Unit 9: Thermodynamics and Electrochemistry

9.1 Introduction to Entropy 9.2 Absolute Entropy and Entropy Change 9.3 Gibbs Free Energy and Thermodynamic Favorability 9.4 Thermodynamic and Kinetic Control 9.5 Free Energy and Equilibrium 9.6 Free Energy of Dissolution 9.7 Coupled Reactions 9.8 Galvanic (Voltaic) and Electrolytic Cells 9.9 Cell Potential and Free Energy 9.10 Cell Potential Under Nonstandard Conditions 9.11 Electrolysis and Faraday’s Law

entropy (ΔS)

amount of disorder in a system

state of mater with most entropy

gas

what effect does increasing moles have on entropy

increases entropy

what happens to entropy when temp increases

increases

how to find overall entropy

$$\Delta S_{rxn}=\Sigma S_{products}-\Sigma S_{reactamts}$$

units of entropy

J / mol * K

gibbs free energy (ΔG)

determines how spontaneous or favorable a reaction is

units of gibbs free energy

kJ / mol

spontaneous reaction

reaction can occur independently without external input

what does ΔG < 0 mean

reaction is spontaneous/favorable

what does ΔG > 0 mean

reaction is unfavorable

enthalpy (ΔH)

heat content of the system

what does ΔH < 0 mean

reaction is exothermic; releases energy to its surroundings

what does ΔH > 0 mean

reaction is endothermic; absorbs energy from its surroundings

gibbs equation

$$\Delta G=\Delta H-T\Delta S$$

most optimal conditions for favorable reaction

-ΔH and +ΔS

if +ΔH and +ΔS, reaction favored at

high temps

if +ΔH and -ΔS, reaction favored at

no temp

if -ΔH and -ΔS, reaction favored at

low temps

limitation of thermodynamics

cannot predict rate of reaction: need kinetics to determine rate

why might a reaction be thermodynamically favorable but have little product formation

favorable reactions don’t necessarily have to be instantaneous; requires sufficient energy to overcome activation energy

relationship between ΔG and K

$$\Delta G=-RT\ln K$$

if ΔG < 0, what does that mean for K

products favored, K > 1

if ΔG > 0, what does that mean for K

reactants favored, K < 1

thermodynamics of dissolution

most dissolutions endothermic, so entropy must be the driving force in order for dissolution reactions to be favorable

coupled reaction

when an unfavorable reaction is powered by a favorable reaction

oxidation

loss of electrons

reduction

gain of electrons

how to balance ion-electron ½ equations (involves polyatomic ions)

when balancing elements, add H2O to balance out the Os (remember to add H+ to the other side after)

chemical cell

reactants separated in two half cells where electrons can move through a wire

salt bridge

allows ions to freely move between half cells; maintains the charge of the system and prevents ion buildup in one beaker

where does oxidation occur

at the anode

where does reduction occur

at the cathode

how to identify cathode/anode *IF RXN NOT GIVEN*

voltage of cathode is more positive than voltage of anode

galvanic/voltaic cells

cells where the reaction is spontaneous and ΔG < 0

cell potential (E_cell)

difference in voltage between two electrodes

units of E_cell

volts (V)

when E_cell > 0

reaction is spontaneous; ΔG < 0

how to calculate E_cell

$$E_{cell}=E_{ca\operatorname{th}ode}-E_{anode}$$

effect of moles on E_cell

nothing, cell potential is an intrinsic property

electrolytic cells

cells where the reaction is spontaneous and ΔG < 0, requires a battery to power

relationship between ΔG and E_cell

$$\Delta G=-nFE$$

relationship between moles of electrons and Faradays

1 mol e^- = 1 Faraday = 96,485 coulombs

nernst equation

$E=E_{cell}-\frac{RT}{nF}\ln Q$ ; describes voltage of cell when not under standard conditions (in equilibrium)

if Q < 1 (in terms of E)

E > E_standard

if Q > 1 (in terms of E)

E < E_standard

if Q = 1 (in terms of E)

E= E_standard and cell in standard conditions

concentration cell

cell with similar ions; ions flow from lower concentration to higher concentration

equation relating moles of e- to electrical current

$q=I\cdot t$ —> $n\cdot F=I\cdot t$