Topic 21: Electrochemistry Flashcards
Description of reduction reactions (3)
a) Gain of e- / O
b) Oxidation number decreases
c) Loss of H
Description of oxidation reactions (3)
a) Loss of e- / O
b) Oxidation number increases
c) Gain of H
What happens in a redox reaction?
One substance always becomes reduced while the other one oxidized
Function of a voltaic cell (2)
a) Use spontaneous redox (ΔG < 0) reactions to produce electrical energy
b) System does work on the surroundings
Structure of a voltaic cell
Two half-cells joined by a wire and a salt bridge
Definition of half-cell
Metal in contact with an aqueous solution of its own ions
Function of a salt bridge
Ions flow through the half-cells to balance the charges and maintain the potential difference.
Electrodes based on REDOX in voltaic cells
a) Anode | Oxidation / - / Left
b) Cathode | Reduction / + / Right
Electrodes based on activity
a) Active - Metals are components of half-reactions
b) Inactive - No reaction components used as electrodes (Pt / Graphite)
Process of voltaic cells
a) External circuit
b) Salt bridge
a) Electricity is produced because e- flow from the oxidized substance to the reduced species where e- are gained
b) The salt bridge completes the circuit and neutralizes any buildup of charge by the following ion movement
Description of ion flow in salt bridge
a) Anions in the salt bridge move from the cathode to the anode
b) Cations move in the salt bridge move from the anode to the cathode
Cell Notation Examples
Pt(s)|H2(g),H+(aq) ‖ Fe3+(aq),Fe2+(aq)| Pt(s)
Zn(s)|Zn2+(aq) ‖ Cu2+(aq)|Cu(s)
Explanation for spontaneity in voltaic cells
Electrons move spontaneously from areas of higher PE to areas of lower PE
a) Each electrode tends to lose or gain e- when it is in contact with the solution of its ions
b) A potential difference between the electrodes will allow the flow of e-
Standard electrode potential Eθ(half-cell)
Electrical potential of a half-reaction to reduce with all components in their standard states
Standard cell potential Eθ(cell)
Potential difference between two electrodes in a voltaic cell to undergo reduction under standard conditions
a) No current flowing
Standard conditions
298K
100kPa
1M
All elements in their standard state
Calculation + Units
Eθcell = Eθ(cathode) - Eθ(anode)
V = J/C
Does cell potential depend on the amount of substance?
No. It is an intensive property
Standard hydrogen electrode definition
Half-cell with solution of 1 M H+ ions at 298 K, H2(g) at 100 kPa and Pt electrode
Importance of SHE
a) Reference point to measure electrode potential of other half cells
b) Linked to a second half-cell through an external circuit
Reaction of SHE
2H+ (aq;1M) + 2e- ↔ H2 (g;1atm)
Eθ = 0 V
Relation between reduction electrode potential and oxiding/reducing strength
a) The higher the electrode potential, the stronger the oxidizing strength
b) The higher the electrode potential, the weaker the reducing strength
Explanation of activity series using electrode potential
a) Metals that can displace H2 from acid
b) Metals that cannot displace H2 from acid
c) Metals that can displace H2 from water
d) Metals that displace other metals from solution
a) Lower electrode reduction potential
b) Higher electrode reduction potential
c) Very low electrode potential
d) Lower electrode potential
Relationship between Ecell and spontaneity
a) Ecell > 0
b) Ecell < 0
c) Ecell = 0
a) Spontaneous
b) Nonspontaneous
c) Equilibrium
Relationship between free energy and standard cell potential
∆G = w(max) = -Ecell × charge
∆G = -nFEcell
Relationship between equilibrium constant and standard cell potential
Ecell = (RT/nF) ln(K)
E°cell = (0.0592/n) log(K) at 298.15 K
Effect of concentration in cell potential
(a) ∆G = ∆G° + RT ln(Q)
(b) Ecell = E°cell - (RT/nF) ln(Q)
Principles of concentration cell
a) Voltaic cell of a concentrated solution of a substance with a dilute solution
b) Final solution has an intermediate concentration
c) Ecell depends on the ratio of concentrations
Process of concentration in
a) Anode (Dilute)
b) Cathode (Concentrated)
a) Oxidation causes solution to become more concentrated
b) Reduction causes solution to become less concentrated
Reaction that measures pH
2H+ (aq;1M) → 2H+ (aq;unknown)
Calculation of pH using Ecell
Ecell = -0.0592/n log([H+unknown)]^2)/([H+standard]^2)
Ecell = 0.0592×pH
Composition of a pH electrode
a) Glass electrode
- Ag/AgCl half-reaction immersed in HCl solution
-Enclosed by a thin glass membrane sensitive to H+ ions
b) Reference electrode
- Pt wire immersed in calomel (Hg2Cl2) paste, liquid Hg, and saturated KCl
Function of electrolytic cells
Use electrical energy from an external source to drive a nonspontaneous redox reaction
Description of electrodes in electrolytic cells
Made up of an inert, heat resistant material like graphite or Pt
Electrodes based on REDOX in electrolytic cells
a) Anode (Oxidation | +)
b) Cathode (Reduction | -)
Structure of electrolytic cells
Electrodes are placed in an electrolyte and connected to the power source by electrical wires
a) Power source = Battery
b) Electrodes cannot touch each other
c) Electrolyte = Liquid substance
Definition of electrolysis
Splitting of a substance by the input of electrical energy into its elements
General reaction during electrolysis
a) Ions attracted by the oppositely charged electrodes
b) Cation is reduced while anion is oxidized
In mixed molten salts, which species is reduced at the cathode?
The more easily reduced
Metals with high Ionization energy
In mixed molten salts, which species is oxidized at the cathode?
The more easily oxidized
Nonmetals with lower electronegativity
Are E values useful in assessing the reduction/oxidation of molten salts?
No. As those refer to species in aqueous solutions
Can pure water easy to electrolyze?
It proceeds rapidly by adding small amount of a salt that cannot be electrolyzed
Oxidation reaction of water
2H2O(l) → O2(g) + 4H+(aq) + 4e-
E = 0.82 V (1.4 V with overvoltage)
Reduction reaction of water
2H2O(l) + 2e- → 2OH-(aq) + H2(g)
E = -0.42 V
Definition of overvoltage
Increment above the expected voltage
Common in the production of gases
Predicting electrode products in cathode
a) Cations of less active metals are reduced (Au/Ag/Cu/Cr/Pt/Cd)
b) Cations of more active metals are not reduced (G1/G2/Al)
=> Production of H2(g) + OH-(aq)
Predicting electrode products in anode
a) Anions that are oxidized include halides, except F-(aq)
b) Anions that are not oxidized include F-(aq) and common oxoanions
- Production of O2(g) and H+(aq)
Faraday law
The amount of substance produced at each electrode is directly proportional to the quantity of charge flowing through the electrolytic cell
Relationship between current and charge
Charge (C) = Current (A) x Time (t)
