⚡ Electrolysis – Full Breakdown
1) What is Electrolysis?
Definition: The chemical decomposition of an electrolyte (in molten or aqueous state) into its constituent elements by passing a direct electric current through it.
Key points:
- It happens in an electrolytic cell (not a voltaic cell — that produces electricity, this consumes it)
- Only works for ionic compounds in molten or aqueous state (so ions can move)
- Direct current (DC) is used — AC would just vibrate the ions, not move them in one direction
Important Definitions (ICSE loves these as 1-2 markers)
| Term | Definition | Example |
|---|---|---|
| Electrolyte | A compound that conducts electricity in molten or aqueous state and decomposes | NaCl (aq or molten), CuSO₄ (aq), H₂SO₄ (aq) |
| Non-electrolyte | A compound that does NOT conduct electricity in any state | Glucose, alcohol, sugar, urea, CCl₄ |
| Strong electrolyte | Completely dissociates in water → conducts well | HCl, NaOH, NaCl, H₂SO₄ |
| Weak electrolyte | Partially dissociates → conducts weakly | CH₃COOH, NH₄OH, H₂CO₃ |
| Electrode | A rod/wire through which current enters/leaves the electrolyte | Pt, Cu, C (graphite) |
| Anode | Positive electrode — connected to positive terminal of battery | Where oxidation happens (loss of electrons) |
| Cathode | Negative electrode — connected to negative terminal of battery | Where reduction happens (gain of electrons) |
| Ions | Charged particles in the electrolyte | Cations (+ve) → move to cathode. Anions (-ve) → move to anode |
| Trick | Meaning |
|---|---|
| PAO | Positive = Anode = Oxidation (loss of electrons) |
| NCR | Negative = Cathode = Reduction (gain of electrons) |
2) Mechanism of Electrolysis — Step by Step
- Ionisation / Dissociation: The electrolyte dissociates into ions (in molten state or in water).
- Migration: Cations move towards cathode (-ve), anions move towards anode (+ve).
- Discharge: At the cathode, cations gain electrons (reduction) and become neutral atoms. At the anode, anions lose electrons (oxidation) and become neutral atoms.
- Liberation / Deposition: The neutral atoms either bubble out as gas or get deposited on the electrode.
3) Active vs Inert Electrodes
| Type | What they do | Examples |
|---|---|---|
| Inert electrodes | Don't participate in reaction — just conduct electricity. The electrolyte itself breaks down. | Platinum (Pt), Graphite (C) |
| Active electrodes | Participate in the reaction — the anode may dissolve into the solution. | Copper (Cu), Silver (Ag), Nickel (Ni) |
Key difference: With Cu electrodes in CuSO₄, the Cu anode dissolves (Cu → Cu²⁺ + 2e⁻) and Cu²⁺ from solution gets deposited at the cathode. This is the principle behind electroplating and electrorefining.
4) Preferential Discharge — The Most Important Concept
When an electrolyte has more than one type of cation or anion, only one gets discharged at each electrode. Which one? That depends on three factors:
Three Factors Governing Preferential Discharge
Factor 1: Position in the Electrochemical Series (most important)
For CATIONS (discharge at cathode):
Lower the reactivity → easier to discharge → discharged first.
For ANIONS (discharge at anode):
Opposite trend. Higher the oxidation potential → easier to discharge.
Factor 2: Concentration of Ions
Higher concentration → more likely to be discharged. This matters especially for Cl⁻ vs OH⁻. In dilute NaCl, OH⁻ is discharged (O₂ gas). In concentrated NaCl, Cl⁻ is discharged (Cl₂ gas). ICSE loves this distinction.
Factor 3: Nature of Electrode
If the anode is active (like Cu), the anode itself dissolves instead of anions being discharged. This is key in electroplating and refining.
1. Identify ALL cations present → which one is lowest in reactivity? That gets discharged at cathode.
2. Identify ALL anions present → which one has highest discharge tendency? Check concentration too.
3. If anode is active (Cu, Ag) → anode dissolves → anions not discharged.
4. If electrolyte is molten → only the ions from the electrolyte exist (no H⁺ or OH⁻ from water).
5) Electrolysis of Key Compounds — One by One
Case 1: Electrolysis of Acidified Water (H₂O + dil. H₂SO₄)
Electrolyte: H₂O with a few drops of H₂SO₄ (to increase conductivity). H₂SO₄ → 2H⁺ + SO₄²⁻. H₂O ⇌ H⁺ + OH⁻.
Electrodes: Pt (inert)
| Acidified Water — Electrode Reactions |
|---|
| Cathode (-): 2H⁺ + 2e⁻ → H₂↑ (burns with pop sound) |
| Anode (+): 4OH⁻ → 2H₂O + O₂↑ + 4e⁻ (relights glowing splint) |
| NOTES: Volume H₂ : O₂ = 2 : 1. H₂SO₄ is NOT decomposed — only water breaks down. |
Detailed Comparison
| Cathode (-) | Anode (+) | |
|---|---|---|
| Ions present | H⁺ (from both acid and water) | SO₄²⁻, OH⁻ |
| Which discharges? | H⁺ | OH⁻ |
| Observation | Colourless gas bubbles. Volume = 2× O₂ | Colourless gas bubbles. Volume = half of H₂ |
Case 2: Electrolysis of Molten NaCl (no water)
Electrolyte: NaCl (l) → Na⁺ + Cl⁻
Electrodes: Pt (inert)
| Molten NaCl — Electrode Reactions |
|---|
| Cathode (-): 2Na⁺ + 2e⁻ → 2Na (s) — greyish-white sodium metal deposited |
| Anode (+): 2Cl⁻ → Cl₂↑ + 2e⁻ — greenish-yellow gas with pungent smell |
| NOTES: Industrial method to produce Na and Cl₂ (Downs process). Molten state = no water → Na⁺ has no competition. |
Case 3: Electrolysis of Aqueous NaCl (brine)
Electrolyte: NaCl (aq) → Na⁺ + Cl⁻. H₂O ⇌ H⁺ + OH⁻.
| Dilute NaCl — Electrode Reactions |
|---|
| Cathode (-): H⁺ discharged → H₂↑ (H⁺ lower in series than Na⁺) |
| Anode (+): OH⁻ discharged → O₂↑ (OH⁻ higher discharge than Cl⁻ in dilute) |
| NOTES: NaOH accumulates in solution (Na⁺ + OH⁻ left behind). |
| Concentrated NaCl (Brine) — Electrode Reactions |
|---|
| Cathode (-): H⁺ discharged → H₂↑ |
| Anode (+): Cl⁻ discharged → Cl₂↑ (high conc. overrides OH⁻ preference) |
| NOTES: This is the KEY comparison: dilute → O₂ at anode, concentrated → Cl₂ at anode. NaOH left in solution. |
Case 4: Electrolysis of CuSO₄ with Inert Electrodes (Pt / Graphite)
Electrolyte: CuSO₄ (aq) → Cu²⁺ + SO₄²⁻. H₂O ⇌ H⁺ + OH⁻.
| CuSO₄ with Pt — Electrode Reactions |
|---|
| Cathode (-): Cu²⁺ + 2e⁻ → Cu (s) — reddish-brown copper deposited (Cu²⁺ lower in series than H⁺) |
| Anode (+): 4OH⁻ → 2H₂O + O₂↑ + 4e⁻ — O₂ gas bubbles (OH⁻ higher discharge than SO₄²⁻) |
| NOTES: Blue colour fades (Cu²⁺ removed from solution). Solution becomes acidic (OH⁻ consumed, H⁺ remains). |
Case 5: Electrolysis of CuSO₄ with Active (Cu) Electrodes
Electrolyte: CuSO₄ (aq). Anode = Cu, Cathode = Cu.
| CuSO₄ with Cu Electrodes — Electrode Reactions |
|---|
| Cathode (-): Cu²⁺ + 2e⁻ → Cu deposited (cathode gains mass, thickens) |
| Anode (+): Cu → Cu²⁺ + 2e⁻ (anode dissolves, loses mass) |
| NOTES: Cu²⁺ concentration stays CONSTANT — for every Cu²⁺ deposited at cathode, one Cu²⁺ enters from anode. Blue colour DOES NOT fade. Basis of electrorefining & electroplating. |
6) Applications of Electrolysis
a) Electroplating
Depositing a thin layer of one metal over another using electrolysis.
- Object to be plated → made the cathode (-)
- Metal to be plated → made the anode (+)
- Electrolyte: A soluble salt of the plating metal
Example — Silver plating a spoon:
- Spoon = cathode, Silver rod = anode, AgNO₃ solution = electrolyte
| Electroplating (Silver on Spoon) |
|---|
| Anode (+): Ag → Ag⁺ + e⁻ (silver rod dissolves) |
| Cathode (-): Ag⁺ + e⁻ → Ag (deposited on spoon) |
| NOTES: Object to be plated = cathode. Plating metal = anode. Salt of plating metal = electrolyte. |
Why electroplate? Prevents rusting, improves appearance, reduces cost (cheap metal coated with expensive metal).
b) Electrorefining of Copper
Anode: Impure copper
Cathode: Pure copper
Electrolyte: CuSO₄ + dil. H₂SO₄
| Electrorefining of Copper |
|---|
| Anode (+): Cu → Cu²⁺ + 2e⁻ (impure Cu dissolves — impurities settle as anode mud) |
| Cathode (-): Cu²⁺ + 2e⁻ → Cu (pure Cu deposited — 99.98%) |
| NOTES: Anode mud contains precious metals (Au, Ag, Pt). Pure Cu at cathode. |
c) Extraction of Metals
Highly reactive metals (Na, K, Mg, Ca, Al) are extracted by electrolysis of their molten compounds because chemical reduction doesn't work.
| Electrolytic Extraction of Reactive Metals |
|---|
| 2Al₂O₃ → 4Al + 3O₂↑ (Hall-Héroult process for Al) |
| 2NaCl (l) → 2Na + Cl₂↑ (Downs process for Na) |
| NOTES: Only for metals above C in reactivity series (K, Na, Ca, Mg, Al). Very energy-intensive. |
7) Important Observations Table (Quick Revision)
| Electrolyte | Electrodes | At Cathode (-) | At Anode (+) |
|---|---|---|---|
| Acidified water | Pt / Graphite | H₂ (gas) | O₂ (gas) |
| Molten NaCl | Pt / Graphite | Na (metal) | Cl₂ (gas) |
| Dil. NaCl (aq) | Pt / Graphite | H₂ (gas) | O₂ (gas) |
| Conc. NaCl (aq) | Pt / Graphite | H₂ (gas) | Cl₂ (gas) |
| CuSO₄ (aq) | Pt / Graphite | Cu (deposit) | O₂ (gas) |
| CuSO₄ (aq) | Cu | Cu (deposit) | Cu dissolves |
| AgNO₃ (aq) | Pt / Graphite | Ag (deposit) | O₂ (gas) |
| CuCl₂ (aq) | Pt / Graphite | Cu (deposit) | Cl₂ (gas) |
1. If electrolyte contains Cu²⁺, Ag⁺, or any metal below H in reactivity series → that metal gets deposited at cathode. H⁺ is NOT discharged because these are less reactive than H.
2. If electrolyte contains Na⁺, K⁺, Ca²⁺, Mg²⁺ (metals above H) → H⁺ from WATER gets discharged at cathode, NOT the metal. That's why aqueous NaCl gives H₂, not Na.
3. At anode, OH⁻ usually beats SO₄²⁻ and NO₃⁻ → so unless the solution is very concentrated in Cl⁻/Br⁻/I⁻, expect O₂ at the anode.
4. In molten state, there's no water → only the ions from the molten salt exist. So the cation has no choice but to discharge (even Na⁺ gets discharged in molten NaCl).
8) Faraday's Laws (Basic Understanding)
First Law: The mass of a substance deposited or liberated at an electrode is directly proportional to the quantity of electricity passed (Q = I × t).
Second Law: The same quantity of electricity deposits or liberates masses of different substances in proportion to their chemical equivalents (equivalent mass).
1 Faraday (F) = 96,500 C = charge of 1 mole of electrons. Passing 1F through an electrolyte deposits 1 equivalent of the substance.
9) Common ICSE Reasoning Questions
Q1: Why is electrolysis of acidified water considered an example of electrolysis of water and not of the acid?
Answer: H₂SO₄ only provides H⁺ ions to increase conductivity. The H₂SO₄ is not consumed — only water decomposes (H₂O → H₂ + ½O₂). The amount of H₂SO₄ remains unchanged. This is why it's called "electrolysis of water" and not of the acid.
Q2: In the electrolysis of CuSO₄ using Pt electrodes, why does the blue colour of the solution fade?
Answer: The blue colour is due to Cu²⁺ ions. Cu²⁺ gets deposited at the cathode as Cu metal. Since no Cu²⁺ is added back (Pt is inert, doesn't dissolve), the concentration of Cu²⁺ decreases → colour fades.
Q3: In the electrolysis of CuSO₄ using Cu electrodes, why does the blue colour NOT fade?
Answer: Because the Cu anode dissolves: Cu → Cu²⁺ + 2e⁻. For every Cu²⁺ deposited at cathode, one Cu²⁺ enters from the anode. The Cu²⁺ concentration remains constant, so the blue colour persists.
Q4: Why is Na deposited at the cathode in molten NaCl but not in aqueous NaCl?
Answer: In molten NaCl, only Na⁺ and Cl⁻ ions exist. Na⁺ has no competition — it must be discharged. In aqueous NaCl, water provides H⁺ ions. H⁺ is lower in the reactivity series (more easily reduced) than Na⁺, so H⁺ gets discharged instead of Na⁺. Hence, H₂ gas is produced, not Na metal.
Q5: Why does dilute NaCl give O₂ at the anode while concentrated NaCl gives Cl₂?
Answer: In the discharge series, OH⁻ has a higher discharge potential than Cl⁻. However, concentration also matters. In dilute NaCl, Cl⁻ concentration is too low to overcome the preference for OH⁻, so OH⁻ is discharged (O₂ gas). In concentrated NaCl (brine), the high concentration of Cl⁻ overrides the preference, and Cl⁻ is discharged (Cl₂ gas).
Q6: Distinguish between electrolytic cell and voltaic (galvanic) cell.
Answer: Electrolytic cell — electrical energy is used to drive a non-spontaneous chemical reaction (electricity → chemical). Voltaic cell — a spontaneous chemical reaction produces electrical energy (chemical → electricity).
Q7: In electroplating, why is the object to be plated made the cathode?
Answer: The metal ions (M⁺) in the electrolyte are positively charged. They migrate towards the cathode (negative electrode). At the cathode, M⁺ + e⁻ → M, the metal deposits on the object. If the object were the anode, it would dissolve instead of gaining a coating.
1. For ANY electrolysis question, first write the ions present (cations and anions separately). Then apply preferential discharge rules.
2. Remember: Cathode = Reduction = Gain of electrons = Metal/H₂ deposited. Anode = Oxidation = Loss of electrons = Non-metal/O₂ liberated.
3. The "molten vs aqueous" distinction is the #1 trick ICSE plays. Always check if the electrolyte is molten or aqueous before answering.
4. For electroplating: object = cathode, plating metal = anode, salt of plating metal = electrolyte. This pattern is universal.
5. Colour changes: CuSO₄ blue fades with Pt electrodes, stays with Cu electrodes. FeSO₄ green fades. CuCl₂ blue-green fades as Cu²⁺ deposits.