Electroplating

Electroplating is the process of depositing a thin, uniform layer of a more valuable, more attractive, or more corrosion-resistant metal onto the surface of another (usually cheaper or more reactive) metal object, by electrolysis.

Three choices set up the cell:

• Cathode = the object being plated. A spoon, a ring, a bumper, any metal item — it must be clean, greaseless, and free of oxide film, or the new metal layer will not stick.
• Anode = a pure rod or plate of the plating metal itself. Silver-plating uses a silver anode, gold-plating a gold anode, nickel-plating a nickel anode — never a different metal, and never an inert electrode.
• Electrolyte = a soluble salt of that same metal. Silver nitrate for silver, copper sulphate for copper, nickel sulphate for nickel.

At the cathode, the metal cation from solution is reduced and joins the object's surface as solid metal:

where: $M^{n+}$ is the plating metal's cation (e.g. Ag⁺, Ni²⁺, Au³⁺), $n$ is its charge, and $M$ is the neutral metal atom added to the coating. At the anode, the pure metal rod dissolves, releasing fresh cations back into the electrolyte:

Because the anode dissolves at essentially the same rate the cathode deposits, the electrolyte's metal-ion concentration stays steady through the whole process — metal is simply ferried from the anode, through the solution, onto the cathode. This is the one structural feature that makes electroplating different from ordinary electrolysis of a fixed electrolyte: the "reactant" being consumed at the anode is manufactured there on purpose, matched to what the cathode is consuming.

Conditions for a good electroplate (smooth, firmly adherent, evenly distributed):

• The cathode surface must be thoroughly degreased and de-oxidised before plating begins.
• Current must be low and steady, with the object rotated or moved, so metal builds up slowly and evenly rather than in a rough, flaky layer.
• The electrolyte concentration must stay steady — automatically arranged by using an anode of the same metal.
• Electrode spacing and shape should give roughly uniform current density over the whole surface being coated.

Electroplating is purely descriptive chemistry at this level — no Faraday's-law mass calculation is expected — but the reasoning behind the set-up is worth being able to explain in full sentences, because ICSE examiners test exactly this logic with "what if" scenarios. The core insight is a closed loop of the same metal. Any electrolytic cell obeys the usual discharge rules: cations go to the cathode and are reduced, anions (or the anode material itself, if it is reactive) supply oxidation at the anode. In an ordinary electrolysis — say, of copper sulphate with inert platinum electrodes — the cathode plates copper while the anode discharges hydroxide ions or oxygen, and the copper ion concentration in the solution steadily falls as electrolysis proceeds, since nothing is putting copper back in. That is fine for extracting copper, but useless for plating, because a shrinking ion supply means an ever-thinner, ever-slower deposit, and eventually the plating stops. Electroplating fixes this by choosing the anode deliberately to be the same metal as what is being deposited. A reactive metal anode is oxidised in preference to the surrounding anions, so instead of oxygen or another anion product forming, the anode itself dissolves: $M \rightarrow M^{n+} + n e^-$. Every mole of metal ion removed from solution at the cathode is matched, electron for electron, by a mole of the same metal ion re-entering solution at the anode. The bath's composition is therefore self-sustaining — you could, in principle, run it for a very long time and never need to top up the electrolyte, because the anode is manufacturing exactly what the cathode consumes. This is also why the anode metal must match the plating metal and cannot be swapped for something "close enough." If the anode is a different, more easily oxidised metal, that metal dissolves instead, and its ions accumulate in the electrolyte while the intended plating ion is depleted with nothing replacing it — the process described precisely in Worked Example 2 below.

Question: Describe how you would electroplate an iron spoon with nickel, naming the electrodes and electrolyte, and writing the electrode reactions.

Step 1. Identify the cathode. The object to be plated is always the cathode, since that is where reduction and deposition occur — so the iron spoon (cleaned and degreased first) is the cathode.

Step 2. Identify the anode and electrolyte. Since nickel is being deposited, the anode must be pure nickel metal, and the electrolyte must be a soluble nickel salt — nickel sulphate, NiSO₄, solution is standard.

Step 3. Write the cathode reaction. Ni²⁺ ions from the electrolyte migrate to the spoon and are reduced:

Step 4. Write the anode reaction. The pure nickel anode dissolves, replacing the Ni²⁺ removed at the cathode:

Step 5. State the outcome. With a low, steady current and the spoon kept moving in the bath for even coverage, a smooth, adherent nickel coating builds up over the whole spoon, protecting the iron underneath from rusting and giving it a bright finish.

Competitive stretch

Question: A student tries to gold-plate a ring using gold chloride electrolyte but a copper anode instead of a gold one. What goes wrong?

Step 1. The cathode is unaffected at first — Au³⁺ ions still migrate to the ring and deposit as gold: $\text{Au}^{3+} + 3e^- \rightarrow \text{Au}$.

Step 2. The anode reaction is wrong for the job — copper, not gold, dissolves: $\text{Cu} \rightarrow \text{Cu}^{2+} + 2e^-$, since copper is being oxidised instead of replenishing gold ions.

Step 3. Consequence: gold ions are steadily consumed at the cathode and never replaced, so the gold-ion concentration falls, plating becomes patchy or stops, and unwanted Cu²⁺ builds up and may co-deposit, contaminating the coating. Fix: the anode must be pure gold, not copper.

• Using an anode of a different metal from the one being plated — the anode metal and the electrolyte's metal must always match the plating job.
• Forgetting to degrease/clean the cathode object first — a greasy or oxidised surface stops the coating from adhering, however correct the electrical set-up is.
• Mixing up which electrode is which — the object being plated is always the cathode (reduction site), never the anode.
• Assuming a high current plates faster "for free" — too high a current gives a rough, flaky, poorly adherent deposit; low and steady wins.

Ask the tutor: Want to work through a "what electrode/electrolyte would you use" plating scenario, or check why a mismatched anode spoils the coating? Ask the AI tutor to walk it through with you.