Download Chemical Bonding (short)

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Kinetics and
Thermodynamics
Section 1
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A) KINETICS – is the study of reaction
Mechanisms and rates of chemical reactions
Recall
that
Molecules
are
constantly
colliding
with
each
Other?
If they collide
with enough
force then
bonds can start
to break.
This is
“collision
theory”
Collision Theory - particles must
collide in order to react
Blue2 + Green2  Blue-Green
Or is it Blue-Green  Blue2 + Green2?
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Not all collisions are effective
In terms of “red”
and “blue”, what
makes the lowest
collision effective?
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ineffective
ineffective
ineffective
Only when red strikes blue can the
new bond form and old bond
ineffective
separate.
effective
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Effective Collisions mean that particles
must collide at proper angle and with
enough energy to break bonds
Molecules must collide in the right “orientation”
to break bonds OR in order to form new ones.
Effective Collisions
particles must also collide with
enough energy to break bonds
Breaking requires an
input of potential energy
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Effective Collisions
particles must also collide with
enough energy to break bonds
Where does the energy come
from to break chemical bonds?
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Effective Collisions
particles must also collide with
enough energy to break bonds
One molecule collides with another
causing them to break apart.
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Chemical Kinetics
Effective Collisions particles must also collide with enough
energy to break bonds
We can’t strike molecules directly of course.
But we can have them hit each other!
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Activation Energy:
Minimum energy needed to
break bonds and start reaction
Do you recall that the
molecules in a sample move
at different speeds?
 Increasing kinetic energy 
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This graph shows the distribution of
energies of the molecules in a sample
Notice that only a few molecules
are moving fast enough to break
the bonds during a collision
Majority
of
Molecules are
moving too slowly
Only a few
Molecules are moving
fast enough
 Increasing kinetic energy 
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This graph shows the distribution of
energies of the molecules in a sample
Majority
of
molecules
Few
molecules
Activation
Energy
Is this
minimum
speed
energy
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This graph shows the distribution of
energies of the molecules in a sample
Majority
of
molecules
Few
molecules
 Collisions
can transfer
more energy
during a
collision but
there is a
minimum that
will break the
bond
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How is
activation
energy like
getting
over a
wall?
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You can think of activation energy also as a hill
Energy is used
to break the bond
Energy is
released
when new
bonds form
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We can graph and see these
changes during the reaction.
Activation energy is added to “get over the wall
(to break the original bonds)
Energy stored
in the molecules
is on the Y axis
Reactants start
at a state of low energy
Energy is
released only
when the new
bonds form
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This is called a
potential energy diagram.
Activation energy is added to “get over the wall
Energy stored
in the molecules
is on the Y axis
Reactants start
at a state of low energy
Energy is released
only when the new
bonds form
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We will work more with these graphs later.
Activation energy is added to “get over the wall
Energy stored
in the molecules
is on the Y axis
Reactants start
at a state of low energy
Energy is released
only when the new
bonds form
Reaction Mechanism – most
reactions take place in steps
ex.
A2 + B2 --> 2 AB
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Reaction Mechanism – most
reactions take place in steps
step 1: A2 + B2  A2B2
step 2:
A2B2  2 AB
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A2B2 (activated complex)
A2 + B2
2AB
A2B2 is the higher energy intermediate
called the “Activated Complex”
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The activated complex is
formed and then used up so its
not in the overall equation
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+
A2 + B2  A2B2
A2B2  2 AB
A2 + B2  2 AB
Notice, the Steps
combine to form
overall reaction
Here’s another example:
Ex: H2 + 2 ICl  I2 + 2HCl
Net reaction
Most reactions proceed through multiple steps
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Section 1B
What determines the
speed of a reaction?
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Reaction rate factors
Reaction Rate:
depends on the
number of effective
collisions occurring
and number steps in
the reaction
mechanism
Sodium metal
in H2O
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Factors:
1) Nature of Reactants
more reactive chemicals
= faster reactions
ex: Na reacts faster than
Mg, since its group 1 its
more active than group 2
Na in H2O
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Factors:
2) Number of bonds broken and formed
Ex
2 H2 + O2  2H2O
OR
NaCl(aq) + AgNO3(aq)  AgCl(s) + NaNO3(aq)
which reaction looks faster?
Hint: how many steps in each?
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Factors:
2) Number of bonds broken and formed
Ex
2 H2 + O2 --> 2H2O
Synthesis
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Factors:
2) Number of bonds broken and formed
Ex
2 H2 + O2 --> 2H2O
Bonds
inside H-H
and O=O
must be
broken first
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Factors:
2) Number of bonds broken and formed
vs. NaCl(aq) + AgNO3(aq) --> AgCl(s) + NaNO3(aq)
Na+ Cl-
Ag+ NO3-
The Bonds are broken when reactants are first dissolved
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Factors:
2) Number of bonds broken and formed
vs. NaCl(aq) + AgNO3(aq) --> AgCl(s) + NaNO3(aq)
Na+ Cl-
Ag+ NO3-
Ag+ (aq) + Cl- (aq)  AgCl (s)
All that’s left to do is form the new bonds.
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Factors:
2) Number of bonds broken and formed
vs. NaCl(aq) + AgNO3(aq) --> AgCl(s) + NaNO3(aq)
One step; as silver and
chloride bond to form
the precipitate.
Ag+ (aq) + Cl- (aq)  AgCl (s)
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Factors:
2) Number of bonds broken and formed
(a) 2 H2 + O2 --> 2H2O
OR
(b)
NaCl(aq) + AgNO3(aq) --> AgCl(s) + NaNO3(aq)
Reaction (b) occurs faster since no
bonds need to be broken
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Reaction Rate Factors:
2) Number of bonds broken and formed
ex.
(a) 2 H2 + O2 --> 2H2O
OR
(b)
NaCl(aq) + AgNO3(aq) --> AgCl(s) + NaNO3(aq)
Double replacement are the
fastest reactions
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Hint: What causes the bonds to break?
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←One step, no bonds to break.
Breaking bonds in KClO3
Breaking bonds of H-H and O=O
Breaking bonds of N≡N and O=O
Which has the fewest bonds to break?
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What can be done to speed up a reaction?
Anything that
increases the number
of effective collisions
(bonds break)
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Temperature
Increase the temperature
Particles move faster so:
More collisions occur
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Temperature
Increase the temperature
Particles move faster and:
Collisions become more effective
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Temperature
Increase the temperature
Particles move faster so:
More effective collisions
More
Effective
Collisions
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Temperature
Increase the temperature
Particles move faster so:
Reaction
Rate
Increases
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What can be done to speed up a reaction?
Concentration
Increase concentration
More particles in a given volume
Particles are closer so:
More collisions
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In which will collisions be more frequent?
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Which reaction will be faster?
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What can be done to speed up a reaction?
Concentration
Increase concentration
More particles in a given volume
Particles are closer so:
More Effective collsions
Reaction rate increases.
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What can be done to speed up a reaction?
Pressure (In gases)
Increase pressure,
Decreases volume so:
Increases concentration:
More collisions
Molecules are closer together
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Increasing pressure forces
molecules closer together
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What can be done to speed up a reaction?
Pressure (In gases)
Increase pressure,
Decreases volume so:
Increases concentration:
More collisions
Molecules are closer together
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What can be done to speed up a reaction?
Pressure (In gases)
Increase pressure,
Decreases volume so:
Increases concentration:
More collisions
Reaction rate increases
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Catalyst
A catalyst is a substance that helps reactions
without being changed itself
It can provide a
surface
So the Molecules don’t
need as much energy
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Catalyst
A catalyst is a substance that helps reactions
without being changed itself
i.e. A catalyst
lowers the
activation energy
required.
(the molecules can move
slower and still react)
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Ex: Catalytic converter
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Ex: Catalytic converter
In your car’s catalytic converter, pollutant
gases are broken down into safer products
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Surface area
In reactions involving
different phases
In Heterogeneous
mixtures
Usually solids
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Surface area
In reactions involving
different phases
Smaller particles
= More surface
= More collisions
= Faster reactions
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Questions:
1. Explain what is meant by “Collision theory”
2. Do all collisions result in a chemical change?
Explain.
3. In what way is activation energy like a wall to
get over?
4. Which reaction is faster? (explain why)
(a) H2SO4 (aq) + Ca(OH)2 (aq)  H2O (l) + CaSO4 (aq)
Or (b) H2 (g) + I2 (s)  2HI (l)
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5. How does each factor affect the rate of a chemical
reaction? Explain each in terms of effective collisions
TemperatureConcentrationPressure –
Particle size (surface area)Catalyst-
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(a) H2SO4 (aq) + Ca(OH)2 (aq)  H2O (l) + CaSO4 (aq)
Or (b) H2 (g) + I2 (s)  2HI (l)
6. Which reaction in #4 would benefit from an increase in
pressure? …An increase in surface area?
7. How could you speed up the reaction of aqueous solutions, as
in reaction 1 in #4?
8. Which reacts faster a 1 cm3 piece of Na or 1 cm3 of Mg? Why?
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Section 2
Thermodynamics
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2A Enthalpy:
Energy changes in chemical reactions
Heat released during a reaction
will come from stored energy
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Enthalpy – energy stored in chemical
substance (mostly PE)
its symbol H (for Heat content)
Reactant’s
energy
Product’s energy
 As the reaction proceeds 
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Enthalpy: Energy changes in chemical reactions
Bond breaking
energy is
absorbed
Bond forming
energy is released
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Enthalpy: Energy changes in chemical reactions
Bond breaking
energy is
absorbed
Bond forming
energy is released
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Energy
IN to break bonds
Energy
OUT as new bonds form
ΔH
Net (overall)
loss of energy
ΔH Heat of Reaction Ref Tbl I
– Net energy absorbed or released in a reaction
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Energy in to break bonds
Energy out as new bonds form
ΔH
Net (overall)
loss of energy
If more energy is released overall, that’s exothermic
If more energy is absorbed overall, that’s endothermic
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Potential Energy Diagrams
Show the change in enthalpy ∆H that occurs during a reaction
Activated complex
Ex: A + B  C + D
Reactants A & B at 40 kJ
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60 kJ added to break bonds
and form the
activated complex at 100 kJ
(this is the activation energy)
80 kJ is released as bonds form products C & D at 20
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Potential Energy Diagrams
Show the change in enthalpy ∆H that occurs during a reaction
Activated complex
ΔH - The difference
between the
reactant's H(PE)
and product's H(PE):
ΔH = H products -- H reactants
ΔH = 20 kJ – 40 kJ = – 20 kJ
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Potential Energy Diagrams
Show the change in enthalpy ∆H that occurs during a reaction
Activated complex
Notice that the
change is negative
since heat
is being lost
Exo-thermic!
ΔH = H products -- H reactants
ΔH = 20 kJ – 40 kJ = – 20 kJ
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Potential Energy Diagrams
Show the change in enthalpy ∆H that occurs during a reaction
Activated complex
Notice also its
ending energy
minus starting
energy
ΔH = H products -- H reactants
ΔH = 20 kJ – 40 kJ = – 20 kJ
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Energy
decreases
Exothermic - products have less
energy than reactants: Heat lost
ΔH is negative ( - ΔH)
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Energy
decreases
in the thermochemical equation:
ex. 2H2 + O2 
ΔH =
2H2O + 573.7 kJ
- 573.7 kJ
Energy as a product: written on right side
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Energy
increases
Endothermic
= products with more energy than reactants:
Heat gained ΔH is positive ( + ΔH)
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Energy
increases
Thermochemical equation:
ex. 2H2O + 573.7 kJ

2H2 + O2
ΔH = + 573.7 kJ
Energy as a reactant on left side of equation
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Activation energy
Energy required to initiate a reaction (for collisions to be effective)
Breaks bonds of reactant molecules to form activated complex
Activation
energy
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A catalyst is a substance that speeds up a chemical reaction,
MnO
but is not consumed by the reaction. Ex: H2O2
H2O + O2
Manganese dioxide MnO2 is the catalyst.
→
2
Adding a catalyst lowers
the activation energy
It provides a different
reaction (mechanism) with a
lower activation energy
But NO change in
H reactants or H products
So ΔH doesn’t change
Notice: ΔH not affected
ΔH
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TYU
1. What is the activation energy for the forward
reaction? The reverse?
2. What is the PE of the reactants? Products?
3. What is the PE of the activated complex?
4. What is the heat of reaction?
5. Which values will change when a catalyst is
added?
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2B) Heat of reaction: Ref tbl I
For 2 moles
Heat released or gained
for the molar quantities shown
in the reaction (coefficients)
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2B) Heat of reaction: Ref tbl I
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Ex: How much heat is released by the complete combustion of
1 mole of octane, C8 H18 ? according to the equation:
2C8 H18 (s) + 25O2 (g)  16CO2 (g) + 18H2O?
10943 joules
2 moles
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÷2
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Exothermic = -H
+ 66.4 becomes
- 66.4 for the
reaction in reverse!
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Exothermic?
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Endothermic
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How much heat is
released for 1.0 mole?
÷2
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Questions:
1. Is bond breaking exothermic or endothermic? Explain.
2. Describe the sign convention that is used in thermochemical
calculations.
3. Sodium acetate dissolves readily in water.
Its H = - 17.3 kJ /mol. Would this process increase or decrease the
temp of the water? Explain
4. How is the activation energy of a reaction like a wall or barrier?
5. If bond breaking is always endothermic, how can a reaction be
exothermic?
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6. For the reaction A + B <==> C, the activation energy of the forward reaction is
5kJ and the total energy change is – 20kJ. What is the activation energy of the
reverse reaction? Draw a PE diagram to illustrate the problem:
7. Sketch an energy profile curve (PE diagram) for this gas phase reaction:
F2(g) + H2(g) <==> 2HF(g) + 103 kJ
H = - 103 kJ and activation energy = 22 kJ
8. Sketch a generic PE diagram for an endothermic reaction. Label the reactants
and products, the heat of reaction (∆H), and activation energy (Eact).
9. On your sketch above, draw a dotted line on your PE diagram and label it to
show the effect of adding a catalyst. Label it Ecat
10. What happens to the sign of H when the reverse of a chemical reaction is
written? Why?
2C) Entropy (Randomness) ΔS
Entropy is Randomness
Entropy tends to increase
“spontaneously”
In simple terms:
Molecules tend to “spread out”
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2C) Entropy (Randomness) ΔS
Entropy is Randomness
Entropy tends to increase
“spontaneously”
Molecules tend to “spread out”
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Phases:
Highest
Gas
far apart
And liquids tend to vaporize
Since
entropy is
increasing
Liquid
close but mobile
Solids tend to melt
Lowest
Solid
close not mobile
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Spontaneous reactions:
Continue on their own once started
Depends on Heat (ΔH) AND Entropy (ΔS)
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Spontaneous reactions:
Heat: exothermic (- ΔH)
heat lost: “Favorable”: tends to keep going
Exothermic can recycle energy
to keep themselves going
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Spontaneous reactions:
Heat: exothermic (- ΔH)
heat lost: “Favorable”: tends to keep going
H is the leftover energy
Spontaneous reactions:
endothermic (+ΔH) heat gained:
“unfavorable” tends not to keep going
Endothermic don’t give off enough
energy to provide activation energy
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Spontaneous reactions:
endothermic (+ΔH) heat gained:
“unfavorable” tends not to keep going
H is the extra energy added to make up the difference
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Spontaneous reactions: Continue on their own once started
Depends on Heat (ΔH) and Entropy (ΔS)
Heat: exothermic (- ΔH) heat lost: “Favorable”: tends to keep going
endothermic (+ΔH) heat gained: “unfavorable” tends not to keep going
Entropy:
increase – more random (+ ΔS) favorable
decrease – less random (- ΔS) unfavorable
Temperature
increases effect of entropy change:
Hotter = more entropy
H
S
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Spontaneous reactions: Continue on their own once started
Depends on Heat (ΔH) and Entropy (ΔS)
Heat: exothermic (- ΔH) heat lost: “Favorable”: tends to keep going
endothermic (+ΔH) heat gained: “unfavorable” tends not to keep going
Is a reaction spontaneous or not spontaneous?
Depends on which is greater contribution
… more favorable
or more unfavorable
H
S
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Examples
Ex 1: burning paper
CH(s) + O2(g)  CO2(g) + H2O(g) + energy
heat released (- ΔH) : exothermic: favorable
solid  gases: (+ ΔS): inc entropy: favorable
Both changes favorable: its spontaneous
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Examples
Ex 2: Boiling H2O (l) + 2260 kJ  H2O (g)
liquid  gases: (+ ΔS): inc. entropy: favorable
heat absorbed (- ΔH) : endothermic: unfavorable
At high temperature (above 100 0C) entropy is greater
(more favorable, water boils)
At lower temps (below 100 0C) heat (ΔH) is greater
(more unfavorable – doesn’t happen)
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Questions
1. Define these in your own words
Entropy –
Spontaneous reaction -
2. What two factors determine if a reaction is spontaneous?
3. What change in enthalpy (ΔH) is favorable? Why?
4. Suppose the products in a spontaneous process are more ordered than the
reactants. Is the entropy change favorable or unfavorable? Explain.
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5. At normal atmospheric pressure, steam condenses to liquid water,
even though the change is an unfavorable entropy change. Explain
why.
6. Determine if each is an increase or decrease in entropy.
(Also, explain why)
CaCO3(s)  CaO(s) + CO2 (g)
2 H2 (g) + O2 (g)  2 H2O (l)
NaCl (s)  NaCl (aq)
2 H2 (g) + O2 (g)  2 H2O (g)