Unit 2 spring_b9944c7a5cb9b7f26e3a54c48b96157c.pdf

chem

1. © 2018 Pearson Education Ltd.
Chapter 10

Gases
James F. Kirby
Quinnipiac University
Hamden, CT
Lecture Presentation

2. © 2018 Pearson Education Ltd.
Characteristics of Gases
•Physical properties of gases are all similar.
•Composed mainly of nonmetallic elements with
simple formulas and low molar masses.
•Unlike liquids and solids, gases
–expand to fill their containers.
–are highly compressible.
–have extremely low densities.
•Two or more gases form a homogeneous mixture.

3. © 2018 Pearson Education Ltd.
Some Common Gases

4. © 2018 Pearson Education Ltd.
Properties That Define the
State of a Gas Sample
1)Temperature
2)Pressure
3)Volume
4)Amount of gas, usually expressed as
number of moles

•Having already discussed three of
these, we need to define pressure.

5. © 2018 Pearson Education Ltd.
Pressure
•Pressure is the
amount of force
applied to an area:
P =
F
A
•Atmospheric
pressure is the
weight of air per
unit of area.

6. © 2018 Pearson Education Ltd.
Units of Pressure
•Pascals: 1 Pa = 1 N/m
2
(SI unit
of pressure)
•Bar: 1 bar = 10
5
Pa = 100 kPa
•mm Hg or torr: These units
are literally the difference in
the heights measured in
millimeters of two connected
columns of mercury, as in the
barometer in the figure.
•Atmosphere:
1 atm = 760. torr = 760. mm Hg =
101.325 kPa = 1.10325 bar

7. © 2018 Pearson Education Ltd.
Standard Pressure
•Normal atmospheric pressure at sea
level is referred to as standard
atmospheric pressure.
•It is equal to
–1 atm
–760 torr (760 mmHg)
–101.325 kPa

8. © 2018 Pearson Education Ltd.
Ideal-Gas Equation
V  1/P (Boyle’s law).
V  T (Charles’s law).
V  n (Avogadro’s law).
•So far we’ve seen that
•Combining these, we get
V 
nT
P
•Finally, to make it an equality, we use
a constant of proportionality (R) and
reorganize; this gives the ideal-gas
equation: PV = nRT.

9. © 2018 Pearson Education Ltd.
R
•The ideal gas
constant makes
the equation and
equality, not only
a proportion.
•PV = nRT

10. © 2018 Pearson Education Ltd.
Density of Gases
If we divide both sides of the ideal-gas
equation by V and by RT, we get
n/V = P/RT.
Also: moles  molecular mass = mass
n  M = m.
If we multiply both sides by M, we get
m/V = MP/RT
and m/V is density, d; the result is:
d = MP/RT.

11. © 2018 Pearson Education Ltd.
Density and Molar Mass of a Gas
•To recap:
–One needs to know only the molecular mass,
the pressure, and the temperature to calculate
the density of a gas.
•d = MP/RT
–Also, if we know the mass, volume, and
temperature of a gas, we can find its
molar mass.
•M = mRT/PV

12. © 2018 Pearson Education Ltd.
Volume and Chemical Reactions
•The balanced equation tells us relative amounts of
moles in a reaction, whether the compared
materials are products or reactants.
•PV = nRT.
•So, we can relate volume for gases, as well.
•Reminder: Gay-Lussac’s law of combining
volumes.
gas data for n
(use PV = nRT)
moles of other
reactant/product
answer
(g, P/V/T)

13. © 2018 Pearson Education Ltd.
Dalton’s Law of
Partial Pressures
•If two gases that don’t react are combined in
a container, they act as if they are alone in
the container.
•The total pressure of a mixture of gases
equals the sum of the pressures that each
would exert if it were present alone.
•In other words,
P
t = P
1 + P
2 + P
3 + …

14. © 2018 Pearson Education Ltd.
Mole Fraction
•Because each gas in a mixture acts as if it is
alone, we can relate amount in a mixture to partial
pressures:



•That ratio of moles of a substance to total moles is
called the mole fraction, χ.

15. © 2018 Pearson Education Ltd.
Pressure and Mole Fraction
•The end result is

16. © 2018 Pearson Education Ltd.
Chapter 14

Chemical Kinetics
James F. Kirby
Quinnipiac University
Hamden, CT
Lecture Presentation

17. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Chemical Kinetics
•The speed at which reactions take place is
called reaction rate.
•The study of reaction rate is called
chemical kinetics.
•A step-by-step view of the change of
reactants to products is called a
mechanism.

18. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Factors That Affect Reaction Rates
1)Physical state of the reactants
2)Reactant concentrations
3)Reaction temperature
4)Presence of a catalyst

19. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Physical State of the Reactants
•The more readily the reactants collide, the
more rapidly they react.
•Homogeneous reactions (all gases or
liquids) are often faster.
•Heterogeneous reactions that involve
solids are slower; the faster of these
reactions occurs if the surface area is
increased; that is, a fine powder reacts
faster than a pellet or tablet.

20. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Reactant Concentrations
•Increasing reactant
concentration
generally increases
reaction rate.
•Since there are more
molecules, more
collisions occur.

21. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Reaction Rate
•Rate is a change in concentration over a time
period: Δ[ ]/Δt.
•Δ means ―change in.‖
•[ ] means molar concentration.
•t represents time.
•Types of rate measured:
–average rate
–instantaneous rate
–initial rate

22. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Following Reaction Rates
•Rate of a reaction is
measured using the
concentration change
for a reactant or a
product over time.
•In this example,
[C
4H
9Cl] is followed to
find average rate.
Rate = –Δ[C
4H
9Cl]/Δt

23. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Plotting Rate Data
•A plot of the data gives
more information
about rate.
•The slope of the curve at
one point in time gives the
instantaneous rate.
•The instantaneous rate at
time zero is called the
initial rate; this is often the
rate of interest to chemists.
•This figure shows
instantaneous and initial
rate of the earlier example.
Note: Reactions
typically slow down
over time!

24. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Relative Rates
•As was said, rates are followed using a reactant or a
product. Does this give the same rate for each reactant
and product?
•Rate is dependent on stoichiometry.
•If we followed use of C
4H
9Cl and compared it to
production of C
4H
9OH, the values would be the same.
Note that the change would have opposite signs — one
goes down in value, the other goes up.

25. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Relative Rates and Stoichiometry
•What if the equation is not 1:1?
•What will the relative rates be for:
2 O
3 (g) 3 O
2 (g)

26. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Determining Concentration Effect on Rate
•How do we determine what effect the concentration
of each reactant has on the rate of the reaction?
•We keep every concentration constant except for
one reactant and see what happens to the rate.
Then, we change a different reactant. We do this
until we have seen how each reactant has affected
the rate.
•For a Reaction: aA+bB --> cC
•Rate : V= k[A]m[B]n
•Where k is
•m and n are; m +n is the global order reaction

27. © 2018 Pearson Education Ltd.
Chemical
Kinetics
An Example of How
Concentration Affects Rate
•Experiments 1–3 show how [NH
4
+
] affects rate.
•Experiments 4–6 show how [NO
2
−
] affects rate.

28. © 2018 Pearson Education Ltd.
Chemical
Kinetics
An Example of How
Concentration Affects Rate
•In experiments 1–3, the concentration of NO
2
–
is
constant, so any change in rate is due to the NH
4
+
.
•As [NH
4
+
] doubles, the rate doubles.
•rate
2/rate
1 = [NH
4
+
]
2/[NH
4
+
]
1
•The math becomes 2 = 2
x
, so the order is 1.

29. © 2018 Pearson Education Ltd.
Chemical
Kinetics
An Example of How
Concentration Affects Rate
•In experiments 4–6, the concentration of NH
4
+
is
constant, so any change in rate is due to the NO
2
–
.
•As [NO
2
–
] doubles, the rate doubles.
•rate
2/rate
1 = [NO
2
–
]
5/[NO
2
–
]
4
•The math becomes 2 = 2
y
, so the order is 1.

30. © 2018 Pearson Education Ltd.
Chemical
Kinetics
An Example of How
Concentration Affects Rate
•Experiments 1–3 show how [NH
4
+
] affects rate.
•Experiments 4–6 show how [NO
2
–
] affects rate.
•Result: The rate law, which shows the relationship
between rate and concentration for all reactants:
Rate = k [NH
4
+
] [NO
2
–
]

31. © 2018 Pearson Education Ltd.
Chemical
Kinetics
More about Rate Law
•The exponents tell the reaction order with respect
to each reactant.
•In our example from the last slide:
Rate = k [NH
4
+
] [NO
2
–
]
•The order with respect to each reactant is 1. (It is
first order in NH
4
+
and NO
2
–
.)
•The reaction is second order (1 + 1 = 2; we just
add up all of the reactants’ orders to get the
overall reaction order).
•What is k? It is the rate constant. It is a
temperature-dependent quantity.

32. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Order ≠ Stoichiometry
•The order of the reaction must be
determined experimentally. It is not
necessarily related to the balanced
equation!

33. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Relative Value of k
•The rate constant is often used to
determine the relative rate of a reaction.
•For reactions with k ~ 10
9
or higher, the
reaction is considered fast.
•For reactions with k ~ 10 or lower, the
reaction is considered slow.

34. © 2018 Pearson Education Ltd.
Chemical
Kinetics
First-Order Reactions
•Some rates depend
only on one reactant to
the first power.
•These are first-order
reactions.
•The rate law becomes:
Rate = k [A]
•The conversion of
methyl isonitrile to
acetonitrile is a first-
order reaction.

35. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Relating k to [A] in a First-Order Reaction
•rate = k [A]
•rate = −Δ [A] / Δt
•So: k [A] = −Δ [A] / Δt
•Rearrange to: Δ [A] / [A] = − k Δt
•Integrate: ln ([A] / [A]
o) = − k t
•Rearrange: ln [A] = − k t + ln [A]
o
•Note: This follows the equation of a line:
y = m x + b
•So, a plot of ln [A] vs. t is linear.

36. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Evidence That Conversion of Methyl
Isonitrile to Acetonitrile
•The equation for the reaction:
CH
3NC CH
3CN
•It is first order. The plot of ln[CH
3CN] vs. time is
linear.
Rate = k [CH
3NC]

37. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Finding the Rate Constant, k
•Besides using the rate law, we can find the
rate constant from the plot of ln [A] vs. t.
•Remember the integrated rate law:
ln [A] = − k t + ln [A]
o
•The plot will give a line. Its slope will
equal −k.

38. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Half-Life
•Definition: The amount of time it takes for one-
half of a reactant to be used up in a chemical
reaction.
•First-Order Reaction:
–ln [A] = − k t + ln [A]
o
–ln ([A]
o/2) = − k t
½ + ln [A]
o
–− ln ([A]
o/2) + ln [A]
o = k t
½
–ln ([A]
o / [A]
o/2) = k t
½
–ln 2 = k t
½ or t
½ = 0.693/k

39. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Second-Order Reactions
•Some rates depend only on a reactant to
the second power.
•These are second-order reactions.
•The rate law becomes:
Rate = k [A]
2

40. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Solving the Second-Order
Reaction for A Products
•rate = k [A]
2

•rate = − Δ [A] / Δ t
•So, k [A]
2
= − Δ [A] / Δ t
•Rearranging: Δ [A] / [A]
2
= − k Δ t
•Using calculus: 1/[A] = 1/[A]
o + k t
•Notice: The linear relationships for first-
order and second-order reactions differ!

41. © 2018 Pearson Education Ltd.
Chemical
Kinetics
An Example of a Second-Order Reaction:
Decomposition of NO
2
•A plot following NO
2
decomposition shows
that it must be second
order because it is
linear for 1/[NO
2], not
linear for ln [NO
2].
•Equation:
NO
2 NO + ½ O
2

42. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Half-Life and Second-Order Reactions
•Using the integrated rate law, we can see
how half-life is derived:
–1/[A] = 1/[A]
o + k t
–1/([A]
o/2) = 1/[A]
o + k t
½
–2/[A]
o −1/[A]
o = k t
½
–t
½ = 1 / (k [A]
o)
•So, half-life is a concentration dependent
quantity for second-order reactions!

43. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Zero-Order Reactions
•Occasionally, rate is
independent of the
concentration of the
reactant:
Rate = k
•These are zero-order
reactions.
•These reactions are linear
in concentration.
•[A]
t = – k t + [A]
0

44. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Factors That Affect Reaction Rate
1)Temperature
2)Frequency of collisions
3)Orientation of molecules
4)Energy needed for the reaction to take
place (activation energy)

45. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Temperature and Rate
•Generally, as
temperature increases,
rate increases.
•The rate constant is
temperature
dependent: It
increases as
temperature increases.
•Rate constant doubles
(approximately) with
every 10 °C rise.

46. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Energy Needed for a Reaction to Take
Place (Activation Energy)
•The minimum energy needed for a reaction to take
place is called activation energy.
•An energy barrier must be overcome for a reaction to
take place, much like the ball must be hit to overcome
the barrier in the figure below.

47. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Transition State (Activated Complex)
•Reactants gain energy as the reaction
proceeds until the particles reach the
maximum energy state.
•The organization of the atoms at this
highest energy state is called the
transition state (or activated complex).
•The energy needed to form this state is
called the activation energy.

48. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Reaction Progress
•Plots are made to show
the energy possessed by
the particles as the
reaction proceeds.
•At the highest energy
state, the transition state
is formed.
•Reactions can be
endothermic or
exothermic after this.
•Rate constant depends
on the magnitude of E
a.

49. © 2018 Pearson Education Ltd.
Chemical
Kinetics
The Relationship between Activation
Energy and Temperature
•Arrhenius noted the relationship between
activation energy and temperature: k = Ae
−Ea/RT

•Activation energy can be determined graphically
by reorganizing the equation: ln k = −E
a/RT + ln A

50. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Law vs. Theory
•Kinetics gives what happens. We call the
description the rate law.
•Why do we observe that rate law? We
explain with a theory called a mechanism.
•A mechanism is a series of stepwise
reactions that show how reactants become
products.

51. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Reaction Mechanisms
•Reactions may occur all at once or through
several discrete steps.
•Each of these processes is known as an
elementary reaction or elementary
process.

52. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Molecularity
The molecularity of an elementary reaction tells
how many molecules are involved in that step of
the mechanism.

53. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Termolecular?
•Termolecular steps require three molecules
to simultaneously collide with the proper
orientation and the proper energy.
•These are rare, if they indeed do occur.
•These must be slower than unimolecular or
bimolecular steps.
•Nearly all mechanisms use only
unimolecular or bimolecular reactions.

54. © 2018 Pearson Education Ltd.
Chemical
Kinetics
What Limits the Rate?
•The overall reaction cannot occur faster than the
slowest reaction in the mechanism.
•We call that the rate-determining step.

55. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Catalysts
•Catalysts increase the rate of a reaction by
decreasing the activation energy of the reaction.
•Catalysts change the mechanism by which the
process occurs.

56. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Types of Catalysts
1)Homogeneous catalysts
2)Heterogeneous catalysts
3)Enzymes

57. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Homogeneous Catalysts
•The reactants and catalyst are in the same phase.
•Many times, reactants and catalyst are dissolved
in the same solvent, as seen below.

58. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Heterogeneous Catalysts
•The catalyst is in a
different phase than
the reactants.
•Often, gases are
passed over a solid
catalyst.
•The adsorption of
the reactants is
often the rate-
determining step.

59. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Enzymes
•Enzymes are
biological catalysts.
•They have a region
where the reactants
attach. That region
is called the active
site. The reactants
are referred to as
substrates.

60. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Lock-and-Key Model
•In the enzyme–substrate model, the
substrate fits into the active site of an
enzyme, much like a key fits into a lock.
•They are specific.

61. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Chapter 21

Nuclear Chemistry
James F. Kirby
Quinnipiac University
Hamden, CT
Lecture Presentation

62. © 2018 Pearson Education Ltd.
Chemical
Kinetics
The Nucleus
•Remember that the nucleus is composed of the
two nucleons, protons and neutrons.
•The number of protons is the atomic number.
•The number of protons and neutrons together is
the mass number.

63. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Isotopes
•Not all atoms of the same element have the
same mass, due to different numbers of
neutrons in those atoms.
•There are, for example, three naturally
occurring isotopes of uranium:
–Uranium-234
–Uranium-235
–Uranium-238

64. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Radioactivity
•Some nuclei change spontaneously,
emitting radiation. They are said to be
radioactive.
•We refer to these as radionuclides.
•There are several ways radionuclides can
decay into a different nuclide.
•We use nuclear equations to show how
these nuclear reactions occur.

65. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Nuclear Equations
•In chemical equations, atoms and
charges need to balance.
•In nuclear equations, atomic number
and mass number need to balance. This
is a way of balancing charge (atomic
number) and mass (mass number) on
an atomic scale.

66. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Most Common Kinds of Radiation Emitted
by a Radionuclide

67. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Types of Radioactive Decay
•Alpha decay
•Beta decay
•Gamma emission
•Positron emission
•Electron capture

68. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Alpha Decay
Alpha decay is the loss of an α-particle
(He-4 nucleus, two protons and two neutrons):





•Note how the equation balances:
–atomic number: 92 = 90 + 2
–mass number: 238 = 234 + 4

69. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Beta Decay
Beta decay is the loss of a β-particle (a
high-speed electron emitted by the nucleus):




Balancing: atomic number: 53 = 54 + (–1)
mass number: 131 = 131 + 0
β
0
–1 e
0
–1
or

70. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Gamma Emission
Gamma emission is the loss of a γ-ray,
which is high-energy radiation that
almost always accompanies the loss of
a nuclear particle:

71. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Positron Emission
Some nuclei decay by emitting a
positron, a particle that has the same
mass as, but an opposite charge to, that
of an electron:



Balancing: atomic number: 6 = 5 + 1
mass number: 11 = 11 + 0

72. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Electron Capture
An electron from the surrounding electron
cloud is absorbed into the nucleus during
electron capture.



Balancing: atomic number: 37 + (–1) = 36
mass number: 81 + 0 = 81

73. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Sources of Some Nuclear Particles
•Beta particles:


•Positrons:


•What happens with electron capture?

74. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Kinetics of Radioactive Decay
•Radioactive decay is a first-order process.
•The kinetics of such a process obey this
equation:

75. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Half-Life
•The half-life of such a process is


•Half-life is the time required for half of a radionuclide
sample to decay.

76. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Energy in Nuclear Reactions
•There is a tremendous amount of
energy stored in nuclei.
•Einstein’s famous equation, E = mc
2
,
relates directly to the calculation of
this energy.

77. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Energy in Nuclear Reactions
To show the enormous difference in energy for
nuclear reactions, the mass change associated
with the α-decay of 1 mol of U-238 to Th-234 is
–0.0046 g.
The change in energy, ΔE, is then
ΔE = (Δm)c
2

E = (–4.6 × 10
–6
kg)(3.00 × 10
8
m/s)
2

E = –4.1 × 10
11
J (410 billion kJ!!)

(Note: the negative sign means heat is released.)

78. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Nuclear Fission
•Commercial nuclear power plants use fission.
•Heavy nuclei can split in many ways. The equations below
show two ways U-235 can split after bombardment with a
neutron.

79. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Nuclear Fission
•Bombardment of the radioactive nuclide with a neutron
starts the process.
•Neutrons released in the transmutation strike other nuclei,
causing their decay and the production of more neutrons.
•This process continues in what we call a chain reaction.

80. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Nuclear Fission
•The minimum mass that must be present for a chain
reaction to be sustained is called the critical mass.
•If more than critical mass is present (supercritical mass),
an explosion will occur. Weapons were created by causing
smaller amounts to be forced together to create
this mass.

81. © 2018 Pearson Education Ltd.
Chemical
Kinetics
Nuclear Fusion
•When small atoms are combined, much energy is
released. This occurs on the Sun. The reactions are
often called thermonuclear reactions.
•If it were possible to easily produce energy by this
method, it would be a preferred source of energy.
•However, extremely high temperatures and pressures are
needed to cause nuclei to fuse.
•This was achieved using an atomic bomb to initiate fusion
in a hydrogen bomb. Obviously, this is not an acceptable
approach to producing energy.