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29 Cards in this Set
- Front
- Back
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First Law of Thermodynamics |
Energy is conserved. Not created/destroyed. Transformed |
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Second Law of Thermodynamics |
Some energy is always wasted (Friction, Heat loss, etc.). Becomes Entropy (Disorder) |
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Spontaneous Process |
Occurs w/o Outside intervention; Exothermic |
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Entropy |
S = k ln(W) k = 1.38 E -23 J/K |
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Entropy Equations (dSuniverse) |
dSuniverse = dSsystem + dSsurroundings
dSsurroundings = -dHsystem/T
dSuniverse = dSsystem - dHsystem/T
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Spontaneity with dSuniverse |
if dSuniverse > 0 Spontaneous
if dSuniverse < 0 and if -dSsystem < dSsurroundings Spontaneous |
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Entropy Equations (dG) |
dG = dHsystem - TdSuniverse |
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Gibb's Free Energy (dG) |
Free Energy Change;
if dGrxn decreases; Process = Spontaneous;
if dGrxn increases; Process != Spontaneous;
Exothermic Reactions = Usually Spontaneous |
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Entropy (dS) and Enthalpy (dH) |
+dS and -dH = Always Spontaneous -dS and +dH = Never Spontaneous +dS and +dH = Spontaneous @ High Temp. -dS and -dH = Spontaneous @ Low Temp. |
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Gibb's Free Energy Equations() |
dG = dG* + RT ln (Q)
Q = Reaction Quotient = Products/Reactants
dG = dGproducts - dGreactants |
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dG vs. dG* vs. dGf |
dG - Free energy for reaction under any condition
dG* - Free energy change for reaction in standard state
dGf - Formation free energy change for a compound from its elements |
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Keq with dH and dS |
ln (k) = (-dH/R)(1/T) + (dS/R)
ln (k2/k1) = -(dH/R)(1/T2 - 1/T1) |
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Redox Reactions |
Oxidation Number of atoms changes |
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Oxidation |
If oxidation number gets bigger |
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Reduction |
If oxidation number gets smaller |
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Oxidation Number Rules (6 Rules) |
1) Pure elements have a number of 0 2) Monatomic Ions have # = Charge 3) Hydrogen = +1 unless w/ metal (then -1) 4) Fluorine is always -1 5) Oxygen = -2 unless attached to Fluorine or itself 6) Halogens = -1 unless attached to F, O, or each other |
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Balancing in Acidic Solution |
1) Split into half reactions (Oxidation means e- are products; Reduction means e- are reactants) 2) Remove Electons 3) Balance elements other than H and O 4) Use H2O to balance O's 5) Use H+ to balance H's 6) Balance Charges 7) Multiply by correct value to ensure e- cancel |
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Balancing in Basic Solution |
1) Split into half reactions 2) Remove Electons 3) Balance elements other than H and O 4) Use H2O to balance O's 5) Use H+ to balance H's 6) Add OH- to cancel out H+; Cancel H2O's 7) Balance Charges 8) Multiply by correct value to ensure e- cancel |
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Galvanic/Voltaic Cells |
Reactions in separate compartments; Electrons forced to move; Creates useful work |
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Voltaic Reactions |
Atoms in one electrode release electrons
Electrons travel from solution through wire and reduce ion on other side.
Salt bridge allows travel of ions |
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Voltaic Cell Rules |
1) Anode = Oxidation = Product electrons 2) Cathode = Reduction = Reactant electrons 3) Electrons flow from anode to cathode 4) Electrons cannot swim |
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Cell Notation |
- Single line = Change of phase - Double line = Salt Bridge - Anode||Cathode - Oxidation||Reduction |
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Amps and Volts |
Amperes (A) = 1 Coulomb (C)/second (s)
Volt (v) = 1 Joule (J)/Coulomb (C)
Spontaneous Reaction = Positive Voltage
More Spontaneous Reaction = Greater Voltage |
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Cell EMF |
- Spontaneous Reaction: Electrons flow from Anode to Cathode
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Standard Reduction Potential |
- Standard Reduction Potential (E*red) is the potential for a half-reaction written as a reduction - E*cell = E*red(Cathode) - E*red(Anode) - if E*red > 0 then reaction is spontaneous - More positive E*red = Stronger Oxidizing Agent - More negative E*red = Stronger Reducing Agent |
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E* and dG* and K |
Spontaneous Rxn: E*>0; dG*<0; K>1 Non-Spontaneous Rxn: E*<0; dG*>0; K<1
Cell Reaction w/ n elections: dG* = -nFE* |
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Calculating E*cell |
E*cell = (RT/nF) ln (K) |
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Nernst Equation |
Ecell = E*cell - (RT/nF) ln (Q) |
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Concentration Cells |
Same element with different concentrations |
