2. PHASE DIAGRAMS (N. AND I.)
■New and Improved!
■Phase Change Diagrams only account for
temperature
■New Phase Diagrams account for
Pressure, as well.
■Terms:
■Critical Point
■Triple Point
*This change comes as a result of our study of Gas Laws!
3. PHASE DIAGRAMS
Temperature →
Pressure
→
A
B
C
D
Solid
Liquid
Vapor
The large white dot
(T) represents the
triple point.
The triple point is the
equilibrium* point of
all three phases–
solid, liquid, gas.
Point C – The Critical Point– at temperatures higher than
this, there can only the gas phase
T
4. MORE TRIPLE POINT
■Along any line,
the two phases
represented are at
equilibrium.
Temperature →
Pressure
→
A
B
C
D
Solid
Liquid
Vapor
5. CHEMICAL LINE UP
Temperature →
Pressure
→
A
B
C
D
Solid
Liquid
Vapor
Each Line has a “meaning”:
Tp
-A: Solid-Liquid Eq.
Tp
-B: Solid-Gas Eq.
Tp
-C: Liquid-Solid Eq.
Tp
-D: Supercooled H2O
6. EQUILIBRIUM
■The state in which two or more processes
are occurring at the same time:
■Phase Changes
■As one molecule melts, another freezes
■As one molecule boils, another condenses
■Triple Point:
■ALL of the phase changes occur at the SAME time, at
Equilibrium
7. RELATIONSHIPS (FOR WATER)
■1 atm:
■BP = 100 C
■FP = 0.016 C
(273.16K)
■0.0063 atm:
■BP = 0.0098 C
■FP = 0.0098 C
Temperature →
Pressure
→
A
B
C
D
Solid
Liquid
V
a
p
o
r
1 atm
100
C
0.0063 atm
0.0098
C
0
C
High-Density
Ice
1000 atm
8. SOME IDEAS TO CONSIDER
■Adding pressure causes a shift to the more
dense phase.
■Adding temp. causes a shift to the less
dense phase
■“Gibbs Phase Rule”
9. WATER’S UNIQUE PROPERTY
■What is the most dense phase of water?
Gibbs’ Phase Rule explains why ice skates work the way they do!!!
The blade of a skate provides pressure to the ice, which should be
near 0 C
This causes a shift to the more dense phase.
This provides a film of water which allows the skate to slide!
10. OTHER PHASE CHANGE GRAPHS
■Carbon
Dioxide
■What can this
tell us about
CO2
?
Temperature →
Pressure
→
Solid
Liquid
Vapor
1 atm
0 C
11. THERMODYNAMICS
■The study of heat/energy as it moves
through a system
■Phase changes: Melting = Freezing?
■Yes-- in one, energy is being added (melting); in the other,
it’s being taken away (freezing)
■Calculated using Calorimetry
■Uses masses and temperature changes to
determine energy flow
12. CALORIMETRY
■Energy can be calculated using the
formula:
Q = m x c x ∆T
Where:
Q = energy (joules, kilojoules, or calories)
m = mass in g
c = Specific Heat (more)
∆T = Change in Temperature (K or o
C)
13. SPECIFIC HEAT
■Term given for the amount of energy
needed to raise 1 g of a substance a given
temperature:
■For water, c = 1 cal / g o
C
■Also c = 4.18 j / g o
C
■From which we can deduce:
■1 cal = 4.18 j
14. FOR ICE? STEAM?
■Ice’s Specific Heat:
■2.09 j/g o
C
■Steam’s Specific Heat:
■1.84 j/g o
C
15. ENERGY OF A PHASE CHANGE
■Heat of Fusion: The amount of heat
required to MELT a solid;
■Also, the heat given off by freezing a liquid
■For H2
O: 333 j / g
■Heat of Vaporization: The amount of heat
needed to BOIL a liquid;
■Also, the heat given off by condensing a gas
■For H2
O: 2230 j / g
17. WE KNOW:
■Chemical Bonds form due to electron
movement to a state of lower energy.
■NOW: we begin learning about the forces
that drive reactions.
18. CHEMICAL KINETICS HAS FOUR
PRIMARY PARTS:
■Kinetic Molecular Theory
■Reaction Mechanisms
■Reaction Rates
■Phases and Phase Changes (already started!)
19. KINETIC MOLECULAR THEORY
■Says that all particles are moving at all
times
■Also called collision theory
■“Proven” (sorta) through our studies of
gases.
20. POSTULATES OF KMT
■All molecules are extremely small in mass.
■All molecules are in constant motion.
■There are so many particles, even in a
small sample, that we can use statistics
(mean, median, mode, etc.) to make
generalizations about particle movement.
21. MORE POSTULATES OF KMT
■Collisions are perfectly elastic (all energy
remains kinetic)
■The kinetic energy of the system is directly
related to the temperature (higher temp. =
more E!)
■… So what?
22. WHAT IT MEANS:
■Heat is a function of the energy of
motion of particles
■When something gets hot, its particles are
moving faster than before.
■Everything, whether you like it or not, is
moving (at a molecular level).
■Chemical reactions and interactions
are all a result of chemicals colliding with
one another… which leads us to…
23. REACTION MECHANISMS
■Chemical Kinetic Theory says that NO
reaction can occur without the effective
collision of two or more molecules.
■Contact
■Electron Clouds Must “touch”
■Proper Alignment
■Molecules Must Align correctly in order to react
■High Enough Energy to Begin Reacting
■Activation E!
■This WHOLE concept is called:
COLLISION THEORY
24. REACTION MECHANISMS
■Are defined as the underlying action in a
chemical reaction
■2C8
H18
+ 25 O2
16CO2
+ 18 H2
O
■What are the odds that both Octanes and ALL 25
O2
collide at the same time?
■Problem: this is a highly exothermic reaction,
rapid reaction-- How can this be?
25. ANSWER: ACTIVATED COMPLEXES
■Molecules that exist ONLY DURING a
reaction
■Parts of molecules that have either:
■Broken apart from reactants
■Formed together to make products
■CO and OH (0) exist during this reaction,
among others.
■Act. Complexes Video
26. SO WHAT?
■If we agree that kinetic theory is true,
and…
■Activated Complexes exist, then…
■We can control how fast some chemical
reactions occur.
27. RATE DETERMINING STEP
■Activated Complexes say that reactions
occur in steps.
■Some steps are slower than others, which
limit how fast a reaction can occur.
■Video here
28. REACTION RATES
■Demonstration Reaction:
■Zn + NH4
NO3
ZnO + N2
+ H2
O
■Slow Reaction was Changed to a quick one!
■Reaction Rate can be Affected!
■Slow Rate: Heat given off
■Fast Rate: EXPLOSION!!
■Watch this guy…
29. FACTORS AFFECTING REACTION
RATE
■There are FOUR you need to know:
■Nature of the Reactants
■Temperature of the System
■Concentration of Reactants and/or Products
■Catalysts
30. NATURE OF REACTANTS
■Depends on what types of things are
being reacted.
■Primarily a measure of the reactivity of
the involved chemicals.
■The more reactive each participant, the more
quickly the reaction occurs.
■Reactivity increases as E! increases
■Affects the number of effective collisions
■Includes Phase Differences: More later
31. TEMPERATURE
■In a vast majority of Reactions, a
higher temp. Results in a faster
reaction
■An increase in effective collisions from
having higher E! available
■To effectively deliver temperature:
■Add Heat: Fire, Sun, etc.
■Add H2
O: Has a high Specific Heat
32. CONCENTRATION AND REACTION
RATE
■An increase in the Conc. of Reactants=
Faster Reaction
■An increase in the Conc. of Products=
Slower Reaction
■Think of a Jr. High dance…
33. CONCENTRATION AND REACTION
RATE (CONTINUED)
■Surface Area Increase = An Increase in
Conc.
■More Molecules Available to Collide
■Phases have variable surface area*:
■Liquid-Liquid
■Liquid-Gas (bubbling)
■Gas-Gas
■Liquid-Solid
■Gas-Solid
■Solid-Solid
34. CATALYSTS AND REACTION RATE
■Catalysts always increase reaction
rate.
■Defined:
■Chemicals that act to speed up a reaction
without being used up themselves in the
reaction.
■We do not fully understand the function
of Catalysts
■Catalysts MAY break down during a
reaction-- but will always re-form by the
end of it.
35. REVERSIBLE REACTIONS
■Chemical Reactions that can occur in
either direction in nature.
■Under certain circumstances, reactants
become products
■Under different circumstances, products
turn back into reactants
■Can (and will) achieve Equilibrium
A + B C + D
Vid Here
36. LECHATELIER’S PRINCIPLE
■LeChatelier said that a reversible reaction
can be affected by adding a stress.
■Stresses include:
■Concentration of Chemicals
■Temperature
■Addition of a Catalyst
37. LECHATELIER’S RESULTS
■Concentration of Chemicals
■More reactant added
■Shifts reaction to the products side
■More product added
■Shifts reaction to the reactants side
■Temperature
■Increasing temp causes:
■Exothermic reactions to move toward reactants
■Endothermic reactions to move towards products
38. LAST FOR LECHATELIER
■A Catalyst will ALWAYS move a reaction
towards the products side
■Catalysts always act to speed up reactions
39. CONTROLLING REACTIONS IN THE FIRST
PLACE?
■We can control how fast reactions take place.
■We know that collisions are the root of all
chemical reactions.
■We know that particles are in constant motion.
■Can we make a reaction occur? Stop a reaction
from occurring?
40. CONTROLLING CHEMICAL REACTIONS
■It is possible to prevent reactions from
occurring, and it is possible to make
reactions occur that normally wouldn’t:
■Activation Energy
■Enthalpy and Entropy
41. ACTIVATION ENERGY
■The Energy Required to START a
Chemical Reaction
■E! act (abbr.)
■Examples: Gasoline; Ba(OH)2
+ NH4
Cl
■Can be shown on an potential E! Graph
■Without enough E!, no reaction will
occur.
42. REACTION RATE FORCES
■Additionally, nature controls whether or
not reactions through two forces:
■Enthalpy, or ΔH
■The total amount of energy in an object/chemical
■Entropy, or ΔS
■The amount of disorder in an object or chemical
43. WHAT NATURE WANTS…
■Nature wants:
■Energy to be given off to the universe
■An decrease in ∆H (- ∆H)
■Entropy to increase
■An increase in ∆S (+ ∆S)
44. PREDICTING REACTION OCCURRENCE
■Given the ΔH and ΔS of a system, you can
predict if a reaction will occur.
■If ΔH is negative and:
ΔS is Positive, the reaction will occur
ΔS is Negative, the reaction will occur if:
ΔH has a greater absolute value than ΔS
■If ΔS is positive and:
ΔH is positive, the reaction will occur if:
ΔS has the larger absolute value
45. MORE PREDICITING
■If ΔH is positive and ΔS is negative the
reaction will never occur.
■Although not “the whole story,” phases are
related to ΔH and ΔS.
46. PHASES AND FORCES
■Recall:
Solids: Very Low ΔS, Very Low ΔH
Orderly, low E!
Liquids: Low ΔS, Low H
Semi-Orderly, moderate E!
Gases: High ΔS, High ΔH
Disorderly, high E!
Plasma: Very High ΔS, Very High ΔH
Ridiculous disorder, ridiculous E!
47. INTERMOLECULAR FORCES
■Forces that act between 2 separate
molecules.
■Phases are directly related to the
strength of the Intermolecular Forces in
a system.
■Solids- Strong IMF
■Plasma- Weak IMF
49. REMEMBER THIS?
■To aide in your understanding of IMF,
check out this blast from the past:
0 2
0.9
3
1.8
1
50. LONDON DISPERSION FORCES
■Occur primarily in Non-Polar Covalent Cpds., but
do occur in all substances.
■Caused by the random “dispersion” of electrons
within a compound.
■As e- move, they create weak positive and negative areas
that are therefore attracted to each other.
■A temporary electro-magnetic force
51. DIPOLE FORCES
■Occur in polar molecules and ionic
compounds only.
■The (∂+) or (+) portion is attracted, quite
strongly, to the (∂-) or (-) part.
■*This force is stronger as substances become
more ionic in character.*
■A permanent, though variable (depending on
the differential), electro-magnetic force
52. HYDROGEN BONDS
■Occur only in Polar Molecules that
contain a hydrogen and highly
electronegative element.
■F, O, Cl <in decreasing order>
■A permanent, average strength (because of
the partial differential) electro-magnetic force
H O
H
H
H
O
Hydrogen Bond
53. GRAVITY
■The pull of two massive bodies towards
each other
■For their size, atoms are quite massive
■Extremely weak force in spite of this
■Gravity is directly tied to mass; size, at the atomic level, is
generally irrelevant
54. IMF AFFECT PHYSICAL PROPERTIES
■According to periodic trends, Water
should boil at a much lower temperature:
Atomic Mass →
Temp
●H2S
●H2Se
●H2Te
●(H2Po)
H2O ●H2S
●H2Se
●H2Te
●(H2Po)
● H2O
● H2O boils at 100 C; H2Te boils at -2 C!
Trends Say:
Real Life Proves:
