AP Chem Unit 9- Electrochemistry

oxidation

loss of electrons

reduction

gain of electrons

oxidation states are written as:

+n or -n

generating electricity is a:

spontaneous redox reduction (K<1)

using electricity is a:

nonspontaneous redox reduction (K>1)

oxidation state- group 1 metals

always +1

oxidation state- group 2 metals

always +2

oxidation state- fluorine

always -1

oxidation state- oxygen

usually -2

oxidation state exception- oxygen

-1 in peroxides, +2 in F₂O

oxidation state- chlorine

usually -1

oxidation state exception- chlorine

not -1 in compound with O & F

oxidation state- hydrogen

usually +1

oxidation state exception- hydrogen

-1 in metal hydrides (group 1)

sum of oxidation states for a neutral compound

0

sum of oxidation numbers for a monatomic or polyatomic ion

ion charge

oxidation number- neutral elements alone

0

oxidation state- ionic compounds

charges

steps for balancing half-reactions

1. separate the half reaction

2. balance the atoms

3. balance the e^- in the equation

   1. e^- must cancel out

4. add the equations together (hess’ law)

balanced for mass & charge

galvantic/voltaic cells

• cells that generate an electrical current (is a battery)

• spontaneous

• moves towards equilibrium

• deltaG= -

• energy is released as heat

• must be separated half-cells

  • oxidizing/reducing agents are separated by wire & salt bridge

    • electrons flow from one container to another

electrolytic cells

• cells driven by externally applied electrical potential (needs a battery)

• nonspontaneous

• pushed from equilibrium

• deltaG= +

• can be all in one cell (electrons forced through wire)

  • which metal is anode or cathode is switched in comparison to galvantic cells

anode

• oxidizes, electrons lost

• metal atom at surface of electrode loses “n” electrons to become an Mⁿ⁺ ion in solution

• lose mass (solid dissolves as atoms become ions)

cathode

• reduces, electrons gained

• metal ion Mⁿ⁺ from solution collides with metal electrode, gaining “n” electrons from it, converting to metal atoms

• gain mass (dissolved ions plate out on solid cathode)

salt bridge

• connects two half-cells to complete the circuit and prevent buildup of charge in half-cells

• keeps solutions in each half-cell electrically neutral

electron flow direction

anode to cathode (e^- lost to e^- gain)

anion (negative ions) flow direction

cathode to anode (e^- gain to e^- lost)

line notation

• solid l aqueous_(M) ll aqueous_(M) l solid

  • anode (oxidation) on left, cathode (reduction) on right

standard reduction potentials (E⁰)

• measure of the potential of a half-reaction to reduce at standard conditions (1M, 1atm)

  • measured in volts (V)

• calculated by comparing them to reduction of hydrogen

  • more easily= +, more difficult= -

elements with the greatest (most positive) reduction potentials are the most easily reduced, becoming the:

cathode

elements with the smallest (most negative) reduction potentials are the most easily oxidized, becoming the:

anode

steps for calculating standard cell pontential

1. identify reactions occurring at anode & cathode

2. flip sign for anode reaction to turn reduction potential into oxidation potential

3. add both potentials (E°_red+E°_ox=E°_cell)

true/false: coefficients from balancing a half-reaction do not change the reduction potential

true

for a cell to be spontaneous, the overall E°_cell must be:

positive

nonstandard cell potential

cell potential represents driving force towards equilibrium

• the further away from equilibrium, the greater E_cell gets

• the closer to equilibrium, the smaller E_cell gets

• reaching equilibrium causes there to no longer be a potential difference, E_cell= 0 (dead battery)

cell potential & concentration

Q for standard cells is always 1 regardless of coefficients/exponents

• Q= anode/cathode

• changes causing decrease in Q bring cell further from equilibrium, making E_cell larger

• changes causing increase in Q bring cell closer to equilibrium, making E_cell smaller

nerst equation

E_cell=E°_cell - (RT/nF) ln (Q_cell)

• E°= standard cell potential

• R= gas constant; 8.314 J mol^-1 K^-1

• T= temp; K

• n= number of moles of e^- transferred

• F= faraday’s constant; 96485 C/mol

• Q= reaction quotient

electrolysis

• similar to running a galvantic cell bacwards

• use voltage bigger than cell potential to reverse direction of redox reaction

  • used for separating metals from ores or plating out metals

• can recover metals from pure molten ionic compounds by passing electric current through it

  • water may be oxidized/reduced instead if present (aqueous solutions)

similarities of voltaic & electrolytic cells

• anode=oxidation

• cathode=reduction

• electrons flow anode to cathode

• anions (-) flow towards anode

in a pure molten ionic compound during electrolysis:

• cation will be reduced (gain e^-, decrease in charge)

• anion will be oxidized (lose e^-_, increase in charge)

in an aqueous solution of an ionic compound or when water is present during electrolysis:

• determine if ions or water is reacting

  • cation vs. water: highest reduction potential generally reduces

    • no group 1 or 2 metal will be reduced; H₂O reduces

  • anion/metal electrode vs. water: lowest reduction potential generally oxidizes

    • no polyatomic ion will be oxidized; H₂O oxidizes

oxidation half-reaction for the electrolysis of H₂O

2H₂O = O₂ + 4H⁺ +4e^- E°_cell= -1.23V

reduction half-reaction for the electrolysis of H₂O

2H₂O + 2e^- = H₂ +2OH^- E°_cell= -0.83V

electroplating

• application by electrolysis of a thin ornamental or protective coating of one metal over another

• cathode

  • item being plated (electrode)

  • solution=desired metal

• anode

  • sacrificial electrode

  • electrode oxidizes, replenishing ions in solution

electric current equation

I=q/t

• I= current (amps)

• q= charge (coulombs)

• t= time (seconds)

gibbs free energy

deltaG=-nFE°

• deltaG= free energy; J/mol

• n= number of e^- transferred

• F= faraday’s constant; 96485 C/mol

• E°= cell potential under standard conditions