Electrochemistry | General Chemistry 3
Electric Energy and Charge
SI units used in electrochemistry:
- Joule: J
- Coulomb: C
- Volt: V (1 V = 1 J.C-1)
- Ampere: A (1 A = 1 C.s-1)
Electrical Energy U (in J):
U = V x Z
V = voltage (in V)
Z = charge (in C)
Electrical Charge Z (in C):
Z = I x t
I = current (in A)
t = time (in s)
Chemical reactions can occur as a result of the passage of an electric current through a solution
Electrochemical Cells
Oxidation-reduction reactions: electrons are transferred from one substance to another
⇒ oxidation-reduction reactions can be used to produce an electric current
Electrochemical cell:
Setup in which an electric current is obtained from a chemical reaction
It consists in 2 electrodes in solution, a salt bridge and an external circuit
Electrode: solid on the surface of which oxidation-reduction reactions occur
Cathode: electrode at which the reduction occurs
Anode: electrode at which the oxidation occurs
Salt bridge: provides an ionic current path between the 2 solutions with electrodes in order to maintain the charge balance in the cell
Zinc-copper electrochemical cell:
The spontaneous chemical reaction is: Zn (s) + Cu2+ (aq) → Zn2+ (aq) + Cu (s)
Zn (s) and Cu (s) are the 2 electrodes
Zn (s) → Zn2+ (aq) + 2 e- [oxidation] ⇒ Zn (s) is the anode
Cu2+ (aq) + 2 e- → Cu (s) [reduction] ⇒ Cu (s) is the cathode
Cell Diagrams
A cell diagram is a representation of an electrochemical cell
Zinc-copper electrochemical cell:
The spontaneous chemical reaction is: Zn (s) + Cu2+ (aq) → Zn2+ (aq) + Cu (s)
The cell diagram is Zn (s) | ZnSO4 (aq) || CuSO4 (aq) | Cu (s)
By convention:
- metal electrodes are written at the ends of the diagram
- left-hand electrode: oxidation reaction ⇒ anode
- right-hand electrode: reduction reaction ⇒ cathode
- single vertical bars: boundaries of phases that are in contact
- double vertical bars: salt bridge
Nernst Equation
Gibbs energy change of the reaction ΔGrxn (in J.mol-1):
ΔGrxn = ZEcell = - νeFEcell
Z = charge per mol (in C.mol-1)
νe = stoichiometric coefficient of the electrons in the 2 half reaction equations
F = Faraday’s constant = 96485 C.mol-1
Ecell = electromotive force emf of the cell (in V)
Relation between the equilibrium constant and the standard cell voltage:
ln K =
K = equilibrium constant
νe = stoichiometric coefficient of the electrons in the 2 half reaction equations
F = Faraday’s constant = 96485 C.mol-1
E0cell = standard cell voltage (in V)
R = ideal gas constant = 8.314 J.mol-1.K-1
T = temperature (in K)
At T = 25°C: ln K = E0cell
Nernst equation:
Relationship between the electromotive force of the cell and the reaction quotient ⇒ electrochemical cells can be used to determine the concentration of ions:
Ecell = E0cell – ln Q
Ecell = electromotive force emf of the cell (in V)
E0cell = standard cell voltage (in V)
R = ideal gas constant = 8.314 J.mol-1.K-1
T = temperature (in K)
νe = stoichiometric coefficient of the electrons in the 2 half reaction equations
F = Faraday’s constant = 96485 C.mol-1
Q = reaction quotient
At T = 25°C: Ecell = E0cell – ln Q
Half-Reaction E° Values
Relationship between E0cell, E0red and E0ox:
E0cell = E0red + E0ox
E0cell = standard cell voltage (in V)
E0red = standard reduction voltage (in V)
E0ox = standard oxydation voltage (in V)
For a particular half reaction: E0ox = - E0red
Cu2+ + 2 e- → Cu [E0 = + 0.34 V]
Cu → Cu2+ + 2 e- [E0 = - 0.34 V]
It is not possible to directly measure the voltage of a single electrode but it can be determined by measuring the voltage difference between a standard electrode and the desired electrode
⇒ in electrochemistry, a Standard Hydrogen Electrode (SHE) is used: nonreactive metal in a 1.0M H+ (aq) solution through which hydrogen gas is bubbled at P = 1.0 bar
By convention, the voltage of the SHE is E0 = 0 V ⇒ SHE is the primary reference electrode
2 H+ (aq) + 2 e- → H2 (g) [E0red = 0 V]
Faraday’s Laws
Electrolysis: process by which a chemical reaction occurs by the passage of an electric current through the solution. It is quantitatively described by Faraday’s laws
First Faraday’s law:
the mass of elements deposited at an electrode is directly proportional to the charge
⇒ m α Q
Second Faraday’s law:
the mass of elements deposited / liberated at an electrode is directly proportional to the molar mass of the elements divided by the change in magnitude of the oxidation state of the substance
⇒ m α
Faraday’s law of electrolysis:
m = x
m = mass deposited as metal or evolved as gas (in g)
I = current (in A)
t = time (in s)
F = Faraday’s constant = 96485 C.mol-1