Electrochemistry Essentials: Electrolysis and Electroplating—Driving Nonspontaneous Reactions

1) Electrolytic cell anatomy: what you must label every time

An electrolytic cell uses an external power supply to force a redox process that would not proceed on its own. Your first job is always to identify the parts and assign polarity correctly.

Core components

• Power supply (DC source): pushes electrons through the external circuit.
• Electrolyte: molten salt or ionic solution that carries charge by ion migration.
• Electrodes: conduct electrons; can be inert (Pt, graphite) or active (made of a metal that can dissolve).

Polarity in electrolytic cells (and how it differs from galvanic)

In an electrolytic cell, the power supply sets the electrode signs:

• Anode is positive (+): connected to the positive terminal of the power supply.
• Cathode is negative (−): connected to the negative terminal.

This is the opposite sign convention from galvanic cells (where the anode is negative and cathode is positive). The reaction roles do not change: oxidation occurs at the anode and reduction occurs at the cathode.

Ion migration checklist

• Cations migrate toward the cathode (−) to be reduced.
• Anions migrate toward the anode (+) to be oxidized.

Inert vs active electrodes: why it matters

• Inert electrode: does not participate chemically; products come from the electrolyte (and possibly water).
• Active electrode: the electrode itself can be oxidized (dissolve) or reduced (plate). This can dominate product formation (especially at the anode in electroplating/electrorefining).

2) Predicting products: molten salts vs aqueous solutions

Product prediction in electrolysis is mainly about which species can be oxidized/reduced at each electrode. The big fork in the road is whether the electrolyte is molten (no water present) or aqueous (water competes).

A. Molten salts: usually straightforward

In a molten ionic compound, the only available redox-active species are its ions.

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Step-by-step method (molten)

• Step 1: Identify the cation and anion present.
• Step 2: At the cathode (−), reduce the cation to its element (often a metal).
• Step 3: At the anode (+), oxidize the anion to its element (often a nonmetal like Cl₂, O₂ from oxide, etc.).

Example: molten NaCl

• Cathode (reduction): Na+ + e- → Na(l)
• Anode (oxidation): 2 Cl- → Cl2(g) + 2 e-

B. Aqueous solutions: water competes at both electrodes

In water-based electrolytes, you must consider that water can be reduced (to H₂) and water can be oxidized (to O₂). Often, the observed products are determined by which process is easier under the conditions.

Key competing half-reactions involving water (conceptual forms)

• Water reduction (cathode possibility): produces H₂(g) and OH⁻ in neutral/basic conditions.
• Water oxidation (anode possibility): produces O₂(g) and H⁺ (or consumes OH⁻ in basic solution).

Step-by-step method (aqueous)

• Step 1: List plausible reductions at the cathode: metal ion reduction vs water reduction to H₂.
• Step 2: List plausible oxidations at the anode: anion oxidation (e.g., halides) vs water oxidation to O₂, plus possible oxidation of an active anode.
• Step 3: Decide which wins based on typical trends:
  • At the cathode: very reactive metal cations (e.g., alkali/alkaline earth) usually do not plate from water; H₂ forms instead. Less reactive metal ions (e.g., Cu²⁺, Ag⁺) commonly plate as metal.
  • At the anode: Cl⁻, Br⁻, I⁻ often oxidize to halogens in concentrated solutions; otherwise O₂ from water oxidation is common. If the anode is an active metal used for plating/refining, it may dissolve instead of evolving O₂.

Example 1: aqueous NaCl with inert electrodes (brine)

• Cathode: water reduction dominates → H₂(g)
• Anode: chloride oxidation can dominate (especially concentrated) → Cl₂(g)
• Solution becomes basic near cathode due to OH⁻ formation.

Example 2: aqueous CuSO₄ with inert electrodes

• Cathode: Cu2+ + 2 e- → Cu(s) (copper plates)
• Anode: sulfate is hard to oxidize; water oxidizes → O₂(g)

Example 3: aqueous CuSO₄ with a copper anode (active)

• Cathode: Cu2+ + 2 e- → Cu(s) (copper plates)
• Anode: Cu(s) → Cu2+ + 2 e- (anode dissolves)
• Result: Cu²⁺ concentration can stay roughly steady while copper transfers from anode to cathode.

3) Electroplating mechanism: how metal coatings actually form

Electroplating deposits a thin layer of metal onto an object by making the object the cathode (−) in an electrolytic cell. Metal ions in solution are reduced to metal atoms on the surface.

Basic setup (conceptual)

• Object to be plated → cathode (−)
• Plating metal source (often the same metal) → anode (+)
• Electrolyte contains ions of the plating metal (plus additives as needed).

What happens at each electrode

Cathode (deposition): metal ions gain electrons and become solid metal on the surface.

M^n+ + n e- → M(s)

Anode (often dissolution if active): the plating metal dissolves to replenish ions.

M(s) → M^n+ + n e-

If the anode is inert, then the anode reaction is usually water oxidation to O₂, and the metal-ion concentration in solution will drop as plating proceeds unless replenished.

Why electrolyte composition matters

• Provides metal ions at a sufficient concentration to sustain deposition.
• Controls pH and conductivity, affecting side reactions (like H₂ evolution) and current distribution.
• Additives (levelers/brighteners) can improve smoothness by influencing nucleation and growth; conceptually, they help prevent rough, dendritic deposits.

Why plating thickness depends on charge passed (Faraday’s idea in practice)

The amount of metal deposited is tied to how many electrons are delivered to the cathode. Each metal ion requires a fixed number of electrons (n) to become metal. Therefore, more total charge (Q) passed means more moles of electrons and more moles of metal deposited.

Step-by-step thickness logic (no heavy math required)

• Step 1: More time at a given current → more charge passes.
• Step 2: More charge → more electrons delivered to the cathode.
• Step 3: More electrons → more metal ions reduced → thicker coating.

In quantitative problems, you typically use Q = I × t and relate charge to moles of electrons, then to moles of metal using the electron requirement n.

4) Reading electroplating and electrorefining diagrams

Diagrams for plating/refining often show two metal pieces in a metal-ion solution connected to a power supply. The key is to identify where metal dissolves and where metal deposits.

Fast identification rules

• Cathode (−): site of metal deposition (mass increases).
• Anode (+): site of metal dissolution if the anode is made of that metal (mass decreases).

Electrorefining: purifying a metal by selective dissolution/deposition

In electrorefining, the impure metal is the anode (+) and a thin sheet of pure metal is the cathode (−). The electrolyte contains ions of the metal being refined.

• Anode: impure metal atoms oxidize to ions and enter solution.
• Cathode: metal ions reduce and plate as pure metal.
• Impurities: some fall off as anode sludge; others may remain dissolved depending on their chemistry.

Diagram interpretation checklist

• Locate the + terminal connection → that electrode is the anode → expect dissolution (if active).
• Locate the − terminal connection → that electrode is the cathode → expect deposition.
• Confirm by ion motion: metal cations drift toward the cathode.

5) Applied exercises (with worked, step-by-step reasoning)

Exercise 1: molten salt electrolysis

System: molten MgCl₂, inert electrodes.

Step 1: Identify ions → Mg²⁺, Cl⁻.

Step 2: Cathode (reduction of cation)

Mg2+ + 2 e- → Mg(l)

Step 3: Anode (oxidation of anion)

2 Cl- → Cl2(g) + 2 e-

Products: Mg(l) at the cathode, Cl₂(g) at the anode.

Exercise 2: aqueous salt with inert electrodes (water competition)

System: aqueous Na₂SO₄, inert electrodes.

Step 1: Candidates at the cathode

• Na⁺ to Na(s) is not favored in water; water reduction to H₂ is typical.

Cathode product: H₂(g) (and OH⁻ builds up locally).

Step 2: Candidates at the anode

• Sulfate is difficult to oxidize; water oxidation to O₂ is typical.

Anode product: O₂(g) (and acidity builds up locally).

What you would observe: gas bubbles at both electrodes; no metal plating.

Exercise 3: electroplating nickel onto steel

System: NiSO₄(aq) electrolyte, steel object as cathode, Ni metal anode.

Step 1: Assign electrodes

• Steel object connected to negative terminal → cathode (−).
• Ni metal connected to positive terminal → anode (+).

Step 2: Write half-reactions

Cathode: Ni2+ + 2 e- → Ni(s)   (nickel deposits on steel)

Anode:   Ni(s) → Ni2+ + 2 e-   (nickel dissolves into solution)

Step 3: Identify products

• Solid formed: Ni(s) coating on the cathode.
• Gas formation: ideally minimal; if current is too high or surface is dirty, competing H₂ evolution can occur at the cathode, causing pitting/poor adhesion.

Exercise 4: interpreting a diagram (electrorefining copper)

Diagram description: A thick impure Cu slab is connected to the + terminal; a thin pure Cu sheet is connected to the − terminal; electrolyte is CuSO₄(aq).

Identify where metal dissolves vs deposits

• Anode (+): impure Cu dissolves: Cu(s) → Cu2+ + 2 e-.
• Cathode (−): Cu deposits: Cu2+ + 2 e- → Cu(s).

Mass changes: anode mass decreases; cathode mass increases.

Exercise 5: improving plating quality (conceptual troubleshooting)

Scenario: You try to plate Zn onto a metal part from an aqueous Zn²⁺ solution, but the coating is rough and patchy with bubbles.

Likely issues and improvements (step-by-step)

• Step 1: Surface preparation
  • Problem: oils/oxides prevent uniform electron transfer and nucleation.
  • Improve: degrease, rinse, mild abrasion/pickling as appropriate, then avoid touching the surface.
• Step 2: Current control
  • Problem: too high current density drives side reactions (often H₂ evolution) and creates rough, dendritic growth.
  • Improve: lower current, use a controlled power supply, increase agitation to reduce concentration gradients.
• Step 3: Electrolyte condition
  • Problem: low metal-ion concentration or wrong pH increases competition from water reduction and causes burning/pitting.
  • Improve: maintain metal-ion concentration, appropriate pH/buffering, and conductivity; filter contaminants.
• Step 4: Geometry and contact
  • Problem: sharp edges get higher current density → thicker/rougher deposits at edges, thin in recesses.
  • Improve: reposition anodes, use shields/thieves, ensure solid electrical contact to the part.