SSS 2: ELECTROCHEMICAL SERIES
ELECTROLYSIS
Electrolysis can be defined as the process by which a solution or a molten substance is broken down into its components parts by the passage of electric current.
Electrolytic cell is an assembly of two electrodes in an electrolyte. It is used for the electrolysis of a substance
Basic Terms in Electrolysis
1. Electrolyte: A solution of acids, bases, salts or substances in molten state that can conducts electricity and is decomposed by it.
Types of electrolytes
There are two types of electrolytes namely:
(a) Strong electrolytes
(b) Weak electrolytes
Strong Electrolytes
These are substances that completely dissociate into ions when dissolved in water, allowing electricity to pass easily.
Examples:
(i) Strong mineral acids: HCl, HNO₃, H₂SO₄
(ii) Strong alkalis: NaOH, KOH
(iii) Soluble ionic salts: NaCl solution, CuSO₄ solution
Equations:
HCl(aq) → H⁺(aq) + Cl⁻(aq)
NaOH(aq) → Na⁺(aq) + OH⁻(aq)
NaCl(aq) → Na⁺(aq) + Cl⁻(aq)
Weak Electrolytes:
These are substances that partially dissociate into ions in water, so they conduct electricity poorly.
Examples: CH₃COOH, H₂CO₃, NH₄OH
Equation:
CH₃COOH ⇌ CH₃COO⁻(aq) + H⁺(aq)
Note:
Most organic compounds, such as organic acids and alcohols, are weak electrolytes.
2. Non-electrolyte: They are solutions of substances that does not conduct electricity, even in molten state e.g. sugar, alcohol, urea, hexene, ethanol, ether, benzene, trichloromethane etc.
3. Electrodes: They are conductors in the form of wires, rods or plates through which current enters or leaves the electrolyte. They are the poles of the cell.
There are two types of electrodes namely:
(i) Anode
(ii) Cathode
Anode
•The anode is the positive electrode where oxidation occurs (loss of electrons).
• Electrons enter the external circuit through the anode
Cathode
•The cathode is the negative electrode where reduction occurs (gain of electrons).
4.Ions: They are the charged particles in an electrolyte. Cations (positive ions) move to the cathode (negative electrode), while anions (negative ions) move to the anode (positive electrode)
Differences between conductors and electrolytes
Mechanism of Electrolysis
When electric current passes through an electrolyte:
1. The electrolyte dissociates into positive and negative ions.
2. The cations (positive ions) move to the cathode, where they gain electrons and are reduced. [At the cathode (–): positive ions (cations) gain electrons → reduction]
3. The anions (negative ions) move to the anode, where they lose electrons and are oxidized. [At the anode (+): negative ions (anions) lose electrons → oxidation]
4. The products of electrolysis are formed at the electrodes.
5. The discharge of ions means the conversion of ions into neutral atoms or molecules by gain or loss of electrons at the electrodes.
The mechanisms of electrolysis is based on the ionic theory and depends on wether the electrolyte is in molten form or an aqueous solution.
Arrhenius Ionic Theory
The Arrhenius Ionic Theory, proposed by Svante Arrhenius (1884), states that when an electrolyte dissolves in water, it dissociates into positive and negative ions. These ions are responsible for conducting electricity and taking part in chemical reactions.
Acids produce H⁺ ions in water.
e g. HCl → H⁺ + Cl⁻
Bases produce OH⁻ ions in water.
e g. NaOH → Na⁺ + OH⁻
Salts produce both positive and negative ions.
e.g. NaCl → Na⁺ + Cl⁻
Strong electrolytes ionize completely, while weak electrolytes ionize partially.
The theory explains electrical conductivity in solutions but is limited to aqueous (water-based) reactions and cannot explain acid-base behavior in non-aqueous solvents.
Factors Affecting the discharge of ions
The factors affecting the discharge of ions are:
1. Position of ions in the electrochemical Series
2. Concentration of the ions
3. Nature of the electrode
1. Position of ions in the electrochemical Series
The electrochemical (reactivity) series shows how easily ions gain or lose electrons.
•Cations higher in the series (like Na⁺, K⁺, Ca²⁺, Mg²⁺) are less easily discharged.
•Cations lower in the series (like Cu²⁺, Ag⁺, H⁺) are more easily discharged.
•Cations migrate to the cathode(negative electrode)
Cations
K⁺
Na⁺
Ca²⁺
Mg²⁺
Al³⁺
Zn²⁺
Fe²⁺
Sn²⁺
Pb²⁺
H⁺
Cu²⁺
Hg²⁺
Ag⁺
Au³⁺
Order of ease of preferencial discharge (for cations):
K⁺ > Na⁺ > Ca²⁺ > Mg²⁺ > Al³⁺ > Zn²⁺ > Fe²⁺ > H⁺ > Cu²⁺ > Ag⁺
Anions
F⁻
SO₄²⁻
NO₃⁻
Cl⁻
Br⁻
I⁻
OH⁻
Order of ease of preferencial discharge (for anions):
F⁻ > SO₄²⁻ > NO₃⁻ > Cl⁻ > Br⁻ > I⁻ > OH⁻
Note:
The lower ion in the electrochemical series is preferentially discharged.
Example: Electrolysis of acidified water
Dilute tetraoxosulphate (VI)acid (H₂SO₄) is added to make water conduct electricity because pure water is a poor conductor.
At the Cathode
Hydrogen ions (H⁺) gain electrons (reduction) to form hydrogen gas.
2H⁺(aq) + 2e⁻→ H₂(g)
At the Anode
Hydroxide ions (OH⁻) or water molecules lose electrons (oxidation) to form oxygen gas.
2H₂O(l) → O₂(g) + 4H⁺ + 4e⁻
Overall reaction:
2H₂O(l) → 2H₂(g) + O₂(g)
Observation:
•Bubbles of hydrogen gas form at the cathode.
•Bubbles of oxygen gas form at the anode.
•The volume of hydrogen collected is twice that of oxygen.
2. Relative Concentration of the Ions
When two ions of similar positions in the electrochemical series are present, the ion with higher concentration is preferentially discharged.
Example: Electrolysis of concentrated NaCl (brine) solution
NaCl(aq) → Na⁺ + Cl⁻
H₂O(l) → H⁺ + OH⁻
At the Anode
Ions present: Cl⁻ and OH⁻
Cl⁻ is discharged because it is more concentrated than OH⁻.
2Cl⁻ → Cl₂(g) + 2e⁻
At the cathode
Ions present: H⁺ and Na⁺
H⁺ (lower in series) is discharged.
2H⁺ + 2e⁻→ H₂(g)
Products:
Cathode: Hydrogen gas (H₂)
Anode: Chlorine gas (Cl₂)
However, if the NaCl solution is dilute, OH⁻ (from water) is preferentially discharged instead of Cl⁻.
3. Nature of the Electrode
The type of electrode (inert or active) can also affect which ions are discharged.e.g.
(a) Use of Inert (inactive or passive) Electrodes
(b) Use of Active Electrodes
(c) Use of Mercury as the Electrode
(a) Use of Inert (inactive or passive) Electrodes
Examples: Platinum (Pt), Graphite (C), and Gold (Au).
•They do not react with the electrolyte.
•They do not affect the products at the electrodes.
•They do not influence which ions are discharged; the discharge depends only on the type of ion and concentration.
Example:
Electrolysis of aqueous CuSO₄ using inert electrodes:
CuSO₄(aq) → Cu²⁺ + SO₄²⁻
H₂O(l) → H⁺ + OH⁻
At the cathode
ions present: Cu²⁺ and H⁺
Cu²⁺ (lower) is discharged.
Cu²⁺ + 2e⁻→ Cu
At the anode
ions present: SO₄²⁻ and OH⁻
OH⁻ (lower) is discharged.
4OH⁻ → 2H₂O(l) + O₂ + 2e⁻
Products:
Cathode: Cu deposited
Anode: O₂ gas evolved from OH⁻
(b) Use of Active Electrodes
Examples: Copper (Cu), Silver (Ag)
•They participate in the reaction.
•The electrode metal itself can dissolve into the solution or deposit on the other electrode.
Example:
Electrolysis of CuSO₄ solution using copper electrodes
CuSO₄(aq) → Cu²⁺ + SO₄²⁻
H₂O(l) → H⁺ + OH⁻
At the cathode
ions present: Cu²⁺ and H⁺
Cu²⁺ + 2e⁻ → Cu
•Copper is deposited at the cathode
•mass of the cathode increases
At the anode
ions present: SO₄²⁻ and OH⁻
Possible reactions:
I. Discharge of SO₄²⁻
II. Discharge of OH⁻
III.Copper anode dissolves into ions
Cu → Cu²⁺ + 2e⁻ (anode dissolves)
Results:
I. Copper anode dissolves into ions
II. Mass of the anode decreases
Observation:
The anode gradually dissolves, and copper is deposited at the cathode.
Application: Used in electroplating and purification of copper.
(c) Use of mercury as the electrode
When mercury is used as an electrode especially as the cathode, it affects the products and reactions during electrolysis due to its unique chemical properties.
(I) Mercury as Cathode
Mercury has a strong tendency to form amalgams (alloys) with certain metals such as sodium, potassium, and zinc.
Instead of the metal being deposited as a solid on the electrode surface, it dissolves in mercury to form an amalgam.
Example: In the electrolysis of aqueous sodium chloride (NaCl) using a mercury as cathode:
At the cathode
Na⁺ + 2e⁻→ Na (amalgam)
Later, this amalgam can react with water to produce sodium hydroxide and hydrogen gas:
2Na(Hg) + 2H₂O(l) → 2NaOH + H₂ + 2Hg
(II) Mercury as Anode
Mercury is generally not used as an anode because it can oxidize to form mercuric ions (Hg²⁺), which is undesirable and toxic.
Electrolysis of molten substances
Electrolysis of a molten substance is the process by which an ionic compound in its molten (liquid) state is decomposed into its elements when an electric current is passed through it.
Examples:
1. Electrolysis of Molten Sodium Chloride (NaCl) using graphite electrodes
At the Cathode
Na⁺ + e⁻ → Na
sodium metal is deposited
At the Anode
2Cl⁻ → Cl₂ + 2e⁻
chlorine gas is evolved
Products:
•Sodium metal is deposited at the cathode
•Chlorine gas is evolved at the anode
2. Electrolysis of Molten Lead(II) Bromide (PbBr₂)
Molten PbBr₂ → Pb²⁺ + 2Br⁻
At the Cathode
Pb²⁺ + 2e⁻ → Pb (lead metal deposited)
At the Anode
2Br⁻ → Br₂ + 2e⁻ (bromine gas evolved)
Products:
•Lead metal (Pb) is deposited at the cathode
•Bromine gas (Br₂) is evolved at the anode
Application or uses of electrolysis
Electrolysis is of great importance in the industry. Some of the main applications are:
1. Electroplating:
A process whereby a thin layer of a metal is deposited on another metal by electrolysis. i.e. coating one metal with another (e.g. coating iron with silver or gold).
I. Electroplating protects against rusting or corrosion of metals
II. It makes objects look more attractive by coating them with shiny metals such as gold, silver, or chromium
III. It improves hardness and wear resistance of the object plated
IV. It repairs worn-out parts of a metal
V. Electroplating can provide a smooth surface that reduces friction
2. Extraction of Metals:
Used to extract metals like sodium, potassium, magnesium, and aluminium from their molten salts.
3. Purification of Metals:
Used to purify copper and zinc.
4. Manufacture of Chemicals:
Electrolysis of brine (NaCl solution) produces chlorine, hydrogen, and sodium hydroxide.
5. It improves corrosion resistance and the appearance of aluminium substance.
Faraday’s Laws of Electrolysis
1. Faraday's First Law of Electrolysis:
Faraday's First Law of Electrolysis states that the mass of an element deposited at the electrode during electrolysis is directly proportional to the quantity of electricity passed.
m ∝ Q
m = zIt
Where: m = mass of substance (g)
I = current in ampere (A)
t = time in seconds (s)
Q = Quantity of electricity. It is
measured in Coulomb [C]
z = electrochemical equivalent in
gram per Coulomb [gC⁻¹]
m=zIt
z=m/It
1/z = It/m
The reciprocal of a, 1/z is referred to as the charge-to-mass ratio. It is expressed in Coulomb per gram [C g⁻¹].
Examples:
1. A current of 0.72A was passed through a solution of an electrolyte for 3 hours 20 minutes. Calculate the quantity of electricity that was passed.
Solution:
I= 0.72A
t= 3 hours 20 minutes
= (3 × 60minutes) + 20 minutes
= 180 minutes+ 20 minutes
= 200 minutes
= 200 × 60 seconds
= 12000 seconds
Q = It
Q = 0.72 × 12000
Q = 8640C
2. When a current of 0.20A was passed through an electrolyte, 1200C of electricity was used. Calculate the time of current flow.
Solution:
Q = 1200C, I= 0.20A
Q = It
t = Q/I
= 1200/0.20
= 6000 seconds
3. Calculate the mass of silver deposited when 0.2A of electricity is passed through a silver chloride solution for 2 hours.
[Chemical equivalent of silver= 0.000098gC⁻¹]
Solution:
t= 2hours = 2 × 60 × 60 seconds
= 7200 seconds
I= 0.2A
m = zIt
= 0.000098 × 0.2 × 7200
= 0.14g
4. Calculate the chemical equivalent of copper tetraoxosulphate (VI) solution for 10 minutes depositing 0.04g of copper.
Solution:
t= 10 × 60 seconds = 600 seconds
I = 0.5A, m = 0.04g
m= zIt
z= m/It
= 0.04 / (0.5 × 600)
z= 0.00013gC⁻¹
2. Faraday's Second Law of Electrolysis
States that when the same quantity of electricity passes through different electrolytes, the masses of substances deposited are directly proportional to their chemical equivalents.
OR
States that when the same quantity of electricity passes through solutions of different electrolytes, the relative number of moles of elements deposited is inversely proportional to the charges on the ions of the elements.
1 mole of electron (e⁻) = 1 Faraday
= 96500 Cmol⁻¹
Note:
RMM= Relative Molecular Mass
RMM= number of mole(s) of eletrons
Examples:
1. Calculate the number of Faraday or number of moles of eletrons that are required to form the following during electrolysis.
(a) One mole of sodium atom
(b) One mole of copper atom
(c) One mole of chlorine atom
(d) Two moles of oxygen molecule
(e) Three moles of silver atom
Solution:
(a) Na(s) + e⁻ → Na⁺(aq)
1 mole of Na atom →1 mole of e⁻
1 mole of Na⁺ = 1 mole of e⁻
= 1 Faraday
(b) Cu(s) + 2e⁻ → Cu²⁺ (aq)
1 mole of Cu atom →2 moles of e⁻
1 mole of Cu²⁺ = 2 mole of e⁻
= 2 Faradays
(c) Cl + e⁻ → Cl⁻(aq)
1 mole of Cl atom →1 mole of e⁻
1 mole of Cl⁻ = 1 mole of e⁻
= 1 Faraday
(d) 2O₂(g) + 8e⁻ →4O²⁻(aq)
2 moles of O₂ molecules→8 moles of e⁻
4 atoms of O (2O₂) = 8 moles of e⁻
4 moles of O²⁻(4O²⁻) = 8 moles of e⁻
= 8 Faradays
(e) 3Ag(s) + 3e⁻ → 3Ag⁺ (aq)
3 moles of Ag atom →3 moles of e⁻
3 moles of Ag⁺ = 3 moles of e⁻
= 3 Faradays
2. Calculate the mass of silver deposited when a current of 2.6A is passed through solution of a silver salt for 70 minutes
[Ag= 108, 1F= 96500Cmol⁻¹]
Solution:
Q = It
= 10,920 C
Ag⁺(aq) + e⁻ → Ag(s)
RMM of Ag→ 1 mole of e⁻
108g of Ag → 1 mole of e⁻
108g of Ag → 1 Faraday
108g of Ag → 96500 C
96500 → 108g of Ag
10920 → (10920/96500) × 108
= 12.2g
3. Calculate the mass of copper formed when 965C of electricity passes through a copper (II) tetraoxosulphate (VI) solution.
[Cu= 64, Faraday= 96500 Cmol⁻¹]
Solution:
Cu²⁺(aq) + 2e⁻ → Cu(s)
RMM of Cu→ 2 moles of e⁻
64g of Cu → 2 Faraday
64g of Cu →2 × 96500 C
2 × 96500 → 64g of Cu
965 → [965/(2×96500)] × 64
= 0.32g
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