MCAT General Chemistry - Electrochemistry

696

Electrochemical cells

contained systems in which Oxidation–Reduction reactions occur

electrodes

locations in an electrochemical cell where oxidation and reduction take place

anode

electrode where oxidation occurs; electrons origin (negative, negative reduction potential) in galvanic cell and anion attractor (positive, positive reduction potential) in electrolytic cell

cathode

electrode where reduction occurs; electrons attractor (positive, positive reduction potential) in galvanic cell and cation attractor (negative. negative reduction potential) in electrolytic cell

electromotive force (emf)

corresponds to the voltage or electrical potential difference of the cell

positive electromotive force

cell is able to release energy; (ΔG < 0); spontaneous

negative electromotive force

cell must absorb energy; (ΔG > 0); nonspontaneous

current in electrochemical cells

inverse of movement of electrons from anode to cathode; direction of flow of a positive charge

galvanic / voltaic cells

nonrechargeable spontaneous batteries; two electrodes of distinct chemical identity are placed in separate compartments connected to each other by a conductive material

ΔG < 0

E_cell > 0

half-cells

one of the two distinct electrodes in a galvanic cell

electrolyte

aqueous ion solution composed of cations and anions

Daniell cell

galvanic cell where zinc is the anode and copper is the cathode; each electrode is bathed in an electrolyte solution containing its cation and sulfate

salt bridge

inert salt that connects the two solutions in a galvanic cell; dissipates charge gradient by permitting the exchange of cations and anions; contains ions that will not react with the electrodes or with the ions in solution

plating / galvanization

precipitation process onto the cathode

cell diagram

shorthand notation representing the reactions in an electrochemical cell

anode | anode solution (concentration) || cathode solution (concentration) | cathode

single vertical line ( | ) indicates a phase boundary.

double vertical line ( || ) indicates the presence of a salt bridge or

some other type of barrier.

ex. Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)

electrolytic cells

house nonspontaneous reactions that require the input of energy to proceed

molten NaCl is decomposed into Cl₂(g) and Na(l)

ΔG > 0

E_cell < 0

electrolysis

Oxidation–Reduction reaction driven by an external voltage source in which chemical compounds are decomposed

Faraday equation of Oxidation–Reduction

the amount of chemical change induced in an electrolytic cell is directly proportional to the number of moles of electrons that are exchanged

Mⁿ⁺ + n e^– → M (s)

Faraday constant, (F)

equivalent to the amount of charge contained in one mole of electrons

1 F = 96,485 C

electrodeposition equation

helps determine the number of moles of element being deposited on a plate; also used to determine the amount of gas liberated during electrolysis

mol M = It/nF

where mol M is the amount of metal ion being deposited at a specific electrode, I is current, t is time, n is the number of electron equivalents for a specific metal ion, and F is the Faraday constant

concentration cell

special type of galvanic cell where the electrodes are chemically identical; current is generated as a function of a concentration gradient established between the two solutions surrounding the electrodes; drives the movement of electrons in the direction that results in equilibration of the ion when the current will stop

emf → 0

Nernst equation

calculates the voltage as a function of concentrations in a concentration cell

E_cell = E^∘_cell − (RT/nF)lnQ = E^∘_cell - 0.0592/n log Q

resting membrane potential (Vm)

Sodium and potassium cations and chlorine anions are exchanged as needed to produce an electrical potential across a cellular membrane; Disturbances may stimulate the firing of an action potential

rechargeable cell / battery

function as both a galvanic and electrolytic cell

lead–acid battery / lead storage battery

rechargeable battery; low energy density

fully charged = Pb anode and porous PbO2 cathode, connected by a conductive material (concentrated 4 M H2SO4)

discharging = electrodes plate with lead sulfate (PbSO4) and dilute the acid electrolyte

fully discharged = two PbSO4 electroplated lead electrodes with a dilute concentration of H2SO4

charging = reverses the electroplating process and concentrates the acid solution

discharge

lose electric potential in a battery by changing the composition of electrodes; may be irreversible or reversible; galvanic

charging

foster electric potential in a battery by restoring the electrodes; electrolytic

Energy density

measure of a battery’s ability to produce power as a function of its weight

Nickel–cadmium batteries

rechargeable cells; two half-cells made of solid cadmium (the anode) and nickel(III) oxide-hydroxide (the cathode) connected by a conductive material, typically potassium hydroxide (KOH); higher energy density

ex. AA and AAA cells

Surge currents

periods of large current (amperage) early in the discharge cycle; preferable in appliances such as remote controls that demand rapid responses

nickel–metal hydride (NiMH) batteries

more efficient, more energy density, more cost effective, less toxic modern alternative to Ni-Cd batteries

Isoelectric focusing

technique used to separate amino acids or polypeptides based on their isoelectric points (pI); positively charged / protonated amino acids migrate toward the cathode; negatively charged / deprotonted amino acids migrate toward the anode

standard hydrogen electrode (SHE),

relative standard of reduction potential; 0 V by convention

reduction potential

the tendency of a species to gain electrons and to be reduced; more positive the potential, the greater the tendency to be reduced

Standard reduction potential (E°red)

measured under standard conditions: 25°C (298 K), 1 atm pressure, and 1 M concentrations; predict the direction of electron flow

oxidation potential

reduction half-reaction and the sign of the reduction potential are reversed

standard electromotive force (emf or E°cell)

difference in potential (voltage) between two half-cells under standard conditions; do NOT multiply them by the number of moles oxidized or reduced

E°cell = E°red,cathode − E°red,anode

Gibbs Free Energy in an electrochemical cell

change in the amount of energy available in a chemical system to do work

ΔG° = –nFE°cell

where ΔG° is the standard change in free energy, n is the number of moles of electrons exchanged, F is the Faraday constant, and E°cell is the standard emf of the cell

ΔG° and E°cell will always have opposite signs

Gibbs Free Energy in equilibrium

ΔG° = –RT ln Keq

where R is the ideal gas constant, T is the absolute temperature, and Keq is the equilibrium constant for the reaction

equilibrium constants less than 1 (favors the reactants), the E°cell will be negative

equilibrium constant greater than 1 (favors the products), the E°cell will be positive

Gibbs Free Energy in nonstandard conditions

ΔG = ΔG° + RT ln Q

where ΔG is the free energy change under nonstandard conditions, ΔG° is the free energy change under standard conditions, R is the ideal gas constant, T is the temperature, and Q is the reaction quotient