dispersed
 16
possible
arrangements
 6.25% chance of
finding the left
empty
 37.5 % chance
of them being
evenly
dispersed
Gases
 Gases
completely fill their chamber
because there are many more ways to do
that than to leave half empty.
Ssolid <Sliquid <<Sgas
 there
are many more ways for the
molecules to be arranged as a liquid than
a solid.
 Gases have a huge number of positions
possible.
Entropy
Entropy
of a
solid
 A solid
Entropy
of a
liquid
Entropy
of a gas
has an orderly arrangement.
 A liquid has the molecules next to each
other.
 A gas has molecules moving all over the
place.
Entropy increases when...
 reactions
of solids produce gases or liquids, or
liquids produce gases.
 a substance is divided into parts -reactions
with less reactants than products increase in
entropy.
 the temperature is raised -the random motion
of the molecules is increased.
 a substance is dissolved. –increase in # of
pieces
 There are tables of standard entropy but our
book doesn’t have one. I will give them to you.
Entropy Calculations
Standard entropy is the entropy at 25ºC and 1 atm
pressure.
 Abbreviated Sº, measure in J/K.
 The change in entropy for a reaction is
DSº= Sº(Products) - Sº(Reactants)
 Determine if the reaction is positive or negative
entropy. If entropy increases then ΔS is positive.
 A solid changing to a gas is an increase in
entropy. Making more products or dissolving a
solid is a positive change as well.

Spontaneity
Will the reaction happen, and how
can we make it happen?
Spontaneous Reaction
 Physical
or chemical change that will
occur without outside intervention.
 Nonspontaneous reactions don’t.
 Even if they do happen, we can’t say how
fast.
 Two factors influence spontaneity,
 Enthalpy (heat) and Entropy (disorder).
Enthalpy and Entropy Factors
 Exothermic
reactions tend to be
spontaneous. *Negative DH.
 Reactions where the entropy of the
products is greater than reactants tend to
be spontaneous. *Positive DS.
 A change with positive DS and negative DH
is always spontaneous.
 A change with negative DS and positive DH
is never spontaneous.
Gibbs Free Energy
 The
energy free to do work is the change
in Gibbs free energy.
 DGº = DHº - TDSº (T must be in Kelvin)
 All spontaneous reactions release free
energy.
 A negative ΔGo like a negative ΔHo
represents a release in energy.
 So DG <0 for all spontaneous reactions.
DG=DH-TDS
DS DH DG
+ -
-
Spontaneous?
At all Temperatures
+ + ?
At high temperatures,
“entropy driven”
-
At low temperatures,
“enthalpy driven”
-
-
?
+ +
Not at any temperature,
Reverse is spontaneous
Reversible Reactions
 Reactions
are spontaneous if DG is
negative.
 If DG is positive the reaction happens in
the opposite direction.
 2H2(g)
+ O2(g)  2H2O(g) + energy
+ energy  2H2(g) + O2(g)
 2H2(g) + O2(g)
2H2O(g) + energy
 2H2O(g)
Equilibrium
 When
I first put reactants together the
forward reaction starts.
 Since there are no products there is no
reverse reaction.
 As the forward reaction proceeds the
reactants are used up so the forward
reaction slows.
 The products build up, and the reverse
reaction speeds up.
Equilibrium
 Eventually
you reach a point where the
reverse reaction is going as fast as the
forward reaction.
 This is dynamic equilibrium.
 The rate of the forward reaction is equal to
the rate of the reverse reaction.
 The concentration of products and
reactants stays the same, but the reactions
are still running.
Equilibrium
 Equilibrium
position- how much product
and reactant there are at equilibrium.
 Shown with the double arrow.

Reactants are favored

Products are favored
 Catalysts speed up both the forward and
reverse reactions so don’t affect
equilibrium position.
Measuring Equilibrium
 At
equilibrium the concentrations of
products and reactants are constant.
 We can write a constant that will tell us
where the equilibrium position is.
 Keq equilibrium constant
coefficients
 Keq = [Products]
[Reactants]coefficients
 Square brackets [ ] means concentration
in molarity (moles/liter)
Writing Equilibrium Expressions
 General
equation
aA + bB
 Keq
cC + dD
= [C]c [D]d
[A]a [B]b
 Write
the equilibrium expressions for the
following reactions.
 3H2(g) + N2(g)
2NH3(g)
 2H2O(g)
2H2(g) + O2(g)
Calculating Equilibrium
 Keq
is the equilibrium constant, it is only
effected by temperature.
 Calculate the equilibrium constant for the
following reaction
3H2(g) + N2(g)
2NH3(g)
if at 25ºC there are 0.15 mol of N2 , 0.25
mol of NH3 , and 0.10 mol of H2 in a 2.0 L
container.
What it tells us
 If
Keq > 1 Products are favored
 If Keq < 1 Reactants are favored
LeChâtelier’s Principle
Regaining Equilibrium
LeChâtelier’s Principle
 If
something is changed in a system at
equilibrium, the system will respond to
relieve the stress.
 Three types of stress are applied.
1a. Changing Concentration
 If
you add reactants (or increase their
concentration).
 The forward reaction will speed up.
 More product will form.
 Equilibrium “Shifts to the right”
 Reactants  products
1b. Changing Concentration
 If
you add products (or increase their
concentration).
 The reverse reaction will speed up.
 More reactant will form.
 Equilibrium “Shifts to the left”
 Reactants  products
1c. Changing Concentration
 If
you remove products (or decrease their
concentration).
 The forward reaction will speed up.
 More product will form.
 Equilibrium “Shifts to the right”
 Reactants  products
1d. Changing Concentration
 If
you remove reactants (or decrease their
concentration).
 The reverse reaction will speed up.
 More reactant will form.
 Equilibrium “Shifts to the left”.
 Reactants  products
 Used to control how much yield you get
from a chemical reaction.
2a. Changing Temperature
 Reactions
either require or release heat.
 Endothermic reactions go faster at higher
temperature.
 Exothermic go faster at lower
temperatures.
 All reversible reactions will be exothermic
one way and endothermic the other.
2b. Changing Temperature
 As
you raise the temperature the reaction
proceeds in the endothermic direction.
 As you lower the temperature the reaction
proceeds in the exothermic direction.
 Reactants + heat  Products at high T
 Reactants + heat  Products at low T
3. Changes in Pressure
 As
the pressure increases the reaction
will shift in the direction of the least
gases.
 At high pressure
2H2(g) + O2(g)  2 H2O(g)
 At low pressure
2H2(g) + O2(g)  2 H2O(g)