12: Electrochemistry Flashcards
electrochemical cells
contained systems in which oxidation-reduction reactions occur… 3 types galvanic, electrolytic, concentration
anode
electrode where oxidation occurs
cathode
electrode where reduction occurs *AN OX & RED CAT
electromotive force (emf)
corresponds to the voltage or electrical potential difference of the cell… if emf is +, cell is able to release energy (ΔG<0). if emf is -, cell absorbs energy so its nonspontaneous
movement of electrons
- current runs from cathode to anode
- electrons move through an electrode in opposite flow of current from anode to cathode because oxidation (loss) occurs at anode
galvanic/voltaic cells
nonrechargeable batteries
- -ΔG and +Ecell meaning the reaction is spontaneous… energy is harnessed by separaed reduction and oxidation half-reactions
- 2 electrodes are placed in half-cells and connected to each other with conductive material
- electrodes are surounding by aqeous electrolyte solution with cations & anions
- solutions are connected by a salt bridge, which is made of an inert salt
Daniell cell
galvanic cell in which a zinc electrode is placd in an aqueous ZnSO4 solution and a copper electrode is placed in an aqueous CuSO4 solution.
- electrons flow from the zinc anode through the wire to the copper cathode
plating/galvanization
precipitation process onto the cathode itself of reduced ions
salt bridge
made of an inert electrolyte which contains ions that will not react with electrodes or ions in solution but rather diffuse into solution to balance out charge of newly created ions and charge of leftover ions
cell diagram
shorthand notation representing reactions in an electrochemical cell
- reactants and products are always listed from left to right in form: anode | anode soln (conc) || cathode soln (conc) | cathode
- single vertical line indicates a phase boundary
- double verical line indicates presence of salt bridge
ex. Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
electrolytic cells
nonspontaneous reactions that require input of energy (+ ΔG)
- called electrolysis because external voltage source causes chemical compounds to decompose
- anode & cathode can consist of any material
- ex. NaCl is decomposed into Cl2 (g) and Na (l)
- Na+ ions migrate toward cathode to be reduced to Na and Cl- ions migrate toward anode to be oxidized to Cl2
Faraday’s law
liberation of gas and deposition of elements on electrodes is directly proportional to the number of electrons being transferred ruing the redox reaction… normality or gram equivalent weight is used
Mn+ + n e- → M (s)
1 e- carries cahrge of 1.6 x 10-19 C
Faraday’s constant
1 F = 96,485 C, which is amount of charge contained in 1 mole of electrons or 1 equivalent
*round up to 105 C/mol e-
electrodeposition equation
sumarizes process of Faraday’s law and helps determine number of moles of element being deposited on a plate
mol M = It/nF
- I is current
- t is time in seconds
- n is number of electron equivalents
- F is 105 C/mol e-
concentration cell
special type of galvanic cell that contains 2 half-cells connected by a conductive material allowing a spontaneous redox reaction to proceed, which generates current and delivers energy
- electrodes are chemically identical… so current is generated as a function of a concentration gradient established between 2 solutions surrounding electrodes
- current stops when the concentrations of ionic species in the half-cells are equal… V=0 when concentrations are equal
cell membrane of a neuron
best representation of a concentration cell… value of electrical potential across membrane depends on concentrations and charges of ions… allow Vm (resting membrane potential) to be maintained
rechargeable cell/battery
one that can function as both a galvanic and electrolytic cell
ex. lead-acid battery
lead acid/storage battery
rechargeable battery
- as a voltaic cell when fully charged consists of 2 half-cells… Ph anode and porous PbO2 cathode connected by conductive material (H2SO4)
- when fully discharged, consists of 2 PbSO4 electroplated lead electrodes with dilute concentration of H2SO4
Nickel-Cadium batteries
rechargeable cells that consist of 2 half-cells made of solid cadium (anode) and nickel(III) oxide-hydroxide (cathode) connected by a conductive material
*AA & AAA cells
- charging reverses the electrolytic cell potentials
- higher energy denisty than lead-acid batteries
- higher surge current (period of large current early in discharge cycle)
electrode charge designations
- galvanic cell: anode is - and cathode is + because anode is source of electrons which move through conductive material to the cathode
- electrolytic cell: anode is + because it is attached to the positive pole of the external voltage source and attracts anions from solution and cathode is - because external source is used to reverse the charge of an electrolytic cell
- oxidation always occurs at anode
- reduction always occurs at cathode
- cations are always attracted to cathode
- anions are always attracted to anode
isolectric focusing
technique used to separate amino acids or polypeptides based on their isolelectric points (pI)
- charged amino acids migrate toward the cathode
- charged amino acids migrate toward the anode
standard hydrogen electdoe (SHE)
0 V
used as a reference point for reduction potential
reduction potential
tendency of species to gain electrons and to be reduced…
- means more likely to be reduced
standard reduction potential (Eºred)
measured under standard conditions (298 K, 1 atm, 1M)
- more positive Eºred means greater relative tendecy for reduction to occur
- less positive Eºred means greater relatvie tendency for oxidation to occur
reduction potentials for galvanic cells
electrode with more positive reduction potential is the cathode… electrode with the less positive reduction potential is the anode
- reaction is spontaneous (-ΔG) because species with a stronger tendency to gain electrons is actually doing so
reduction potentials for electrolytic cells
electrode with the more positive reduction potential is forced by external voltage to be oxidized and is therefore the anode… electrode with the less positive reduction potential is forced to be reduced as is therefore the cathode
- movement of electrons is in the direction against the tnedency of the the electrochemical species so the reaction is nonspontaneous (+ΔG)
standard electromotive force (emf/ E°cell)
difference in potential (voltage) between 2 half-cells under standard conditions
Eºcell = Eºred, cathode - Eºred, anode
*potential of each electrode does not depend on the size of the elctrode, but rather the identity of the material… Eºred only changes if chemical identity of that electrode is changed
Gibbs free energy (ΔGº)
standard change in free energy
ΔGº=-nFEºcell
- n is number of moles of electrons exchanged
- F= faraday constant = 105 if expressed in coulumbs (J/V), then ΔGº expressed in J not kJ
- Eºcell=standard emf of cell
- galvanic cells have -ΔGº and + Eºcell
- electrolytic cells have + ΔGº and - Eºcell
nerst equation
used when conditions deviate from standard conditions (298 K, 1 atm, a M)
- assume 298 K
- Q=[C]c[D]d/[A]a[B]b
- only species in solution are included
potentiometer
voltmeter that draws no current and gives a more accurate reading of emf of a cell
equilibria
ΔGº = -nFEºcell = -RTlnKeq
*log/ln is + if K>1, 0 if K=1, and - if K<1
Gibbs free energy of an electrochemical cell with varying concentrations
ΔG = ΔGº + RTlnQ
