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Flashcards in ch 12 - Electrochemistry Deck (50)
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1

electrochemical cells

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

2

electrodes

area in electrochemical cells where oxidation and reduction take place

3

anode

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

4

cathode

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

5

electromotive force (emf)

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

6

galvanic (voltaic) cells

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

7

how galvanic cells work

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

8

salt bridge

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

9

Daniell cell

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

10

plating or galvanizing

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

11

cell diagram

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

12

similarities in all electrochemical cells

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

13

differences between electrolytic cells and galvanic cells

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)

14

electrolysis

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

15

equation for number of moles of electrons exchanged in electrolytic cells

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+)

16

charge of one electron

1.6 x 10^-19 coulombs (C)

17

Faraday constant

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

18

electrodeposition equation

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

19

concentration cells

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

20

resting membrane potential (V sub m)

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

21

rechargeable cell (rechargeable battery)

cell that can function as both galvanic and electrolytic

22

lead-acid battery

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

23

net equation for a discharging lead-acid battery

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

24

charging lead-acid cell

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

25

energy density

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

26

nickel-cadmium batteries

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

27

surge currents

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

28

nickel-metal hydride (NiMH) batteries

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

29

isoelectric focusing

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

30

characteristics of Ni-Cd discharging battery

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