Thermodynamics:

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Thermodynamics
The science of the transformation of energy
during chemical and physical processes. energy-
The capacity to do work.   system- Part of
universe of interest or under study.   surrounding
s- The universe outside of the system.
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Types of Systems
open system- System that exchanges mass and
energy(usually heat).   closed system- System
that allows the exchange of energy(heat) but not
mass.   isolated system- System that does not
allow the transfer of energy or mass.
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Heat vs. Work
Heat(q) Transfer of thermal energy between a
system and its surroundings.   Work(w) - Transfer
energy between a system and its surroundings when
a force is applied over a distance.
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Internal Energy(U)
Internal energy is the sum of the kinetic and
potential energies.   U state function. Can not
measure directly. Can measure change.
?U Ufinal Uinitial
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Internal Energy(U)
?U q w ?U energy enters system -
energy leaves system q system absorbs
heat - system emits heat w work done on
system - system does work
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First Law of Thermodynamics
Total energy of the universe is constant. Energy
is neither created or destroyed. U constant for
an isolated system.
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Enthalpy(H)
H Measure of heat content. Can not measure
directly. Can measure change(?H). ?H expressed
in J or kJ. ?H - (exothermic) ?H
(endothermic)   When pressure is constant ?H
q.
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Enthalpy(H) cont
Ex Consider a container of gas. If 100. kJ of
heat is added, the gas will expand if allowed.
When the gas expands, 10 kJ of work is done.
Calculate ?U and ?H.
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Spontaneous Processes and Entropy
Spontaneous process is a process which has a
tendency to occur without external influence. Ex
rusting of iron.   Originally spontaneous
processes were defined as exothermic processes.
Not so!
Spontaneous reactions occur because matter and
energy tend to become more disordered.
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Entropy
Entropy(S) Thermodynamic property related to
the degree of disorder. Can not measure S
directly. Can measure ?S. ?S  SFINAL -
SINITIAL ?S system becoming more
disordered. -?S system becoming less disordered
or more organized. Ex Solid ? Liquid ?S
System is becoming less ordered.
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Entropy cont
Second Law of Thermodynamics In any spontaneous
process there is always an increase in the
entropy of the universe.   For a spontaneous
process ?Suniverse gt 0. Where ?Suniverse ?SSYS
?SSUR ?SSYS entropy change for the
system. ?SSUR entropy change for the surroundings
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Spontaneity is Temperature Dependant
For a spontaneous process ?Suniverse gt 0.
?Suniverse ?SSYS ?SSUR
Consider a system of Liquid Water
H2O(l) ? H2O(g)
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Quantifying ?S
qrev Heat changes for a reversible process.
Process that can be reversed in direction by a
small change in a system property. T
temperature in Kelvin
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Entropy
Assuming a constant pressure.
T temperature in Kelvin ?S has units of kJ/K or
J/K.
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Entropy cont
Third Law of Thermodynamics The entropy for a
pure perfect crystal is zero at absolute zero(O
K).
Ex Calculate the change in entropy for the
melting of ice at 0 ?C. H2O(s) ? H2O(l) ?H
6.0 kJ
To determine if a process is spontaneous, can not
measure ?Suniverse and ?SSUR.
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Gibbs Free Energy(G)
G H TS since we measure change ?G ?H
T?S ?G lt 0 process is spontaneous ?G gt 0
process is not spontaneous ?G 0 process is at
equilibrium
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Gibbs Free Energy(G)
Ex Calculate the change in Gibbs free energy
for the melting of ice if ?H 6.0 kJ and ?S
22. J/K at 0.0?C. H2O(s) ? H2O(l)
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?G Temperature Dependant
For ?G ?H T?S
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?G cont
 Ex Consider the following process H2O(l) ?
H2O(g)   where ?H 44.0 kJ and ?S 118
J/K Calculate ?G at 90 ?C, 100 ?C, and 110 ?C.
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Standard Gibbs Free Energy Change(?G?)
"?" Indicates standard conditions (1 atm and
25?C) Can determine ?G? from ?H? and ?S?. ?G?
?H? T?S?
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Enthalpies of Formation(?Hf? )
Can use ?Hf? values to determine ?H? of any
reaction.   ?H? ??Hf?(prod) - ??Hf?(react)
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Enthalpies of Formation(?Hf? )
Ex Calculate ?H? for the following 2NH3(g)
3Cl2(g) ? N2(g) 6HCl(g) Given
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Standard Entropies
Similar method used for ?S?. ?S? ??S?(prods) -
??S?(reacts)   Ex Calculate ?S? for the
following N2O4(g) ?2NO2(g) ?S? ? Given
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Calculating ?G?
Ex Calculate ?G? for the following N2(g)
3H2(g) ?2NH3(g) ?G? ?   Given
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Standard Free Energies of Formation
Similar method can be used for ?G?. ?G?
??Gf?(prod) - ??Gf?(reacts)
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Relationship between ?G and ?G?
?G ?G? RTInQ   R Gas Constant(8.314
J/K?mole) T temperature in Kelvin Q reaction
quotient
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?G? cont
Ex Calculate ?G for the following CO(g)
2H2(g) ? CH3OH(l)   at 25 ?C if ?G? -29.0 kJ
and p(CO) is 5.0 atm and p(H2) is 3.0 atm.
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?G and Equilibrium
For systems at equilibrium ?G? -RTInK   Can
now determine the equilibrium constant from ?G?.
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Electrochemistry
Oxidation Number- Arbitrary number which
indicates the number of electrons a species will
lose or gain.
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Assigning Oxidation Numbers
1. Oxidation number of a free element or diatomic
molecule is zero. Ex Na(s), Cu(s), H2(g), F2(g)
2. In most cases the oxidation number of hydrogen
is 1, oxygen is -2, and fluorine is -1 when
combined with another element.
3. The sum of the oxidation numbers of each of
the elements in a molecule or ion must equal the
charge.
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Using Oxidation Numbers

• Ex
• Zn(s) Cu2 ? Zn2 Cu(s)

Zn(s) oxidized(lost electrons). Cu2(aq)
reduced(gained electrons).
Ex 2H2(g) O2(g) ? 2H2O(l)
H2(g) oxidized(lost electrons). O2(g)
reduced(gained electrons).
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Reduction-Oxidation Reactions(REDOX)
Oxidation- Process in which oxidation state of an
element increases. Species loses
electrons.   Reduction- Process in which
oxidation state of an element decreases. Species
gains electrons.
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REDOX cont
OXIDATION Zn(s) ? Zn2 2e- REDUCTION Cu2
2e- ? Cu(s) REDOX Zn(s) Cu2 ? Zn2
Cu(s)
Zn(s) oxidized/reducing agent. Cu2(aq)
reduced/oxidizing agent.
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Writing Balanced Redox Reactions
Oxidation and reduction reactions occur together.
Occur in acidic or basic medium. Ex (acidic) SO32
- MnO4- ? SO42- Mn2 STEP 1 Identify
the oxidized and reduced species and write the
corresponding half reactions.
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Writing Balanced Redox Reactions cont
STEP 2 Balance each of the half reactions. First
atoms other than H and O. Balance O atoms by
adding H2O molecules and then balance H atoms by
adding H ions.
STEP 3 Balance the number of electrons.
STEP 4 Add both half reactions and simplify.
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Writing Balanced Redox Reactions cont
Balance the following redox reaction which occurs
in a basic medium.  CrO42- S2-? Cr(OH)3(s)
S(s)   NOTE In basic medium add an equal
number of OH- ions to both sides to neutralize H
ions. OH- H ? H2O
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Electrochemical Cells
Consider, Zn(s) Cu2(aq) ? Zn2(aq)
Cu(s)   Zn(s) ? Zn2(aq) 2e- Cu2(aq) 2e-
? Cu(s)
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Electrochemical Cells
Electrode- Strip of metal. Half cell-Strip of
metal in contact with its ion. Salt bridge-
Allows passage of charge but not reactants.
Anode Zn(s) ? Zn2(aq) 2e- Cathode Cu2(aq)
2e- ? Cu(s)
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Electrochemical Cells cont

-
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Describing Electrochemical Cells(Cell Diagram)
ANODE(OXIDATION) on left. CATHODE(REDUCTIO
N) on right. ? indicates phase
change. ?? indicates salt bridge.   M(s) ?Mn(aq)
??Mn(aq)?M(s) (anode)
(cathode)
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Electrode Potentials
The voltage recorded by an electrochemical cell
is referred to as the electromotive force(emf)
and given the following symbol. Ecell
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Standard Electrode Potentials
Reaction occurs under standard conditions 25?C
and all substances are at unit concentration (1 M
for all ions and 1 atm for all gases). Measured
potential given the following symbol.
The cell potential can be broken up into two
components or half cells.
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Standard Hydrogen Electrode(SHE)
2H(aq)(1 M) 2e- ? H2(g)(1 atm) Eo 0 V
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SHE/Cu
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Zn/SHE
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Standard Reduction Potentials
Can now determine the potential for each half
cell. Tabulated and are referred to as standard
reduction potentials.
ExCalculate the standard cell potential for the
following Cd(s) Cu2(aq) ? Cd2(aq)
Cu(s)   Given Ereduc?(V) Cu2(aq)
2e- ? Cu(s) 0.337 V Cd2(aq) 2e- ?
Cd(s) -0.403 V
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Standard Reduction Potentials cont
Ex Calculate the standard cell potential for the
following 2Al(s) 3Cu2(aq) ? 3Cu(s)
2Al3   Given Ered?(V) Cu2(aq) 2e-
? Cu(s) 0.337 V Al3(aq) 3e- ?Al(s)
-1.66 V
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Predicting Spontaneous Redox Reactions
welectrical nF Ecell ?G -nFEcell
welectrical electrical work n moles of
electrons transferred F Faraday constant(96485
C/mole) Ecell Voltage of cell. NOTE 1 J 1
C?V ?G Gibbs Free Energy
?Glt0 process is spontaneous ?Ggt0 process is not
spontaneous
Ecell process is spontaneous Ecell -
process is not spontaneous
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Zn(s) Cu2 ? Zn2 Cu(s)
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Cell Potential as a Function of Concentration
Nernst Equation
Where if aA bB ? cC dD
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Cell Potential as a Function of Concentration
cont
Ex Calculate Ecell for the following voltaic
cell. Zn(s)?Zn2(0.15 M)??Cu2(0.20 M) ?Cu(s)
Ex2 Calculate Ecell for the following voltaic
cell. Pt(s)?H2(g)(0.5 atm)?H(0.50 M)??Ag(0.20
M)?Ag(s)
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Electrolysis
Electricity is used to cause a non-spontaneous
redox reaction to occur. Ex Electrolysis of
molten sodium chloride. 2Na(aq) 2Cl-(aq) ?
2Na(s) Cl2(g) Ecell? -4.07 V NOT
SPONTANEOUS!!!
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Electrolysis of Water
Ex 2H2O(l) ? 2H2(g) O2(g) Ecell? -1.23V
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Quantitative Electrolysis
Direct relationship between the amount of
electricity used and the amount of products
obtained from an electrolysis.   E IR E
voltage I current(in amps) R resistance(in
ohms) 1 mole of electrons 1 F(faraday) 1 F
96500 C 1A 1 C?s-1
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Quantitative Electrolysis Examples
Ex 2Na 2Cl- ? 2Na(s) Cl2(g)   In the
electrolysis of NaCl, how much Na(s) and Cl2(g)
is produced by a current of 0.500 A in 10.0
minutes.
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Quantitative Electrolysis Examples
Ex 2 Cu2 H2O ? O2(g) Cu(s) Balance the
following redox reaction and calculate the mass
of O2(g) and Cu(s) produced in 30 sec by a
current of 1.50 A.
Ex 3 Zn(s) Cu2 ? Zn2 Cu(s) If a current
of 1.00 A is used, how many minutes does it take
to electrolyze 0.0146 g of copper?
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Corrosion
Deteriation of metals. Rust forms when Fe(s) is
oxidized to Fe2O3 and O2(g) is reduced to
H2O. 4Fe(s) 3O2(g) ? 2Fe2O3
(brittle)   Ecell occurs
spontaneously Ag(s) ? Ag2S Cu(s) ? CuCO3(patina)
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Physical Properties of Solutions
Dissolution of a solute in a solvent. saturated
A solution that contains the maximum amount of
solute per given volume of solvent. unsaturated
Given volume of solvent can hold more
solute. supersaturated Given volume of solvent
is holding more solute than a saturated
solution. Like dissolves like.
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Concentration Units
Molarity(M)
Molality(m)
Ex Calculate the molality and molarity of a
solution prepared by dissolving 0.20 moles of
ethanol in 500 g of water.
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Concentrations cont
Ex2 Calculate the molality of a solution
prepared by dissolving 5.05 g of C10H8 in 75.0 mL
of benzene (d 0.879 g/mL).
Percent by Mass(w/w) Percent by Volume(w/v)
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Concentration cont
Ex A solution is prepared by dissolving 15.0 g
of pure ethanol , C2H5OH, (d 0.800 g/mL) in
20.0 g of water(d 1.00 g/mL). Assuming the
volumes are additive, calculate the concentration
in w/w and w/v.
Parts per Million(ppm)
Ex If 0.090 g or 90.0 mg of solute is added to
2.0 L of solution, calculate the concentration in
ppm.
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Concentration cont
Parts per Billion(ppb)     Mole Fraction
 
Enthalpy of Solution(?Hsol)   ?Hsol ?H1 ?H2
?H3   ?Hsol lt 0 exothermic ?Hsol gt 0 endothermic
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Henry's Law
S k?PGAS S solubility of the dissolved gas
in the solvent. PGAS partial pressure of gas
above solution. k Henry's constant.
Ex The Henry's law constant of methyl bromide is
k 0.159 mol?L-1?atm-1 at 25 ?C. What is the
solubility, in mol/L, of methyl bromide at 25 ?C
and a partial pressure of 125 mm Hg?
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Colligative Properties
Colligative properties are physical properties of
solutions that are dependant on the number of
dissolved solute particles. Independent of the
type of particle.
Raoult's Law Presence of a solute lowers the
vapor pressure of a solvent. Psolv
Xsolv?Psolv?   Psolv vapor pressure of solvent
above a solution. Xsolv mole fraction of
solvent. Psolv? vapor pressure of pure solvent.
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Colligative properties cont
Ex The vapor pressure of pure water is 17.5 mm
Hg at 20.0? C. Calculate the vapor pressure of
water above a 1.00 m urea solution at 20.0? C.
Ex2 Vapor pressure of benzene(C6H6) at 25? C is
95.1 mm Hg. What is the vapor pressure of benzene
at the same temperature if 5.05 g of benzoic
acid, C6H5COOH, is added to 245 g of benzene?
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Freezing Point Depression and Boiling Point
Elevation
?TB TB(solution) - TB(solvent) ?Tf
Tf(solution) - Tf(solvent)   ?TB KB?m ?Tf
-Kf?m   KB, Kf are constants for a particular
solvent. m molality of solute.
Ex Pure water freezes at 0.00?C. Calculate the
freezing point of a solution consisting of 25 g
of C12H22O11 and 100.0 g of water. Kf 1.86?C/m
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Freezing Point Depression and Boiling Point
Elevation cont
Can use freezing point depression to determine
molecular weight of a solute.
Ex An aqueous solution containing 1.00 g of
sorbitol in 100.0 g of water has a freezing point
of 0.102 ?C. Calculate the molar mass of
sorbitol.
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Osmotic Pressure(?)

• The pressure required to stop osmosis.
• MRT
• ? osmotic pressure in atm.
• M molarity
  
• R 0.0821 L?atm/K?mole
  
• T temperature in Kelvin

Ex A 202 mL solution containing 2.47 g of a
polymer has an osmotic pressure of 8.63 mm Hg at
21.0?C. Calculate the molar mass of the polymer.
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Tonic Solutions
Isotonic Two solutions seperated by a
semipermeable membrane with the same
concentration and thus the same osmotic pressure.
Hypertonic Solution with a higher concentration
and thus higher osmotic pressure.
Hypotonic Solution with a lower concentration
and thus lower osmotic pressure.