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Syllabus
Chemistry 102 Spring 2009
Sec. 501, 503 (MWF 9:10-10:00, 12:40-1:30)
RM 100 HELD
Professor: Dr. Earle G. Stone
Office: Room 123E Heldenfels (HELD)
Telephone: 845-3010 (no voice mail) or leave message at 845-2356
email: [email protected]
(put CHEM 101-Sec. # + subject in subject line of your email)
Office Hours: HELD 408: Wed. 8:00-10:50 AM
I.A. Esther Ocola
S.I. Leader: Analise Castellano
Suggested Course Materials:
“Chemistry and Chemical Reactivity,
Any Edition”, by Kotz
Ebook includes
Online tutorial
Solution manual
$45 per semester
Hardbound ~$160
Solution Manual ~$40
Online Tutor ~$45
Helpful
1. Dictionary of Chemistry
Or online dictionary
2. Mastering the
Fundamental Skills –
General Chemistry I
as a Second Language
Useful Later
As A Second Language Organic
Chemistry I by Klein, There is a O-chem
II and a Physics as a Second Language
(Algebra based or Calculus based) for
those who will have to take those classes.
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493
149
118
57
35
55
18
20
10
25
0
5
100%
30%
24%
12%
7%
11%
4%
4%
2%
5%
0%
1%
501
College
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Ag BICH, NUSC, GENE
GEST
Engineering
Education
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whoop
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2012
2011
2010
2009
2008
260
193
33
9
0
495
82
56
28
13
28
9
14
9
9
0
2
82
33%
22%
11%
5%
11%
4%
6%
4%
4%
0%
1%
33%
73% pre-something
27% widely diverging
academic foci
503
242
67
62
22
31
27
9
6
5
10
0
3
100%
28%
26%
13%
9%
11%
4%
2%
2%
4%
0%
1%
BIMS
Science
GEST
Ag BICH, NUSC, GENE
Engineering
Education
Geosciences
Liberal Arts
Agriculture other
Architecture
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BIMS
Week
1
Date
21-Jan
End of Chapter
Questions 6th
Syllabus
9
28-Jan
2-Feb
4-Feb
6-Feb
Chapter 19
1,5,29,39,49,59
11-Feb
10
Chapter 14
Sect 14.1-14.4
18-Feb
1,2,11,21,31,35,
49, 51,93
20-Feb
1,3,7,9,11,17,23,
27,41,43,47,53,55,
87,89
23-Feb
6
7
Chapter 16
4-Mar
9-Mar
11-Mar
13-Mar
8-Apr
10-Apr
Reading Day
13-Apr
Chapter 18
15-Apr
Exam # 3 Chapters 17, 18
17-Apr
20-Apr
1,5,9,19,23,25,33,
49,63
22-Apr
Chapter 20
1,3,5,13,25,31,45
27-Apr
15
6-Mar
8
1,3,9,15,19,33,35,
37,43,53,69,75,85,
99
24-Apr
27-Feb
2-Mar
13
14
25-Feb
Chapter 18
6-Apr
Exam #1 Chapters 14, 19
Chapter 15
61,71,93,107,109
30-Mar
3-Apr
13-Feb
5
Chapter 17
1-Apr
12
16-Feb
25-Mar
Chapter
27-Mar
11
9-Feb
4
End of Chapter
Questions 6th
14-Mar through 22-Mar Spring Break
Date
23-Mar
30-Jan
3
Week
23-Jan
26-Jan
2
Chapter
29-Apr
Exam # 4 Chapter 20
1-May
Reading Day
Exam #2 Chapters 15, 16
16
4-7 May
Reading Days
Chapter 17
17
11-May
Final Sect 501 8-10 a.m.
Final Sect 503 10:30 a.m.-12:30 p.m
7,11,15,23,27,35
May 11, Monday 8-10 a.m. MWF 9:10-10 a.m.
May 11, Monday 10:30 a.m.-12:30 p.m. MWF 12:40-1:30 p.m.
Grading:
Your grade will be based on
•Four one-hour examinations (each worth 200 points)
•A final examination (400 points)
Major Examination Schedule Spring 2009:
Wed. Feb. 11 Major Exam No.1
Mon. Mar. 9 Major Exam No.2
Wed. Apr. 15 Major Exam No.3
Wed. Apr. 29 Major Exam No.4
Final Exams
Section 501
Mon. May 11 8:00 to 10:00
Section 503
Mon. May 11 10:30 to 12:30
What you are used to
The way the real world works
+3%
80%
72%
70%
60%
C
B
60%
48%
D,F,Q,W
90%
84%
A
D
C
2% 16%
B
100%
96%
A
50% 84%
98% Percentile Rank
A is > average + 1 s
B is > average but less than average + 1 s
C = > average - 1 s but less than average
The mere formulation of a problem is far more often essential
than its solution, which may be merely a matter of mathematical
or experimental skill. To raise new questions, new possibilities, to
regard old problems from a new angle requires creative
imagination and marks real advances in science.
Albert Einstein
Problem - A situation that presents difficulty, uncertainty, or perplexity:
Question - A request for data: inquiry, interrogation, query.
Answer - A spoken or written reply, as to a question.
Solution - Something worked out to explain, resolve, or provide a method
for dealing with and settling a problem.
1. Numbers – Significant Figures, Rounding Rules, Accuracy, Precision, Statistical
Treatment of the Data
2. Units – 5 of the 7
1. Time – seconds
2. Length – Meters
Density?
3. Mass – grams
Molecular Weight (Mass)
4. Amount – Moles
Mole Ratio, Molarity, molality
5. Temperature – Kelvins
3. Vocabulary – Approximately 100 new terms or words and applying new or
more rigid definitions to words you may already own.
4. Principles (Theories and Laws) – Stoichiometry, Quantum Theory, Bonding,
Chemical Periodicity, Solutions, Thermodynamics, Intermolecular Forces, Gas
Laws, Collogative Properties, Kinetics, Equilibrium, Electrochemistry
cp = q/mDT
DG = DH – TDS
PV = nRT
DT = Kmi
rate = k[A]m[B]n
∆E = q + w
Eocell = Ecathode = Eanode
[C]c[D]d
%yield = actual/theoretical * 100% K = [A]a[D]b
c (ms-1)
E = n =
l (m)
Chemistry Review
The prediction of Chemical Reaction in general relies on
1. The Law of Conservation of Mass – this leads to
•
•
Stoichiometry that allows us to compare apples and oranges
Equilibrium predictions of reversible reactions which leads to
• Kinetics allowing us to determine how fast the reaction will occur
2. The Law of Conservation of Energy – this leads to
Thermodynamics which is stated in 3 laws
1. First Law – the energy of the Universe is constant
Some Thermodynamic Terms
Thermodynamics - The study of the relationship between heat, work, and other forms
of energy.
Thermochemistry - A branch of thermodynamics which focuses on the study of heat
given off or absorbed in a chemical reaction.
Temperature - An intensive property of matter; a quantitative measurement of the degree
to which an object is either "hot" or "cold".
1.There are 3 scales for measuring temperature
•Fahrenheit - relative
•32 F is the normal freezing point temperature of water; 212 F is the normal boiling
point temperature of water.
•Celsius (centigrade) - relative
•0 C is the normal freezing point temperature of water; 100 C is the normal boiling
point temperature of water.
•Kelvin - absolute
•0 K is the temperature at which the volume and pressure of an ideal gas extrapolate
to zero.
Some Thermodynamic Terms
Heat (q) - A form of energy associated with the random motion of the elementary
particles in matter.
Heat capacity - The amount of heat needed to raise the temperature of a defined
amount of a pure substance by one degree.
Specific heat - The amount of heat needed to raise the temperature of one gram of a
substance by 1 C (or 1 K)
•SI unit for specific heat is joules per gram-1 Kelvin-1 (J/g-K)
Calorie - The specific heat of water = 4.184 J/g-K
Molar heat capacity - The amount of heat required to raise the temperature of one
mole of a substance by 1 C (or 1 K)
•SI unit for molar heat capacity is joules per mole-1 Kelvin-1 (J/mol-K)
Btu (British thermal unit) - The amount of heat needed to raise the temperature of
1 lb water by 1 F.
NOTE: The specific heat of water (4.184 J/g-K) is very large relative to other
substances. The oceans (which cover over 70% of the earth) act as a giant "heat sink,"
moderating drastic changes in temperature.
Our body temperatures are also controlled by water and its high specific heat.
Perspiration is a form of evaporative cooling which keeps our body temperatures from
getting too high.
Some Thermodynamic Terms
Latent Heat versus Sensible Heat
Sensible heat - Heat that can be detected by a change in the temperature of a
system.
Latent heat - Heat that cannot be detected because there is no change in
temperature of the system. e.g. The heat that is used to melt ice or
to evaporate water is latent heat.
There are two forms of latent heat:
•Heat of fusion - The heat that must be absorbed to melt a mole of a solid.
•e.g. melting ice to liquid water
•Heat of vaporization - The heat that must be absorbed to boil a mole of a liquid.
•e.g. boiling liquid water to steam
Some Thermodynamic Theories
Caloric Theory of Heat
•Served as the basis of thermodynamics.
•Is now known to be obsolete
•Based on the following assumptions
•Heat is a fluid that flows from hot to cold substances.
•Heat has a strong attraction to matter which can hold a lot of heat.
•Heat is conserved.
•Sensible heat causes an increase in the temperature of an object when it flows into
the object.
•Latent heat combines with particles in matter (causing substances to melt or boil)
•Heat is weightless.
The only valid part of the caloric theory is that heat is weightless.
Heat is NOT a fluid, and it is NOT conserved.
Some Thermodynamic Theories
Kinetic Theory of Heat
1. Divides the universe into two parts:
A.
System. - The substances involved in the chemical and
physical changes under investigation: In chemistry lab, the
system is the REACTANTS inside the beaker.
B.
Surroundings - Everything not included in the system, i.e.
the rest of the universe.
2. A BOUNDARY separates the system and the surroundings
from each other and can be tangible or imaginary.
A. Heat is something that is transferred back and forth across
boundary between a system and its surroundings
B. Heat is NOT conserved.
Some Thermodynamic Theories
The kinetic theory of heat is based upon the last postulate in the kinetic molecular
theory which states that the average kinetic energy of a collection of gas
particles is dependent only upon the temperature of the gas.
where R is the ideal gas constant (0.08206 L-atm/mol-K) and T is temperature
(Kelvin) The kinetic theory of heat can be summarized as follows:
When heat enters a system, it causes an increase in
the speed at which the particles in the system move.
•
The set of conditions that specify all of the properties of the system is called the
thermodynamic state of a system.
•
For example the thermodynamic state could include:
– The number of moles and identity of each substance.
– The physical states of each substance.
– The temperature of the system.
– The pressure of the system.
Standard States and Standard
Enthalpy Changes
1. Thermochemical standard state conditions
• The thermochemical standard T = 298.15 K.
• The thermochemical standard P = 1.0000 atm.
– Be careful not to confuse these values with STP.
2. Thermochemical standard states of matter
• For pure substances in their liquid or solid phase the
standard state is the pure liquid or solid.
• For gases the standard state is the gas at 1.00 atm of
pressure.
• For gaseous mixtures the partial pressure must be 1.00 atm.
• For aqueous solutions the standard state is 1.00 M
concentration.
Some Thermodynamic Terms
1.
State Functions are independent of pathway:
– T (temperature), P (pressure), V (volume), DE
(change in energy), DH (change in enthalpy – the
transfer of heat), and S (entropy)
2.
Examples of non-state functions are:
– n (moles), q (heat), w (work)
The Three Laws of Thermodynamics
There are two basic ideas of importance
for thermodynamic systems.
1. Chemical systems tend toward a state
of minimum potential energy.
2. Chemical systems tend toward a state
of maximum disorder.
The First Law of Thermodynamics
• The first law is also known as the
Law of Conservation of Energy.
Energy is neither created nor destroyed in
chemical reactions and physical changes.
•The energy of the universe does not change.
•The energy in a system may change, but it must be complemented by a change in the
energy of its surroundings to balance the change in energy.
The term internal energy is often used synonymously with the energy of a system. It is the
sum of the kinetic and potential energies of the particles that form the system. The last
postulate in the kinetic molecular theory states that the average kinetic energy of a collection
of gas particles is dependent only upon the temperature of the gas.
The First Law of Thermodynamics
Esys = KEsys + PEsys
1.
2.
KE – kinetic energy: translational, rotational, vibrational
PE – energy stored in bonds (Bond energy)
The First Law of Thermodynamics
If a system is more complex than an ideal gas, then the internal energy must
be measured indirectly by observing any changes in the temperature of the
system. The change in the internal energy of a system is equal to the
difference between the final and initial energies of the system:
The equation for the first law of thermodynamics can be rearranged to show
the energy of a system in terms of the energy of its surroundings.
This equation indicates that the energy lost by one must equal the energy
gained by the other:
The First Law of Thermodynamics
The energy of a system can change by the transfer of work and or
heat between the system and its surroundings. Any heat that is taken,
given off, or lost must be complemented by an input of work to make
up for the loss of heat. Conversely, a system can be used to do any
amount of work as long as there is an input of heat to make up for
the work done.
This equation can be used to explain the two types of heat that can be
added to a system:
1. Heat can increase the temperature of a system. This is sensible heat.
2. Heat that does ONLY WORK on a system is latent heat.
The First Law of Thermodynamics
1. Exchange of heat (q) Endothermic and exothermic
2. Work is performed (w)
DE = q + w
Solids, Liquids, Solutions
Changes in volume are negligible
Therefore w is effectively zero
DE = q + 0 = DH
DH is change in enthalpy which is the
transfer of heat and is measured
experimentally by determining changes
in temperature.
Gases
Why only gases?
Because changes in volume
results in work
w = Fd
F = Pressure x Area d = Dh
W = P (A Dh) = DV
h
heat transfer in
(endothermic), +q
heat transfer out
(exothermic), -q
SYSTEM
∆E = q + w
w transfer in
(+w)
Compression of system
w transfer out
(-w)
Expansion of system
By convention except for some engineers whose frame
of reference is the work done on the surroundings.
hi
hf
A(hf-hi)<0 DV
hi hf
A(hf-hi)>0 DV
w = -PDV
DE =DH + w = DH – PDV = DH – D(PV)
DE = DH – D(PV)
Constant Volume
Constant Pressure
w = -PDV
DV = 0
DE = DH
Check the temperature change
Apply some stoichiometry
And the Ideal Gas Law
PV=nRT
D(PV)=D(nRT)
Hold Temperature constant k1
D(PV)=D(nRk1)
Combine constants and
multiply through by -1
-D(PV) = -R1Dn
w = -PDV = -R1Dn
DE = DH + w = DH - R1Dn
DE
DH
Dn
DE = DH
exothermic
No change
DE = DH
endothermic
No change
DE > DH
exothermic
increase
DE > DH
endothermic
decrease
DE < DH
exothermic
decrease
DE < DH
endothermic
increase
Thermochemical Equations
• Thermochemical equations are a balanced chemical reaction
plus the DH value for the reaction.
– For example, this is a thermochemical equation.
C5 H12( )  8 O 2(g)  5 CO 2(g)  6 H 2 O (  )  3523 kJ
1 mole
8 moles
5 moles
6 moles
• The stoichiometric coefficients in thermochemical equations must
be interpreted as numbers of moles.
• 1 mol of C5H12 reacts with 8 mol of O2 to produce 5 mol of CO2,
6 mol of H2O, and releasing 3523 kJ is referred to as one mole of
reactions.
Thermochemical Equations
Write the thermochemical equation for
CuSO4(aq) + 2NaOH(aq)
50.0mL of 0.400 M CuSO4 at 23.35 oC
50.0mL of 0.600 M NaOH at 23.35 oC
Tfinal 25.23oC
CH2O = 4.184 J/goC
Density final solution = 1.02 g/mL
Cu(OH)2(s) + Na2SO4(aq)
The Second Law of Thermodynamics
• The second law of thermodynamics states, “In
spontaneous changes the universe tends
towards a state of greater disorder.”
• Spontaneous processes have two requirements:
1. The free energy change of the system must be
negative.
2. The entropy of universe must increase.
•
Fundamentally, the system must be capable of doing
useful work on surroundings for a spontaneous process
to occur.
Changes in DS are usually quite small compared to DE and DH.
Notice that DS has units of only a fraction of a kJ while DE and
DH values are much larger numbers of kJ.
The Second Law of Thermodynamics
Entropy (S) - A measure of the disorder in a
system. Entropy is a state function.
where k is a proportionality constant equal to the ideal gas
constant (R) divided by Avogadro's number (6.022 x 10-23)
and lnW is the natural log of W, the number of equivalent
ways of describing the state of a system.
In this reaction, the number of ways of describing a system is
directly proportional to the entropy of the system.
The Second Law of Thermodynamics
Number of Equivalent Combinations for Various Types of Poker Hands
Hand
W
ln W
Royal flush (AKQJ10 in one suit)
4
1.39
36
3.58
624
6.44
3,744
8.23
Flush (five cards in the same suit)
5,108
8.54
Straight (five cards in sequence)
10,200
9.23
Three of a kind
Two pairs
One pair
No pairs
54,912
123,552
1,098,240
1,302,540
2,598,960
10.91
11.72
13.91
14.08
Straight flush (five cards in
sequence in one suit)
Four of a kind
Full house (three of a kind plus a
pair)
Total
The Second Law of Thermodynamics
Entropy of Reaction (DS)
•The difference between the sum of the entropies of
the products and the sum of the entropies of the
reactants:
In the above reaction, n and m are the coefficients
of the products and the reactants in the balanced
equation.
As with DH, entropies have been measured and tabulated in
Appendix L as So298.
When:
DS > 0 disorder increases (which favors spontaneity).
DS < 0 disorder decreases (does not favor spontaneity).
The Second Law of Thermodynamics
Natural processes that occur in an isolated system are spontaneous
when they lead to an increase in the disorder, or entropy, of the
system.
Isolated system - System in which neither heat nor work can be
transferred between it and its surroundings. This makes it possible to
ignore whether a reaction is exothermic or endothermic.
If DSsys > 0, the system becomes more disordered through the
course of the reaction
If DSsys < 0, the system becomes less disordered (or more
ordered) through the course of the reaction.
The Second Law of Thermodynamics
There are a few basic principles that should be remembered to
help determine whether a system is increasing or decreasing in
entropy.
•Liquids are more disordered than solids.
•WHY? - Solids have a more regular structure than
liquids.
•Gases are more disordered than their respective liquids.
•WHY? - Gases particles are in a state of constant
random motion.
•Any process in which the number of particles in the
system increases consequently results in an increase in
disorder.
• In general for a substance in its three states of matter:
Sgas > Sliquid > Ssolid
The Second Law of Thermodynamics
Does the entropy increase or decrease for the following reactions?
•
•
•
•
Answers:
•INCREASES - The number of particles in the system increases, i.e. one particle decomposes into two. In
addition, one of the products formed is a gas which is much more disordered than the original solid.
•DECREASES - The number of particles in the system decreases, i.e. there are four moles of gas reactants and
only 2 moles of gas products.
•INCREASES - The number of particles in the system increases, i.e. the single reactant dissociates into two ion
particles. In addition, the ions in the ionic solid are organized in a rigid lattice structure whereas the ions in aqueous
solution are free to move randomly through the solvent.
•DECREASES - The reactant changes from a gas to a liquid, and gases are more disordered than their respective
liquids.
•
•
Entropy, S
The Third Law of Thermodynamics states, “The entropy of a pure, perfect,
crystalline solid at 0 K is zero.”
This law permits us to measure the absolute values of the entropy for
substances.
– To get the actual value of S, cool a substance to 0 K, or as close as
possible, then measure the entropy increase as the substance heats from
0 to higher temperatures.The coldest place in nature is the depths of
outer space. There it is 3 degrees above Absolute Zero.
– Notice that Appendix L has values of S not DS.
Predicted 1924......Created 1995
Bose-Einstein Condensation in a gas: a
new form of matter at the coldest
temperatures in the universe...
A. Einstein S. Bose
Cornell and Wieman cooled a small
sample of atoms down to only a
few billionths (0.000,000,001) of a
degree above Absolute Zero
Entropy, S
BEC
Entropy and Temperature
S increases
slightly with T
S increases a large
amount with
phase changes
Entropy, S
• Entropy changes for reactions can be
determined similarly to DH for reactions.This is
only true, i.e. conserved, for the system. This is
not included for the surroundings.
DS
o
298
=  nS
o
products
n
-  nS
o
reactants
n
Entropy, S
• Calculate the entropy change for the following reaction at 25oC.
Use Appendix L.
2 NO2(g)

 N 2 O 4(g)
You do it!
Entropy, S
• Calculate DSo298 for the reaction below. Use Appendix L.
3 NO(g  
 N 2 O (g   NO2(g 
You do it!
Free Energy Change, DG, and
Spontaneity
•
•
In the mid 1800’s J. Willard Gibbs determined the relationship
of enthalpy, H, and entropy, S, that best describes the maximum
useful energy obtainable in the form of work from a process at
constant temperature and pressure.
– The relationship also describes the spontaneity of a system.
The relationship is a new state function, DG, the Gibbs Free
Energy.
DG = DH-TDS
at constant T and P
Free Energy Change, DG, and
Spontaneity
•
•
The change in the Gibbs Free Energy, DG, is a reliable
indicator of spontaneity of a physical process or chemical
reaction.
– DG does not tell us how quickly the process occurs.
• Chemical kinetics, the subject of Chapter 16, indicates
the rate of a reaction.
Sign conventions for DG.
– DG > 0 reaction is nonspontaneous
– DG = 0 system is at equilibrium
– DG < 0 reaction is spontaneous
Free Energy Change, DG, and Spontaneity
• Changes in free energy obey the same type of
relationship we have described for enthalpy, DH,
and entropy, DS, changes.
DG
0
298
=  nDG
n
0
products
-  nDG
n
0
reactants
Free Energy Change, DG, and Spontaneity
• Calculate DGo298 for the reaction in Example 15-8. Use
Appendix L.
C3H8( )  5 O 2(g)  3 CO2(g)  4 H 2O( )
You do it!
The Temperature Dependence of
Spontaneity
• Free energy has the relationship
DG = DH -TDS.
• Because 0 ≤ DH ≥ 0 and 0 ≤ DS ≥ 0, there
are four possibilities for DG.
Forward reaction
DH
<0
<0
>0
>0
DS
>0
<0
>0
<0
DG
<0
T dependent
T dependent
>0
spontaneity
at all T’s.
at low T’s.
at high T’s.
Nonspontaneous
at all T’s.
Equilibrium
Spontaneous
Non Spontaneous
Spontaneity is favored when
DH < 0
DS > 0
DG = DH -TDS
DG
DH
..
)
)
High T
..
..
)
..
..
..
)
)
..
)
)
Low T
..
)
..
DS
)
..
)
DG = 0
DG < 0
DG > 0
The Temperature Dependence of
Spontaneity
•
•
•
Calculate DSo298 for the following reaction
C3H8(g) + 5 O2(g) )  3 CO2(g) + 4 H2O(g)
We know that DHo298= -2219.9 kJ,
and that DGo298= -2108.5 kJ.