OME General Chemistry
Lecture 10: Electrochemistry
Dr. Vladimir Lesnyak
Office: Physical Chemistry, Erich Müller-Bau, r. 111
Email: [email protected]
Phone: +49 351 463 34907
Outline
Oxidation–Reduction Reactions
1 Electrolytes
2 Balancing Oxidation–Reduction Equations
Voltaic Cells
3 Construction of Voltaic Cells
4 Notation for Voltaic Cells
5 Cell Potential
6 Standard Cell Potentials and Standard Electrode Potentials
7 Equilibrium Constants from Cell Potentials
8 Dependence of Cell Potential on Concentration
9 Some Commercial Voltaic Cells
Electrolytic Cells
10 Electrolysis of Molten Salts
11 Aqueous Electrolysis
2
Electrolytes and Nonelectrolytes
Electrolyte is a substance that dissolves in
water to give an electrically conducting
solution. In general, ionic solids that
dissolve in water are electrolytes.
Example: NaCl, table salt
Nonelectrolyte is a substance that
dissolves in water to give a nonconducting
or very poorly conducting solution.
Example: sucrose C12H22O11, ordinary table
sugar
3
Electrical Conductivity of a Solution
Pure water
Solution of sodium chloride
4
Strong and Weak Electrolytes
Strong electrolyte exists in solution almost
entirely as ions:
most ionic solids that dissolve in water.
H2 O
NaCl(s) → Na+(aq) + Cl−(aq)
Weak electrolyte dissolves in water to
give a relatively small percentage of ions:
most soluble molecular substances.
NH3(aq) + H2O(l) ⇄ NH4+(aq) + OH−(aq)
5
Solubility Rules for Ionic Compounds
Rule
Statement
Applies to
Exceptions
1
Group IA and ammonium
compounds are soluble
Li+, Na+,
K+, NH4+
—
2
Acetates and nitrates are
soluble
C2H3O2−,
NO3−
—
3
Most chlorides, bromides, and − − − AgCl, Hg2Cl2, PbCl2, AgBr, HgBr2, Hg2Br2, PbBr2,
Cl , Br , I
AgI, HgI2, Hg2I2, PbI2
iodides are soluble
4
Most sulfates are soluble
5
Most carbonates are insoluble CO32−
Group IA carbonates, (NH4)2CO3
6
Most phosphates are insoluble PO43−
Group IA phosphates, (NH4)3PO4
7
Most sulfides are insoluble
Group IA sulfides, (NH4)2S
8
Most hydroxides are insoluble OH−
SO42−
S2−
CaSO4, SrSO4, BaSO4, Ag2SO4, Hg2SO4, PbSO4
Group IA hydroxides, Ca(OH)2, Sr(OH)2, Ba(OH)2
6
Oxidation–Reduction Reactions
Reactions involving a transfer of electrons from one species to another:
Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
Fe(s) + Cu2+(aq) → Fe2+(aq) + Cu(s)
7
Oxidation Numbers
Oxidation number (or oxidation state) of an atom in a substance is the
actual charge of the atom if it exists as a monatomic ion, or a hypothetical
charge assigned to the atom in the substance by simple rules.
In oxidation-reduction reaction one or more atoms change oxidation
number → transfer electrons.
oxidized
0
0
+2 ‒2
reduced
2Ca(s) + O2(g) → 2CaO(s)
8
Oxidation-Number (ON) Rules
Rule
Applies to
Statement
1
Elements
ON of an atom in an element is 0
2
Monatomic
ions
ON of an atom in a monatomic ion = the charge on the ion
3
Oxygen
ON of oxygen = −2 in most of its compounds (Exception: O in H2O2
and other peroxides = −1)
4
Hydrogen
ON of H = +1 in most of its compounds (−1 in binary compounds
with a metal, such as CaH2)
5
Halogens
ON of F = −1 in all of its compounds. Each of the other halogens (Cl,
Br, I) has ON = −1 in binary compounds, except when the other
element is another halogen above it in the periodic table or the
other element is oxygen.
6
Compounds The sum of ONs of the atoms in a compound = 0. The sum of the
and ions
ONs of the atoms in a polyatomic ion equals the charge on the ion.
9
Oxidation-Reduction Reactions
oxidation
0
+2
+2
0
Fe(s) + Cu2+(aq) → Fe2+(aq) + Cu(s)
reducing
agent
oxidizing
agent
reduction
Half-reactions:
0
+2
Fe(s) → Fe2+(aq) + 2e‒
+2
0
Cu2+(aq) + 2e‒ → Cu(s)
(electrons lost by Fe)
oxidation
(electrons gained by Cu2+)
reduction
Oxidation is the half-reaction in which there is a loss of electrons
by a species (or an increase of oxidation number of an atom).
Reduction is the half-reaction in which there is a gain of electrons
by a species (or a decrease in the oxidation number of an atom).
10
Common Oxidation-Reduction Reactions
1. Combination reaction: two substances combine to form a third substance.
2Na(s) + Cl2(g) → 2NaCl(s)
not oxidation-reduction: CaO(s) + SO2(g) → CaSO3(s)
2. Decomposition reaction: a single compound reacts to give two or more substances.
∆
2HgO(s) → 2Hg(l) + O2(g)
∆
not oxidation-reduction: CaCO3(s) → CaO(s) + CO2(g)
3. Displacement reaction: an element reacts with a compound, displacing another
element from it.
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
4. Combustion reaction: a substance reacts with oxygen, usually with the rapid
release of heat to produce a flame.
4Fe(s) + 3O2(g) → 2Fe2O3(s)
11
Activity Series of the Elements
12
Balancing Oxidation-Reduction Equations
0
+1
+2
0
Zn(s) + Ag+(aq) → Zn2+(aq) + Ag(s)
Zn → Zn2+
Ag+ → Ag
(oxidation)
(reduction)
Zn → Zn2+ + 2e‒
Ag+ + e‒ → Ag
(oxidation half-reaction)
(reduction half-reaction)
1 × (Zn → Zn2+ + 2e‒)
2 × (Ag+ + e‒ → Ag)________________________
Zn(s) + 2Ag+(aq) + 2e‒ → Zn2+(aq) + 2Ag(s) + 2e‒
Zn(s) + 2Ag+(aq) → Zn2+(aq) + 2Ag(s)
13
Skeleton Oxidation-Reduction Equations
To set up the skeleton equation and then balance it, we need to
answer the following questions:
1. What species is being oxidized (or, what is the reducing agent)?
What species is being reduced (or, what is the oxidizing agent)?
2. What species result from the oxidation and reduction?
3. Does the reaction occur in acidic or basic solution?
+2
+7
+3
+2
Fe2+(aq) + MnO4‒(aq) → Fe3+(aq) + Mn2+(aq)
(acidic solution)
14
Balancing Oxidation-Reduction Equations
in Acidic Solution
Step 1: Assign oxidation numbers to each atom.
Step 2: Split the skeleton equation into 2 half-reactions: oxidation and reduction.
Step 3: Complete and balance each half-reaction.
a. Balance all atoms except O and H.
b. Balance O atoms by adding H2O’s to one side of the equation.
c. Balance H atoms by adding H+ ions to one side of the equation.
d. Balance electric charge by adding electrons (e‒) to the more positive side.
Step 4: Combine 2 half-reactions to obtain the final balanced oxidation-reduction
equation.
a. Multiply each half-reaction by a factor such that when the half-reactions
are added, the electrons cancel.
b. Simplify the balanced equation by canceling species that occur on both
sides, and reduce the coefficients to smallest whole numbers.
15
+2
Balancing Oxidation-Reduction Equations
in Acidic Solution
Fe2+(aq)
↓
+2
Fe2+(aq)
+2
+7
+ MnO4
‒(aq)
→
Fe3+(aq)
+2
+ Mn2+(aq)
(acidic solution)
+3
→ Fe3+(aq)
+3
Fe2+(aq) → Fe3+(aq) + e‒
↓
+7
+3
(oxidation half-reaction)
+2
MnO4‒(aq) → Mn2+(aq)
↓
MnO4‒(aq) → Mn2+(aq) + 4H2O(l)
↓
MnO4‒(aq) + 8H+(aq) → Mn2+(aq) + 4H2O(l)
↓
MnO4‒(aq) + 8H+(aq) + 5e‒ → Mn2+(aq) + 4H2O(l) (reduction half-reaction)
↓
5 × (Fe2+ → Fe3+ + e‒)
1 × (MnO4‒ + 8H+ + 5e‒ → Mn2+ + 4H2O)__________
5Fe2+ + MnO4‒ + 8H+ + 5e‒ → 5Fe3+ + Mn2+ + 4H2O + 5e‒
↓
5Fe2+(aq) + MnO4‒(aq) + 8H+(aq) → 5Fe3+(aq) + Mn2+(aq) + 4H2O(l)
16
Additional Steps for Balancing Oxidation-Reduction
Equations in Basic Solution
Step 5: Note the number of H+ ions in the equation. Add this number of OH– ions to both
sides of the equation.
Step 6: Simplify the equation by noting that H+ reacts with OH– to give H2O. Cancel any
H2O’s that occur on both sides of the equation and reduce the equation to
simplest terms.
MnO4‒(aq) + SO32‒(aq) → MnO2(s) + SO42‒(aq)
(basic solution)
2MnO4‒ + 3SO32‒ + 2H+ → 2MnO2 + 3SO42‒ + H2O
2MnO4‒ + 3SO32‒ + 2H+ + 2OH– → 2MnO2 + 3SO42‒ + H2O + 2OH–
2H2O
2MnO4‒(aq) + 3SO32‒(aq) + H2O(l) → 2MnO2(s) + 3SO42‒(aq) + 2OH–(aq)
17
Construction of Voltaic Cells
Voltaic cell consists of 2 half-cells that are electrically connected.
Electrochemical cell consists of electrodes that dip into an electrolyte and
in which a chemical reaction either uses or generates an electric current.
Voltaic (or galvanic cell) is an electrochemical cell in which a spontaneous
reaction generates an electric current.
Electrolytic cell is an electrochemical cell in which an electric current drives
an otherwise nonspontaneous reaction.
18
Construction of Voltaic Cells
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Zn electrode and Cu electrode,
without an external circuit ‒ no cell
reaction
2 electrodes are connected by an external
circuit ‒ chemical reaction occurs
Electrode at which oxidation occurs – anode.
Electrode at which reduction occurs – cathode.
19
Notation for Voltaic Cells
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
salt bridge
always on the left
Zn(s)|Zn2+(aq) || Cu2+(aq)|Cu(s)
anode
always on the right
cathode
phase
boundary
Zn(s)
anode
terminal
|
Zn2+(aq)
anode
electrolyte
20
Notation for Voltaic Cells
When the half-reaction involves a gas
Hydrogen electrode: hydrogen gas bubbles over a
platinum surface, where the half-reaction occurs:
2H+(aq) + 2e‒ ⇄ H2(g)
cathode: H+(aq)|H2(g)|Pt
anode: Pt|H2(g)|H+(aq)
Cathode
Cathode reaction
Cl2(g)|Cl‒(aq)|Pt
Cl2(g) + 2e‒ ⇄ 2Cl‒(aq)
Fe3+(aq), Fe2+(aq)|Pt
Cd2+(aq)|Cd(s)
Fe3+(aq) + e‒ ⇄ Fe2+(aq)
Cd2+(aq) + 2e‒ ⇄ Cd(s)
21
Cell Potential
Potential difference is the difference in electric potential between two points.
Volt V is the SI unit of potential difference.
The electrical work expended in moving a charge through a conductor:
Electrical work = charge × potential difference
Joules
= coulombs × volts
The Faraday constant, F, is the magnitude of charge on one mole of electrons;
F = 9.6485×104 C per mole of electrons (96,485 C/mol e‒).
faraday is a unit of charge = 9.6485×104 C.
In moving this quantity of charge (1 faraday) from one electrode to another, the
work done by a voltaic cell:
w = ‒F × potential difference
22
Cell Potential
The maximum potential difference between the electrodes of a voltaic
cell is the cell potential or electromotive force (emf) of the cell Ecell.
It can be measured by an electronic digital voltmeter.
The maximum electrical work of a voltaic cell for molar amounts
of reactants (according to the cell equation as written) is:
wmax = ‒nFEcell
n – the number of moles of electrons transferred in the overall cell equation,
Ecell – the cell potential,
F – the Faraday constant, 9.6485×104 C/mol e‒.
23
Standard Cell Potentials and
Standard Electrode Potentials
Cell potential is a measure of the driving force of the cell reaction.
reduced species → oxidized species + ne‒ (oxidation/anode)
oxidized species + ne‒ → reduced species (reduction/cathode)
Ecell = oxidation potential + reduction potential
Oxidation potential for a half-reaction = −reduction potential for the reverse half-reaction
Reduction potentials are tabulated as electrode
potentials, E.
24
Standard Cell Potentials and
Standard Electrode Potentials
Zn(s)|Zn2+(aq) || Cu2+(aq)|Cu(s)
Zn(s) → Zn2+(aq) + 2e‒ −EZn
Cu2+(aq) + 2e‒ → Cu(s)
Zn2+(aq) + 2e‒ → Zn(s)
EZn
ECu
Ecell = ECu + (−EZn) = ECu − EZn
Ecell = Ecathode − Eanode
The cell potential depends on the concentrations of substances
and the temperature of the cell.
The standard cell potential, Eocell, is the emf of a voltaic cell operating
under standard-state conditions (solute concentrations 1 M, gas
pressures 1 atm, temperature usually 25oC).
25
Tabulating Standard Electrode Potentials
The reference chosen for comparing electrode potentials is the
standard hydrogen electrode, it is assigned a potential of 0.0 V.
Zn(s)|Zn2+(aq) || H+(aq)|H2(g)|Pt
Zn(s) → Zn2+(aq) + 2e‒; −EoZn
2H+(aq) + 2e‒ → H2(g); EoH2
Ecell = EoH2 + (−EoZn)
EoZn = − 0.76 V
26
Standard Electrode Potentials in Aqueous Solution at 25oC
27
Strengths of Oxidizing and Reducing Agents
oxidized species + ne‒ → reduced species
The strongest oxidizing agents in a table of standard electrode potentials are the oxidized
species corresponding to half-reactions with the largest (most positive) Eo values.
reduced species → oxidized species + ne‒
The strongest reducing agents in a table of standard electrode potentials are the reduced
species corresponding to half-reactions with the smallest (most negative) Eo values.
28
Calculation of Cell Potentials Using
Standard Reduction Potentials
Cd2+(aq) + 2e‒ → Cd(s); EoCd = −0.4 V
Ag+(aq) + e‒ → Ag(s); EoAg = 0.8 V
Eocell = Eocathode − Eoanode
Eocell = EoAg − EoCd = 0.8 − (−0.4) = 1.2 V
29
Equilibrium Constants from Cell Potentials
∆𝐺𝐺 = 𝑤𝑤𝑚𝑚𝑚𝑚𝑚𝑚
𝑜𝑜
∆𝐺𝐺𝑜𝑜 = −𝑛𝑛𝑛𝑛𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
Free energies of reaction are generally of the order of 10s to 100s of kilojoules.
Example:
Zn(s) + 2Ag+(aq) → Zn2+(aq) + 2Ag(s)
−Eo = 0.76 V
Zn(s) → Zn2+(aq) + 2e‒
2Ag+(aq) + 2e‒ → 2Ag(s)
Eo = 0.80 V__________
Zn(s) + 2Ag+(aq) → Zn2+(aq) + 2Ag(s) Eocell = 1.56 V
−
𝑜𝑜
∆𝐺𝐺𝑜𝑜 = −𝑛𝑛𝑛𝑛𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
= −2 𝑚𝑚𝑚𝑚𝑚𝑚 𝑒𝑒 × 96,485
𝐶𝐶
5
− × 1.56 𝑉𝑉 = −3.01 × 10 𝐽𝐽
𝑚𝑚𝑚𝑚𝑙𝑙 𝑒𝑒
30
Equilibrium Constants from Cell Potentials
𝑜𝑜
∆𝐺𝐺𝑜𝑜 = −𝑛𝑛𝑛𝑛𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
∆𝐺𝐺𝑜𝑜 = −𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅
𝑜𝑜
= 𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅
𝑛𝑛𝑛𝑛𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
𝑜𝑜
𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
𝑜𝑜
𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
𝑅𝑅𝑅𝑅
2.303𝑅𝑅𝑅𝑅
=
𝑙𝑙𝑙𝑙𝑙𝑙 =
𝑙𝑙𝑜𝑜𝑜𝑜𝐾𝐾
𝑛𝑛𝑛𝑛
𝑛𝑛𝑛𝑛
0.0592
=
𝑙𝑙𝑜𝑜𝑜𝑜𝐾𝐾 (𝑖𝑖𝑖𝑖 𝑣𝑣𝑣𝑣𝑣𝑣𝑣𝑣𝑣𝑣 𝑎𝑎𝑎𝑎 25𝑜𝑜𝐶𝐶)
𝑛𝑛
31
Relationships among K, ∆Go, and Ecell
32
Dependence of Cell Potential on Concentration:
Nernst Equation
The cell potential of a cell depends on the concentrations of ions and on gas
pressures → cell potentials provide a way to measure ion concentrations.
∆𝐺𝐺 = ∆𝐺𝐺𝑜𝑜 + 𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅
𝑜𝑜
∆𝐺𝐺𝑜𝑜 = −𝑛𝑛𝑛𝑛𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
∆𝐺𝐺 = −𝑛𝑛𝑛𝑛𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
𝑜𝑜
−𝑛𝑛𝑛𝑛𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐 = −𝑛𝑛𝑛𝑛𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
+ 𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅
𝑜𝑜
𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐 = 𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
−
𝑅𝑅𝑅𝑅
2.303𝑅𝑅𝑅𝑅
𝑜𝑜
−
𝑙𝑙𝑙𝑙𝑄𝑄 = 𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
𝑙𝑙𝑜𝑜𝑜𝑜𝑜𝑜
𝑛𝑛𝑛𝑛
𝑛𝑛𝑛𝑛
first derived by the German chemist Walther Nernst (1864–1941)
𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐 =
𝑜𝑜
𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
0.0592
−
𝑙𝑙𝑜𝑜𝑜𝑜𝑜𝑜 (𝑖𝑖𝑖𝑖 𝑣𝑣𝑣𝑣𝑣𝑣𝑣𝑣𝑣𝑣 𝑎𝑎𝑎𝑎 25𝑜𝑜𝐶𝐶)
𝑛𝑛
Cell potential Ecell decreases as the cell reaction proceeds going to 0.
33
Determination of pH
Glass electrode
+
𝑝𝑝𝑝𝑝 = −log[𝐻𝐻 ]
pH of a solution can be obtained very accurately
from cell potential measurements, using the Nernst
equation to relate cell potential to pH.
The electrode solution is separated from the test
solution by a thin glass membrane, which
develops a potential across it depending on the
H+ concentrations on its inner and outer surfaces.
Glass electrode is an example of an ion-selective electrode.
Many electrodes have been developed that are sensitive to a
particular ion, such as K+, NH4+, Ca2+, or Mg2+. They can be
used to monitor solutions of that ion.
34
Commercial Voltaic Cells
Zinc–carbon (Leclanché dry cell)
Zn(s) → Zn2+(aq) + 2e‒
(anode)
2NH4+(aq) + 2MnO(s) + 2e‒ → Mn2O3(s) + H2O(l) + 2NH3(aq) (cathode)
The first battery was invented by Alessandro
Volta about 1800. He assembled a pile
consisting of pairs of zinc and silver disks
separated by paper disks soaked in salt water.
A battery cell that became popular during the
19th century was constructed in 1836 by the
English chemist John Frederick Daniell.
The voltage of this cell is initially about 1.5 V
35
Commercial Voltaic Cells
Alkaline dry cell
Zn(s) → Zn2+(aq) + 2e‒
2MnO2(s) + H2O(l) + 2e‒ →
Mn2O3(s) + 2OH−(aq)
Lithium–iodine battery
(anode)
(cathode)
used to power heart pacemakers
36
Commercial Voltaic Cells
Lead storage cell
Pb(s) + HSO4−(aq) → PbSO4(s) + H+(aq) + 2e‒ (anode)
PbO2(s) + H+(aq) + HSO4−(aq) + 2e‒ → PbSO4(s) + 2H2O(l)
(cathode)
After the lead battery is discharged, it is recharged from an external electric current.
Each cell delivers about 2 V, a battery consisting of six cells in series gives about 12 V
37
Commercial Voltaic Cells
Nickel-cadmium cell (nicad cell)
Cd(s) + 2OH−(aq) → Cd(OH)2(s) + 2e‒
(anode)
NiOOH(s) + H2O(l) + 2e‒ → Ni(OH)2(s) + OH−(aq) (cathode)
These half-reactions are reversed when the cell is recharged.
Nicad batteries can be recharged and discharged many times.
38
Commercial Voltaic Cells: Hydrogen–oxygen fuel cell
Battery with a continuous supply of energetic reactants (fuel).
2H2(g) + O2(g) → 2H2O(l)
H2(g) → 2H+(aq) + 2e‒
(anode)
Potential = 0.7 V
O2(g) + 4H+(aq) + 4e‒ → 2H2O(l)
(cathode)
39
Rusting is also an Electrochemical Process
O2(g) + 2H2O(l) + 4e‒ → 4OH−(aq)
Fe(s) → Fe2+(aq) + 2e‒
Fe2+(aq) + 2OH−(aq) → Fe(OH)2(s)
4 Fe(OH)2(s) + O2(g) → 2Fe2O3·H2O(s) + 2H2O(l)
40
Cathodic Protection against Corrosion
f
41
Electrolysis of Molten Salts: NaCl
In electrolytic cell electric current drives an otherwise nonspontaneous reaction.
The process of producing a chemical change in an electrolytic cell is electrolysis.
Many important substances, including aluminum and chlorine, are produced
commercially by electrolysis.
Na+(l) + Cl−(l) → Na(l) + ½Cl2(g)
Downs cell
Li, Mg, and Ca metals are obtained by
the electrolysis of their chlorides.
42
Aqueous Electrolysis: NaCl
The electrolysis of aqueous NaCl is the basis of the chlor-alkali industry ‒
the major commercial source of Cl2 and NaOH.
Chlor-alkali membrane cell
reduction:
2H2O(l) + 2e‒ → H2(g) + 2OH−(aq)
oxidation:
2H2O(l) → O2(g) + 4H+(aq) + 4e‒
possible cathode half-reactions:
Na+(aq) + e‒ → Na(s); Eo = −2.71 V
2H2O(l) + 2e‒ → H2(g) + 2OH−(aq); Eo = −0.83 V
2H2O(l) + 2Cl−(aq) → H2(g) + Cl2(g) + 2OH−(aq)
43
Electroplating of Metals
Purification of copper by electrolysis
44
Stoichiometry of Electrolysis
Michael Faraday (1831): the amounts of substances released at the electrodes during
electrolysis are related to the total charge that has flowed in the electric circuit.
1 faraday (9.6485×104 C) = the charge on one mole of electrons
Electric charge = electric current × time lapse
Ampere (A) is the SI unit of current. Coulomb (C) is the SI unit of
electric charge, equivalent to an ampere-second.
A current of 0.5 amperes flowing for 84 seconds gives a charge of
0.5 A × 84 s = 42 A·s, or 42 C.
From the amount of substance produced at an electrode and
the time of electrolysis, we can determine the current.
From the current and the time of electrolysis, we can calculate
the amount of substance produced at an electrode.
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𝑜𝑜
𝑤𝑤𝑚𝑚𝑚𝑚𝑚𝑚 = ∆𝐺𝐺𝑜𝑜 = −𝑛𝑛𝑛𝑛𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
𝑜𝑜
𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
=
𝑜𝑜
𝐸𝐸𝑐𝑐𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎
−
𝑜𝑜
𝐸𝐸𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎
Summary
𝑜𝑜
𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐 = 𝐸𝐸𝑐𝑐𝑐𝑐𝑐𝑐𝑐𝑐
−
0.0592
𝑙𝑙𝑜𝑜𝑜𝑜𝑜𝑜 (𝑖𝑖𝑖𝑖 𝑣𝑣𝑣𝑣𝑣𝑣𝑣𝑣𝑣𝑣 𝑎𝑎𝑎𝑎 25𝑜𝑜𝐶𝐶)
𝑛𝑛
Oxidation-reduction reactions involve a transfer of electrons from one species to another.
Electrochemical cells are of 2 types: voltaic and electrolytic.
Voltaic cells use a spontaneous chemical reaction to generate an electric current. Electrolytic
cells use an external voltage source to push a reaction in a nonspontaneous direction.
Electrons flow in the external circuit from the anode to the cathode.
Cell potential is the maximum voltage of a voltaic cell.
The standard free-energy change, standard cell potential, and equilibrium constant are all
related. Electrochemical measurements can provide equilibrium or thermodynamic information.
Electrode potential depends on concentrations of the electrode substances, according to the
Nernst equation.
Voltaic cells are used commercially as batteries. The basic principle of the voltaic cell is
employed in the cathodic protection of buried pipelines and tanks.
The electrolysis of an aqueous solution often involves the oxidation or reduction of water.
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