Notes 1.1
Exothermic reactions give off heat
120
100
80
Heat
Content
60
40
20
0
0
2
4
6
Since reactions want to minimize energy you would think that the
reaction would be spontaneous like a ball rolling down a hill
C6H12O6 + 6O2  6H2O + 6CO2 + 2806kJ/mol is exothermic
We don’t worry about sugar cubes exploding on contact with
air…why not?
• Demo sugar cub and air, and flame then powdered sugar and
flame. Cornstarch and ABS.
Something prevents or slows the reaction from occurring
160
140
120
100
80
60
40
20
0
Heat
Content
0
2
4
6
Reaction kinetics
• The study of the rates of reactions and the factors that affect
the rates
•
o amount can be products or reactants,
• Ex if 16g of HCl are used up in 12 minutes of a reaction the
average rate is
•
• Try
o If a reaction between CaCO3 and HCl produces 245ml
of CO2 in 17s then what is the reaction rate?
o
Assign Exercises 1-5 page 2
Read 1.2 together and answer 6 as a class…
The slope of the graph is the rate (rise over run = amount
over time)
How can we determine reaction rate?
• There are four ways to measure the rate of the following
reaction
•
Cu(s) + 4HNO3(aq)  Cu(NO3)2(aq) + 2H2O(l) + 2NO2(g) + heat
• Red-brown (none)
blue
brown
1. Colour change
a. Use a spectrophotometer to measure the intensity of the
blue colour (amount of Cu(NO3)2(aq))
i.
2. Temperature change (heat is a product)
a.
3. Pressure change (a gas is produced and could be held in a
sealed container)
a.
4. Mass change (run the reaction for different periods of time
and weigh the copper after each trial)
a.
Assign 7-9 and read Activity 18A
1.3 Factors Affecting Reaction Rates
Your experiments should have shown that HCl and Mg :
• Reacted faster at higher temperatures
• Reacted faster in more concentrated solution
higher rate or less
 time
• Reacted faster with more surface area
reactions-and-rates.jnlp
Other experiments have found that:
• When gaseous reactant pressure increases the rate increases
o (Pressure is just concentration of a gas)
• The nature of some reactants (chemical properties) makes
them react very slowly or quickly regardless of the
conditions.
o We can’t easily control these rates (can’t slow down the
explosion of nitroglycerine much by cooling it)
o More or stronger bonds being broken or formed tends
to slow down reactions
o Look at molecules on page 7 and notice many bonds
formed and broken in the first slow reaction and only
an electron transferred in the second fast reaction
o Assign 10-11
• The ability of reactants to meet (remind of sugar demo)
o Increasing the surface area increases the sites that the
reaction can occur on
 Basically raises the concentration of reaction sites
o Homogeneous reaction
 The reactants are in the same phase and perfectly
mixed
• The reaction can happen everywhere at once
• 2 gases
• 2 substances dissolved in water
• two liquids completely dissolved in each
other
o Heterogeneous reaction
 Reactants are in different phases and not
completely mixed
• Reaction can only happen on the surfaces
where the reactants meet
• Solid and liquid
• Liquid and gas
• Solid and gas
• Two immiscible liquids
• Phase
o Reactants can’t move easily in solids
o Reactants are closest in ionic solution
o Aqueous ions > (gases or liquids)> solids
Assign 12-14
Catalysts
• Chemical that can be added to a reaction to increase the rate
of reaction
o After the reaction is complete there is the same amount
of the catalyst present
• Demo 2H2O2  2H2O + O2 with MnO4
• Most enzymes are catalysts
Inhibitors
• Reduces reaction rate by combining with a catalyst or
reactant
• Demo same reaction with dishsoap
• Examples are antibiotics and poisons
Assign 15-17
Notes 1.4
Experimental reaction rates
• A plot of [reactant] vs time doesn’t give a strait line
• The reaction goes fast at the beginning because there is lots
of reactant
• The reaction slows at the end as the reactant is used up
• Average rate looks at start and finish
• Exact rate at a given time is from the slope of the tangent at
that time
Try18-19
1.5 Collision theory (aka Kinetic Molecular Theory)
• Molecules act like small hard spheres that bounce off one
another and transfer energy during collision
• Molecules must collide before they react
o Effective collisions produce the products
 Correct orientation and energies
• http://www.chem.iastate.edu/group/Greenbowe/sections/proj
ectfolder/animations/NO+O3singlerxn.html
• reactions-and-rates.jnlp
• This explains why two conditions changes can change the
rate
o Increasing the concentration of reactants means
collisions are more frequentrate increases as
collisions/sec increases
o Increasing temperature makes the molecules move
fastercollide more often and with more energyrate
increases as collisions/sec increases.
 Rule of thumb is as temp increases 10o the
reaction rate doubles
Assign 20-22 and Worksheet up to 10
Quiz tomorrow on factors controlling reaction rate
01-06 Enthalpy
demo endothermic and exothermic with ammonium nitrate and
water and sulphuric acid and sugar
Notes 1-06
Enthalpy changes
• As two atoms approach each other the nuclei are
attracted to the electrons but the electrons repel each
other
H+H
PE
H2
Reaction
• They settle at a distance that minimizes the repulsive
forces and maximizes the attractive forces
• The potential energy of the system is minimized and we
say that a bond has formed.
 The atoms ‘want’ to stay together and it will take
energy to separate them
Potential energy
• Stored energy that results from position in space
Kinetic energy
• Energy from the movement of molecules or atoms
within a molecule
Bond energy
• The energy required to break the bond
 The same amount of energy is released when that
bond forms
 Cl2 + 243kJ --> 2Cl
 2Cl --> Cl2 + 243kJ
assign 23
Most reactions are much more complicated
• many bonds break and form
• heat can move in or out
• work can be done by or on the system
Enthalpy
• the sum of all the kinetic and potential energies that
exist in a system when at constant pressure
• ∆H = Hprod-Hreact is the change in enthalpy during the
course of a reaction
• Draw diagram on page 14
• It is ∆H not the values of the reactants or products that
is important
Endothermic reactions
• Hprod > Hreact and ∆H >0
• 2N2 + O2 + 164 KJ --> N2O
• or 2N2 + O2 --> N2O; ∆H =+164 KJ
2N2O
Enthalpy
2N2 + O2
∆H=+164kJ
• ∆H points up to the products and is positive
• Energy enters the system and it is endothermic
• demo ammonium nitrate and water
Exothermic reactions
• Hprod < Hreact and ∆H <0
• H2 + Cl2 --> 2HCl + 184 kJ
• H2 + Cl2 --> 2HCl; ∆H = -184 kJ
H2+Cl2
Enthalpy
(H)
∆H= -184kJ
2HCl
• ∆H points down to the products and is negative
• Energy exits the system and is exothermic
• demo sulphuric acid and sugar
Assign 24-28
Notes 1.7 Kinetic energy distributions
For a reaction to occur you need collisions and sufficiently
energetic molecules.
• At room temp the average gas molecule has 1010 collisions a
second
o Frequency isn’t the problem, energy is.
o Even at low temperatures some molecules react
 Kinetic energies vary from molecule to molecule
# of
molecules
Average energy
Kinetic energy 
At different temperatures the average kinetic changes:
Number of
molecules
with a
particular
energy
25o
200o
400o
Minimum energy
to successfully
react
10o increase equals an approximate doubling of rate for slow rxns
• Both collision frequency and number with the minimum
energy increase
Assign 29-32
Notes 1.8 Activation energies
The existence of a minimum energy means there is an energy barrier even for exothermic
reactions.
Molecules without enough energy just roll back down
PE
Activated complex
R
P
Think of rolling a ball up a hill. You increase PE as you climb but
your kinetic enrgy decreases as you slow down. After you crest
the hill you speed up as your potential energy is converted into
kinetic energy
Activated complex
• The arrangement of atoms which occurs when the reactants
are in the process of rearranging to form products
• An intermediate molecule
Activation energy Ea
• Minimum energy required to change reactants into products
If KE is less than Ea then the atoms’ electron repel before the
activated complex forms
• Ineffective collision
If KE = Ea then the activated complex forms but the molecules
stop and could fall down either side of the hill
• Reaction is possible but not guaranteed
If KE > Ea then the activated complex forms and the molecule
continues to move and can go to the products
Draw diagram from hebden p 21 to relate the size of the energy hill
to the number of molecules with Ea at a given temperature.
When the hill is high or the temp is low very few molecules
have sufficient energy to form the activated complex and vice
versa
Assign 34-37
Energy is conserved during a reaction so the total E (KE
and PE) stays constant through the reaction
• as one rises the other falls
• Draw page 22 Hebden
A successful collision requires both sufficient KE to make
contact and correct alignment of the molecules.
• If the alignment is off the energy hill is higher because
they still need KE to “wiggle into position”
• Draw and discuss the diagram at the bottom of 22.
Assign 38-40
Looking at the Potential Energy diagrams there is no
visible reason why molecules can’t go over the hill from
left to right and right to left
• Reactants can form products, but at the same time…
• Products can form reactants!
REACTANTS ↔ PRODUCTS
Ea(for)
PE
Ea(rev)
Reactants
∆H
Reaction proceeds
Products
Ea(for) is the activation energy for the forward reaction
Ea(rev) is the activation energy for the reverse reaction
The activation energy is always endothermic
• It takes energy to start the forward and reverse
reactions
Look at and draw the other 3 diagrams on pages 24 and 25
and convince yourself that
∆H = Ea(for) – Ea (rev)
Note
• If you are confused about the formula or you forget it
just sketch a quick PE diagram to determine the sign
of ∆H.
Assign 41-45.
Notes 1.9
Reaction mechanisms
• The sequence of steps in an overall reaction
• The odds of 4 or more molecules coming together and
colliding at exactly the same time with exactly the right
orientation are virtually nil
o Complex reactions must happen in simpler steps to be
possible
4HBr + O2  2H2O + 2 Br2
has 3 experimentally determined steps (you won’t be asked to
predict these on your own!)
Step 1: (slow)
The dots are bonds that are in the process of breaking or forming.
Step 2 (fast)
Step 3 (fast)
There are 3 elementary processes (steps) in the reaction
mechanism
Information you get from this mechanism:
1. the slowest step is the rate determining step (step one in this
mechanism)
2. if HOOBr and HOBr are used up quickly but form slowly
their concentrations can’t build up
3. Since HOOBr and HOBr aren’t used in the rate determining
step adding more of these will not affect the overall reaction
rate
4. to determine the overall reaction add all the steps together
and cross off whatever appears on both sides of the arrow.
a. HBr + O2
 HOOBr
b. HBr + HOOBr  2 HOBr
c. 2HBr + 2 HOBr  2 H2O + 2Br2
d. 4HBr + O2
 2H2O +2Br2
5. HOOBr and HOBr are intermediates that can exist on their
own but react quickly in this mechanism
6. The formula of the activated complex is found by adding the
atoms of the reacting molecules (order isn’t important)
Assign 46-53
Break
Notes 1.10
Using:
a. HBr + O2
 HOOBr (slow)
b. HBr + HOOBr  2 HOBr (fast)
c. 2HBr + 2 HOBr  2 H2O + 2Br2 (fast)
Each step has it’s own activated complex and therefore its own
activation energy.
• There are 3 humps in the overall PE diagram for
• 4HBr + O2
 2H2O +2Br2
HOOBr
HOBr
H2O
PE
HBr + O2
Br2
Step 1
Step 2
Step 3
The activation energy for any one step = PE (activated complex) –
PE (reactants for the step)
• draw on the graph
• each Eact starts from a new level!
Assign 54, 55
Notes 1.11
The effects of catalysts
Catalyst
• a substance that provides an overall reaction with an
alternative mechanism having a lower activation energy
• ∆H is not changed but the energy hill is
uncatalyzed
PE
catalyzed
reaction proceedes
• A catalyst increases both the forward and reverse reaction
rates because more molecules have enough kinetic energy to
form the activated complex
uncatalyzed
catalyzed
KE
PE
# of molecules
reaction proceeds
Keep in mind that both pathways are still available and some may
still use the uncatalyzed pathway but most take the easy route.
Draw Pt catalyst diagram from page 32
Notes 1.12
The effect of a catalyst on the reaction mechanism
The reaction
OCl- + I-  lO- + Cl- has huge activation energy because it
involves the approach and collision of two negatively charged
species.
[I--O--Cl]2-
IO- +Cl-
PE
I- + OCl-
Reaction proceeds
The reaction is therefore very unlikely to occur.
A catalyst (water) provides an alternative mechanism with a lower
overall activation energy
Notes
• The catalyst does take part in the reaction
• It is simply regenerated in a later step of the mechanism so at
the end there is the same amount as was started with
• ∆H for the overall reaction is the same in catalyzed and
uncatalyzed reactions
• both intermediate species and catalysts cancel out in the overall
reaction
• intermediate species are produced and then used up
• catalysts are first reactants that are later produced again
Assign 56-61 and assign the worksheet
Notes 1.13
Some uses of catalysts
Catalysts are vital to industry and one of the main reasons life is
possible
Enzymes
• biological catalysts that control virtually all the chemistry in an
organism
• act on a substrate
• made of protein
• the complex shapes bind and then orient reactants
•
•
• C12H22O11 + H2O  2 C6H12O6 is catalyzed by maltase
Producing sulfuric acid
• S + O2  SO2
• The sulfur dioxide is passed over a vanadium pentoxide or
platinum catalyst
• 2SO2 + O2  2 SO3
• The sulfur trioxide in passed through water to make sulfuric
acid
• SO3 + H2O  H2SO4
Catalytic converters
• Automobile engines are hot enough to combine N2 and O2 to
make NO2
• 2NO + O2  NO2 (smog)
• They also release unburned hydrocarbons
• http://auto.howstuffworks.com/catalytic-converter2.htm
• The catalytic converter contains
• platinum, palladium and rhodium which catalyze the
oxidation of the hydrocarbons to less harmful CO2 and H2O
• Use transition metals to catalyze the reduction of NO back to
N2 and O2
Assign 62-63
Test next double block