ch 12 - Electrochemistry

Question 1

electrochemical cells

Answer 1

contained systems in which oxidation-reduction reactions occur; three types are galvanic (voltaic), electrolytic, and concentration

Question 2

electrodes

Answer 2

area in electrochemical cells where oxidation and reduction take place

Question 3

anode

Answer 3

electrode at which oxidation takes place; always attracts anions (no matter type of cell)

Question 4

cathode

Answer 4

electrode at which reduction takes place; always attracts cations (no matter type of cell)

Question 5

electromotive force (emf)

Answer 5

corresponds to voltage or electrical potential difference of the cell; if positive, the cell is able to release energy (delta G <0), which means it is spontaneous; if emf is negative, cell must absorb energy (delta G > 0) and it is nonspontaneous; current (I) runs from cathode to anode and movement of electrons is from anode to cathode

Question 6

galvanic (voltaic) cells

Answer 6

batteries that are nonrechargeable; reactions must be spontaneous (G <0) and electromotive force (E sub cell) must be positive; electromotive force and free energy change always have opposite signs

Question 7

how galvanic cells work

Answer 7

two electrodes placed in separate compartments call half-cells; two electrodes surrounded by aqueous electrolyte solution composed of cations and anions are connected to each other by a conductive material; when electrodes are connected to each other by conductive material charge flows bc of spontaneous oxidation-reduction reaction

Question 8

salt bridge

Answer 8

a structure made of an inert salt (usually KCl or NH4NO3) that connects the two half-cell solutions in a galvanic cell; permits the exchange of cations and anions so excess negative charge does not build up on the cathode and vice versa for the anode causing flow of electrons to stop

Question 9

Daniell cell

Answer 9

type of galvanic cell that has the cations in the two half-cell solutions the same element as the respective metal electrode

Question 10

plating or galvanizing

Answer 10

the precipitation process onto cathode of anions from the salt bridge diffusing into solution on anode side to balance out charge of newly created zinc ions and cations of the salt bridge flow into solution on cathode side to balance out charge of sulfate ions left in solution when copper ions are reduced to copper and precipitate onto the electrode

Question 11

cell diagram

Answer 11

shorthand notation representing reactions in an electrochemical cell: for Daniell cell: Zn (s) | Zn(2+) (1 M) || Cu (2+) (1 M) | Cu (s); follows rules: reactants and products are always listed from left to right as anode | anode sol’n (concentration) || cathode sol’n (concentration) | cathode; 2. single vertical line indicates a phase boundary; 3. double vertical line indicates presence of a salt bridge or some other type of barrier

Question 12

similarities in all electrochemical cells

Answer 12

have reduction reaction occurring at the cathode, oxidation reaction occurring at the anode, current flowing from cathode to anode, and electron flow from anode to cathode

Question 13

differences between electrolytic cells and galvanic cells

Answer 13

electrolytic cells house nonspontanesous reactions that require input of energy to proceed. change in free energy is positive; half-reactions are not separated into different compartments; key: in galvanic, anode is negative and cathode is positive; in electrolytic, anode is positive and cathode is negative (but reduction is at cathode and oxidation is at anode for both)

Question 14

electrolysis

Answer 14

oxidation-reduction reaction driven by an external voltage source (as in electrolytic cells) in which chemical compounds are decomposed

Question 15

equation for number of moles of electrons exchanged in electrolytic cells

Answer 15

can be determined from the balanced half-reaction; for reaction that involves transfer of ‘n’ electrons per atom M: M^(n+) + n (e-) -> M (s); one mole of metal (M (s)) will be produced if ‘n’ number of moles of electrons are suppolied to one mole of M (n+)

Question 16

charge of one electron

Answer 16

1.6 x 10^-19 coulombs (C)

Question 17

Faraday constant

Answer 17

one faraday (F) is equivalent to the amount of charge contained in one mole of electrons (1 F = 96, 485 C) or one equivalent. Number should be rounded to 10^5 C/mol e-; number is derived by multiplying charge of one electron by Avogadro’s number (6.02 x 10^23) for one mole of electrons

Question 18

electrodeposition equation

Answer 18

mol M = (It)/(nF); where mol M = amount of metal ion being deposited at a specific electrode; I = current; t = time in seconds, n = number of electron equivalents for a specific metal ion (ex - oxidation state of metal in solution), F = Faraday constant; can also be used to determine that amount of gas liberated during electrolysis

Question 19

concentration cells

Answer 19

special types of galvanic cells; like galvanic cells, contains two half-cells connected by conductive material, allowing spontaneous redox reaction which generates current and delivers energy; distinguished by design: electrodes are chemically identical and have same reduction potential - current generated by concentration gradient bt two solutions surrounding electrodes resulting in potential difference and driving the electrons in direction to get equilibrium of ion gradient (current stops when ionic species in half-cells are equal); voltage (V) = 0

Question 20

resting membrane potential (V sub m)

Answer 20

sodium and potassium cations, and chlorine anions are exchanged as needed to produce electrical potential in concentration cells; disturbances may stimulate firing of an action potential

Question 21

rechargeable cell (rechargeable battery)

Answer 21

cell that can function as both galvanic and electrolytic

Question 22

lead-acid battery

Answer 22

also called lead storage battery: type of rechargeable battery; as galvanic cell when fully charged, consists of two half-cells (Pb anode and porous PbO2 cathode) connected by conductive material (concentrated 4 M H2SO4); when fully discharged, consists of two PbSO4 electroplated lead electrodes with dilute concentration of H2SO4

Question 23

net equation for a discharging lead-acid battery

Answer 23

Pb (s) + PbO2 (s) + 2H2SO4 (aq) -> 2 PbSO4 (s) + 2H2O; E, degree sign, sub cell = 1.685 - (-0.356) = 2.041 V

Question 24

charging lead-acid cell

Answer 24

part of an electrolytic circuit; equation for this is opposite net equation for discharging battery, as is the electrode charge designation

Question 25

energy density

Answer 25

energy-to-weight ratio; measure of battery’s ability to produce power as a function of its weight (ex. lead-acid batteries need a heavier amount of material to produce a certain output as compared to other batteries

Question 26

nickel-cadmium batteries

Answer 26

rechargeable cells consisting of two half-cells made of solid cadmium (the anode) and nickel (III) oxide-hydroxide (the cathode) connected to a conductor (typically potassium hydroxide (KOH)); charging reverses electrolytic cell potentials; have higher energy density than lead-acid batteries

Question 27

surge currents

Answer 27

periods of large current (amperage) early in the discharge cycle.

Question 28

nickel-metal hydride (NiMH) batteries

Answer 28

have largely replaced nickel-cadmium (Ni-Cd) batteries as they have more energy density, are more cost effective and are less toxic; in lieu of pure metal anode, a metal hydride is used instead

Question 29

isoelectric focusing

Answer 29

a technique used to separate amino acids or polypeptides based on their isoelectric point (pI): positively charged amino acids (protonated at solution’s pH) will migrate toward cathode; negatively charged amino acids (deprotonated at solution’s pH) will migrate to the anode

Question 30

characteristics of Ni-Cd discharging battery

Answer 30

galvanic; anode material is Cd, anode charge is negative; cathode material is NiO(OH) and cathode charge is positive

Question 31

characteristics of Ni-Cd charging battery

Answer 31

electrolytic; Cd(OH)2 is anode material which is positively charged; Ni(OH)2 is cathode material which is negatively charged

Question 32

characteristics of Molten NaCl discharging battery

Answer 32

electrolytic; any anode material which is positively charged; any cathode material which is negatively charged

Question 33

characteristics of Daniell cell discharging battery

Answer 33

galvanic; zinc is the anode material and is negatively charged; copper is the cathode material and is positively charged

Question 34

characteristics of lead-acid charging battery

Answer 34

electrolytic with positively charged PbSO4 anode; negatively charged PbSO4 cathode

Question 35

characteristics of lead-acid discharging battery

Answer 35

galvanic with negatively charged Pb anode and positively charged PbO2 cathode

Question 36

cell diagram for discharging state of lead-acid battery

Answer 36

Pb (s) | H2SO4 (4 M) || H2SO4 (4 M) | PbO2 (s)

Question 37

standard hydrogen electrode (SHE)

Answer 37

reference for defining a reduction potential which is measured in volts; this is given potential of 0 V by convention

Question 38

reduction potential

Answer 38

tendency of a species to gain electrons and be reduced; allows determination of which species will be oxidized or reduced; each species has its own reduction potential; the more positive the potential, the greater the tendency to be reduced

Question 39

standard reduction potential (E, degree sign, sub red)

Answer 39

reduction potential measured under standard conditions (25 degrees celsius, 1 atm, 1 M concentrations); positive means greater relative tendency for reduction to occur (cathode in galvanic cells); less positive means greater relative tendency for oxidation to occur (anodes in galvanic cells) which allows to predict flow of electrons

Question 40

standard electromotive force (emf or E, degree sign, sub cell)

Answer 40

the difference in potential (voltage) between two half-cells under standard conditions calculated from standard reduction potentials; determined by calculating difference in reduction potentials bt the two half-cells: E, degree sign, sub cell = E, degree sign, sub (red, cathode) - E, degree sign, sub (red, anode)

Question 41

relationship between gibbs free energy and emf

Answer 41

delta G, degree sign = -nFE, degree sign, sub cell; delta G, degree sign = standard change in free energy; n = number of moles of electrons exchanged; F = Faraday constant; E, degree sign, sub cell = standard emf of the cell; if Faraday constant is expressed in coulombs (J/V) then standard change in free energy must be expressed in J, not kJ

Question 42

Nernst equation

Answer 42

calculates voltage as a function of concentrations; E sub cell = E, degree sign, sub cell - (RT)/(nF) x lnQ; E sub cell = emf of the cell under nonstandard conditions; E, degree sign, sub cell = emf of the cell under standard conditions; R = ideal gas constant; T = temp in kelvins; n = number of moles of electrons; F = Faraday constant; Q = reaction quotient for reaction at a given point in time; assuming T = 298 K simplified to: E sub cell = E, degree sign, sub cell - 0.0592/n x logQ

Question 43

reaction quotient reminder

Answer 43

for reaction aA + bB -> cC + dD, Q = ([C]^c[D]^d)/([A]^a[B]^b) only those species that take part in equation

Question 44

a method to determine standard free energy change

Answer 44

delta G, degree sign = -RT (ln Keq); R = ideal gas constant; T = absolute temp; Keq = equilibrium constant for reaction

Question 45

relationship between two expressions for delta G, degree sign

Answer 45

delta G, degree sign = -nFE, degree sign, sub cell = -RT (ln Keq) or nFE, degree sign, sub cell = RT (ln Keq)

Question 46

reminders about natural log (ln) in equations for free energy

Answer 46

ln of any number between 0 and 1 is negative; log is positive when equilibrium constants are greater than 1, negative when equilibrium constants are less than 1 and 0 when equilibrium constants are equal to 1

Question 47

negative E, degree sign, sub cell

Answer 47

redox reactions with equilibrium constants less than 1, equilibrium state favors the reactants; ln of any number between 0 and 1 is negative (nonspontaneous)

Question 48

positive E, degree sign, sub cell

Answer 48

the equilibrium constant for redox reaction is greater than 1 (equilibrium state favors products); positive because ln of any number greater than 1 is positive

Question 49

E, degree sign, sub cell = 0

Answer 49

Keq = 1 and concentrations of products and reactants are equal at equilibrium. ln of 1 is 0.

Question 50

change in Gibbs free energy of an electrochemical cell with varying concentrations

Answer 50

delta G = delta G, degree sign + RT (ln Q); where delta G = free energy change under nonstandard conditions; delta G, degree sign = free energy change under standard conditions; R = ideal gas constant; T = temp; Q = reaction quotient