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Unit 9 Application of thermodynamics - Coggle Diagram
Unit 9 Application of
thermodynamics
9.1 Entropy
Entropy increases when
matter becomes more dispersed.
phase change from solid to liquid or
from liquid to gas: S(g) > S(l) > S(s) :red_flag:
gas volume increases
(at constant temperature)
number of moles of gas-phase products is greater than the total number of moles of gas-phase reactants. :red_flag:
energy is
dispersed
the entropy of the system increases with an increase in temperature
because based on KMT: the distribution of kinetic energy among the particles of a gas
broadens
as the temperature increases
9.2 Absolute entropy
and entropy change
ΔS° reaction = Σ n S° products – Σ m S° reactants
Every substance has a nonzero
value for absolute entropy
Entropy is often measured in
J/K
:warning:
9.3 Gibbs Free Energy and
Thermodynamic Favorability
Temperature dependence
of ΔG° :red_flag:
ΔG°=ΔH°-TΔS°
ΔH°>0, ΔS°>0: favored at high temperature
when T >(∆H°)/(∆S° ). Entropy driven reaction
ΔH°<0, ΔS°<0: favored at low temperature
when T <(∆H°)/(∆S° ). Enthalpy driven reaction
ΔH°>0, ΔS°<0: both enthalpy and entropy are unfavorable
ΔH°<0, ΔS°>0: both enthalpy and entropy
are favored
standard Gibbs
free energy change
standard Gibbs free energy change: ΔG°
standard state: pure substances, 1.0 M of solutions, 1.0 atm of gases
ΔG° reaction = Σ nΔGf°products – Σ mΔGf°reactants
Thermodynamic
Favorability
ΔG° > 0, nonspontaneous /
thermodynamically unfavored.
ΔG°< 0, spontaneous /
thermodynamically favored.
ΔG° = 0, system at equilibrium
9.4 Thermodynamic
and Kinetic Control
Kinetic control
processes that are thermodynamically favored,
do not occur to any measurable extent, or they occur at extremely slow rates.
High activation energy
is a common reason for a process to be under kinetic control.
drive thermodynamically
unfavorable processes (ΔG >0).
Apply external sources of energy
9.6 coupled reactions
9.7 Galvanic / Voltaic
and Electrolytic cells
Galvanic cell
Operational
characteristics
Anode: oxidation, Cathode: reduction
Anode: lose mass, cathode: gain mass or gas evolution
electrons: from anode to cathode
Anions to anode, Cations to cathode
Component
thermodynamically favored
redox reaction, ΔG<0, E>0
electrodes: anode (-), cathode (+)
solutions in the half-cells
salt bridge
form closed circuit
maintain charge balance
voltage/current measuring device
electrolytic cell
Component
thermodynamically unfavored
redox reaction, ΔG>0, E<0
Voltage source
electrodes: anode (+), cathode (-)
electrolyte solutions
Operational
characteristics
Anode: lose mass
cathode: gain mass or gas evolution
electrons from negative terminal to cathode,
electrons from anode to positive terminal.
Anode: oxidation, Cathode: reduction
Anions to anode, Cations to cathode
9.10 Electrolysis and
Faraday’s Law
Number of electrons transferred
Current
Mass of material deposited on
or removed from an electrode
Time elapsed
Charge of ionic species
9.8 Cell potential
and free energy
voltaic cell: Ecell° = E°cathode -E°anode
(E° always positive)
all redox reaction:
E° = E°red (reduction) – E°red (oxidation)
E>0, spontaneous redox
E<0, nonspontaneous redox
ΔG° = −nFE°,
n
is the number of electrons in the balanced half-reactions for a process :star:
9.5 Free Energy
and Equilibrium
ΔG°<0, ⇒ K>1; products are favored over reactants
ΔG°>0, ⇒ K<1; reactants are favored over products
ΔG° = −RT ln K
ΔG°=0, ⇒ K=1; neither reactants nor products are favored.
9.9 Cell potential under
nonstandard conditions
Smaller Q,
larger E
When Q is
increased above 1
, the system is
moved closer
to equilibrium and the
cell potential decreases
relative to standard conditions.
When Q is
decreased below 1
, the cell is
further from equilibrium
and the
cell potential increases
relative to standard conditions.
The cell potential is the driving force toward equilibrium
Concentration cell
the direction of spontaneous electron flow can be determined by considering the
direction needed to reach equilibrium
Anode: oxidation, increase M+, dilute solute
Cathode: reduction, decrease M+, concentrated solution
E = E° − (RT/nF) ln Q
E= E°-(0.0592/n) lgQ (at 298 K)
7.14 Free energy
of dissolution
breaking solute-solute interaction,
ΔH>0, ΔS>0
breaking solvent-solvent interaction,
ΔH>0, ΔS>0
△G°=△H°-T△S°
G°<0, thermodynamically favorable ⇒soluble salt
G°>0, thermodynamically unfavorable ⇒insoluble /slightly soluble salt
forming solute-solvent interaction
ΔH<0, ΔS<0