Download CHEM102 December 11, 2003 1 K = 1 occurs when ∆G° = 0

CHEM102
December 11, 2003
Review for the Final Exam
Structure of the final exam.
6 problems + 1 extra credit problem
1.
2.
3.
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6.
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Short answers. Definitions and concepts
Thermodynamic calculations
Nomenclature of organic compounds
Functional groups in organic compounds; hybridization
Chemical equilibria: calculations of equilibrium constants; LeChatelier’s principle
Combined question: rates of chemical reactions, reaction mechanisms and thermodynamics
Extra credit problem: generic
2. Example of Problem 2. (see problem set on Thermodynamics from December 10)
1. Consider the following chemical reaction:
N2(g) + 2 O2(g) → 2NO2(g)
a) Using the tables provided below, calculate the standard enthalpy, standard entropy and standard free
energy change (∆
∆Ho, ∆So, and ∆Go, respectively) for the above reaction at 60 °C:
∆H0 = 2 ∆H0f(NO2(g))-( ∆H0f(N2(g)) + 2∆H0f(O2(g)) = 2 x 33.84-2 x 0 – 2 x 0 = 67.68 kJ/mol
∆S0 = 2 S0(NO2(g))-( S0(N2(g)) + 2S0(O2(g)) = 2 x 240.45 – 191.50-2 x 205.0 =-120.5 J/mol-K
∆G0 = ∆H0 - T ∆S0 = 67.68 - 333 (-0.1205) = 107.807 kJ/mol
b) Calculate the equilibrium constant for this reaction at 60 oC.
∆G0 = -RT lnK; K = e-∆G0/RT= e-107807/8.31x333 = 1.2 x 10-17
very small number- equilibrium favors reactants at this temperature
c) Will an increase in temperature produce an increase or decrease in the mole fraction of NO2(g) at
equilibrium? Explain.
The reaction is endothermic (∆H0<0); increasing temperature will favor formation of products
d) At what temperature if any will this reaction have an equilibrium constant equal to 1? You may assume
that ∆H°° and ∆S°° are temperature-independent.
K = 1 occurs when ∆G° = 0; therefore, ∆H° - T∆S° = 0, or T = ∆H° / ∆S° = 67680/(-120.5) = -561.66 K ,
which physically means that there is no temperature at which K can be equal to 1.
In other words, both the enthalpic and entropic terms are unfavorable (endothermic reaction with negative
entropy change- always non spontaneous at standard conditions).
3. Example of problem 3. (see workshop 2)
Give the IUPAC names for each of the following compounds:
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CHEM102
December 11, 2003
CH3
CH3CH2
CH2CH3
HC
C
CH2Cl
CH3
CH3CH2
CCH2CCH3
CH2CHCH2CH3
C
a)
CH3
4,4-Dimethyl-1-hexyne
H
c)
b)
C
H
1-Chloro-2-methyl-2-phenyl-butane
or (1-chloro-2-methyl)-2-butylbenzene
Cis-6-methyl-3-octene
4. Example of problem 4. The compound whose structure is shown below is acetylsalicylic acid, better known as
aspirin:
O
OH
O
O
CH3
a) Identify the functional groups in aspirine:
carboxyl
O
OH
phenyl
O
O
CH3
ester
b) Give the hybridization of each carbon atom in the aspirin molecule:
O
OH
O
O
CH3
2
all carbons except for the CH3 carbons- sp hybridized; CH3 carbons- sp3 hybridized
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CHEM102
December 11, 2003
5. Example of problem 5. Consider the following reaction at equilibrium:
A(g) V 2B (g)
The equilibrium constant of the reaction was measured at several temperatures and the following data
were obtained:
Temperature (ºC) [A], M [B], M
200
0.0125 0.843
300
0.171 0.764
400
0.250 0.724
a) Based on the data, calculate the equilibrium constant (Kc) at each temperature.
Kc = [B]2/[A]
At 200 oC, Kc = 0.8432/0.0125 = 56.85
At 300 oC, Kc = 0.7642/0.171 = 3.41
At 400 oC, Kc = 0.7242/0.250 = 2.10
b) Is the reaction exothermic or endothermic? Explain your answer
Since the equilibrium constant decreases with increasing the temperature, the reaction is exothermic, according to
the LeChatelier’s principle (formation of reactants is favored at higher temperatures).
6. Example of problem 6. Consider the following reaction:
H2(g) + 2 ICl (g) 2 HCl (g) + I2 (g)
The rate law for this reaction is first order in both H2 and ICl (meaning, Rate = k x [H2] x [ICl]).
a) Which of the following mechanisms are consistent with the observed rate law?
(a) 2ICl(g) + H2(g) 2HCl (g) + I2(g) (termolecular reaction
(b) H2(g) + ICl(g) HI(g) + HCl(g)
HI(g) + ICl(g) HCl(g) + I2(g)
(slow)
(fast)
(c) H2(g) + ICl(g) HI(g) + HCl(g)
HI(g) + ICl(g) HCl(g) + I2(g)
(fast)
(slow)
(d) H2(g) + ICl(g) HICl(g) + H (g)
H(g) + ICl(g) HCl(g) + I(g)
HICl(g) HCl(g) + I(g)
I(g) + I(g) I2(g)
(slow)
(fast)
(fast)
(fast)
Answer: b) and d) are consistent, since the stoichiometry of the slow (rate-determining) step will yield the rate law
equation.
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CHEM102
December 11, 2003
b) Based on the mechanism that is consistent with the rate law and on the thermodynamic tables
provided, calculate the thermodynamic parameters (∆
∆Hº, ∆Sº, and ∆Gº) of the individual steps, and
of the overall reaction at 25 ºC. Is the reaction spontaneous at 25 ºC?
For the overall reaction:
∆Hº= 2∆Hºf(HCl (g)) + ∆Hºf(I2 (g)) -∆Hºf(H2 (g)) -2∆Hºf(ICl (g)) = 2 (-92.3) + 62.25 –0 – 2 17.782 = -157.91
kJ/mol
∆Sº= 2Sº(HCl (g)) + Sº(I2 (g)) - Sº(H2 (g)) -2 Sº(ICl (g)) = 2 186.69 + 260.57 –130.58 – 2 247.44 = 8.49 J/mol-K
∆Gº= ∆Hº-T ∆Sº =157910 J/mol – 298.15 K x 8.49 J/mol-K = 155379 J/mol = 155.379 kJ/mol
The reaction is not spontaneous in the forward direction at 25 oC (∆Gº>0).
c) Calculate the equilibrium constant of the overall reaction at 25 ºC. Which direction will the
equilibrium be shifted upon increasing the temperature? Explain your answer.
∆Gº= -RT lnK; K = e-∆Gº/RT= e-155379/(8.31 x 298.15) = 5.8 x 10-28
This is a very small equilibrium constant. Since the overall reaction is exothermic, increasing temperature
will favor formation of the reactants, based on the LeChatelier’s principle.
Appendix 1. Thermodynamic quantities for selected substances at 298.15 K (25 oC)
Substance
Carbon
C(s, graphite)
CO(g)
CH3OH(l)
CH3COOH(l)
Hydrogen
H(g)
H2(g)
Iodine
I2(g)
I(g)
HI(g)
ICl(g)
Chlorine
Cl2(g)
HCl(g)
Nitrogen
N2(g)
N(g)
NO(g)
N2O(g)
Oxygen
O2(g)
O(g)
H2O(l)
H2O(g)
∆Hof
(kJ/mol)
∆Gof
(kJ/mol)
So
(J/mol-K)
0
-110.5
-238.6
-487.0
0
-137.2
-166.23
-392.4
5.69
197.9
126.8
159.8
217.94
0
203.26
0
114.60
130.58
62.25
106.6
25.94
17.782
19.37
70.16
1.30
-5.44
260.57
180.66
206.3
247.44
0
-92.30
0
-95.27
222.96
186.69
0
472.7
90.37
81.6
0
455.5
86.71
103.59
191.50
153.3
210.62
220.0
0
247.5
-285.83
-241.82
0
230.1
-237.13
-228.57
205.0
161.0
69.91
188.83
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December 11, 2003
Definitions. Define each of the following:
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Solution, solvent, solute, solubility
Saturated solution
Molarity
Molality
Mass percent
Mole fraction
Henry’s law
Raoult’s law
Colligative properties. Vapor pressure lowering, boiling-point elevation, freezing-point depression
Osmosis, Osmotic Pressure
Hypo-, hyper-, and isotonic solutions
Organic compounds
Hybridization: sp3, sp2, sp
σ-bond, π-bond
Hydrocarbons
Alkanes
Alkenes
Alkynes
Straight-chain hydrocarbons
Branched-chain hydrocarbons
Structural isomers
Cis-trans isomers
Functional groups
Aromatic compounds
Alcohols
Ethers
Esters
Amines
Carboxylic Acids
Carbonyl group
Addition reactions
Substitution reactions
Markovnikov’s rule
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Chemical kinetics
Reaction rate
Rate law
Rate constant
First order reaction
Expression for the first order reaction rate
Second order reaction
Expression for the second order reaction rate
Zero-order reaction
Expression for the zero-order reaction rate
Half-life of a reaction
Expression for half-life of the first-order reaction
Expression for half-life of the second-order reaction
Collision theory
Activation energy
Activated complex
Arrhenius equation
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CHEM102
December 11, 2003
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Reaction mechanism
Elementary step
Reactants
Products
Intermediates
Catalysts
Molecularity of a reaction
Unimolecular reaction
Bimolecular reaction
Termolecular reaction
Rate-determining (rate-limiting) step
Homogeneous catalysis
Heterogeneous catalysis
Chemical equilibrium
Equilibrium constant: definition, expression for a given chemical reaction
Homogeneous equilibrium
Heterogeneous equilibrium
Reaction quotient
Equilibrium position
Le Châtelier’s principle
Bronsted acid, Bronsted base
Conjugate acid-base pairs
Strong acid, strong base
Weak acid, weak base
Acid ionization constant, base ionization constant
Ion product of water
pH, pOH
Neutral, acidic, basic solutions
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Buffer solution
Henderson-Hasselbach equation
Solubility product
Molar solubility
Soubility (in grams per liter)
First Law of Thermodynamics
Heat of reaction, internal energy, enthalpy
Hess’s law
State function
System, surroundings, Universe
Thermochemical equations
Standard state
Standard enthalpy of formation
Standard enthalpy of reaction
Hess’s law
Second Law of Thermodynamics
Entropy
Spontaneous processes
Entropy of the system, surroundings, Universe
Standard entropy of reaction
Third Law of Thermodynamics
Gibbs Free Energy
Equilibrium constant in terms of standard Gibbs Free Energy
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