Electrochemistry

Electrochemistry links electrical energy to chemical change. The same principles that explain how batteries work also explain how aluminium is extracted from its ore, how jewellery is silver-plated, and why iron rusts faster when in contact with a less reactive metal.

Conductors and Non-Conductors¶

Substances that allow electric current to pass through them are conductors; those that do not are non-conductors (insulators).

Metallic and Electrolytic Conduction¶

There are two fundamentally different ways a substance can conduct electricity:

Strong and Weak Electrolytes¶

An electrolyte is a molten or aqueous substance containing mobile ions that can carry electrical charge.

A strong electrolyte ionises completely, producing a high concentration of ions and conducting electricity well. Examples: NaCl, HCl, H₂SO₄, NaOH.

A weak electrolyte ionises only partially, so the ion concentration is low and conductivity is poor. Examples: ethanoic acid (CH₃COOH), ammonia solution (NH₃(aq)), carbonic acid.

A non-electrolyte produces no ions at all and does not conduct. Examples: sugar solution, distilled water, ethanol.

Definitions Related to Electrolysis¶

Electrolysis is the decomposition of an electrolyte by the passage of electricity.

Cathode is the negative electrode. Reduction occurs here — positively charged cations are attracted to it and gain electrons.

Anode is the positive electrode. Oxidation occurs here — negatively charged anions are attracted to it and lose electrons.

Memory aid: AN OX (Anode = Oxidation) and RED CAT (Reduction at Cathode).

Cations (positive ions) move toward the cathode.

Anions (negative ions) move toward the anode.

Ion Drift During Electrolysis¶

When a potential difference is applied across an electrolyte:

• Cations drift toward the negative cathode, where they gain electrons and are discharged (reduced).
• Anions drift toward the positive anode, where they lose electrons and are discharged (oxidised).

Which ion is discharged when more than one is present depends on position in the electrochemical series and concentration.

The Electrochemical Series¶

The electrochemical series ranks elements by their tendency to lose electrons (be oxidised). It is closely related to the reactivity series. Metals high in the series lose electrons readily; metals low in the series hold electrons more tightly.

At the cathode, the ion that is most easily reduced (lowest in the series, closest to the noble metals) is preferentially discharged. At the anode, the ion most easily oxidised is discharged first — generally, halide ions (Cl⁻, Br⁻) are discharged before OH⁻ when present in high concentration.

Electrolysis of Selected Substances¶

Molten Lead(II) Bromide¶

Electrolyte: molten PbBr₂ (contains Pb²⁺ and Br⁻ ions)

Acidified Water (Dilute Sulfuric Acid)¶

Electrolyte: dilute H₂SO₄ (contains H⁺ and OH⁻ ions; SO₄²⁻ is not discharged)

The volume of hydrogen collected is twice the volume of oxygen (reflecting the 2 : 1 molar ratio in the reactions).

Concentrated Brine (Saturated NaCl Solution)¶

Electrolyte: concentrated NaCl(aq) (contains Na⁺, H⁺, Cl⁻, OH⁻)

At high chloride concentration, Cl⁻ is preferentially discharged at the anode rather than OH⁻.

The three products — hydrogen, chlorine, and sodium hydroxide — are all industrially important.

Copper(II) Sulfate Solution¶

The outcome depends on the electrode material.

With inert (carbon or platinum) electrodes:

With copper electrodes (used in copper purification and electroplating):

The Faraday Constant¶

The Faraday constant ($F$) is the quantity of electric charge carried by one mole of electrons:

$F = 96\,500 \text{ C mol}^{-1}$

The quantity of electricity (charge) passed during electrolysis depends on the current and the time:

$Q = It$

where $Q$ is charge in coulombs (C), $I$ is current in amperes (A), and $t$ is time in seconds (s).

Calculations Involving Electrolysis¶

The number of moles of electrons transferred is found from:

$n(\text{electrons}) = \frac{Q}{F} = \frac{It}{F}$

Then the mole ratio from the electrode equation gives moles of product, and mass or volume follows.

Industrial Applications of Electrolysis¶

Extraction of Aluminium¶

Aluminium is too reactive to be reduced by carbon, so electrolysis is used instead. The ore is bauxite (impure aluminium oxide, Al₂O₃).

Process:

1. Bauxite is purified to give aluminium oxide.
2. The Al₂O₃ is dissolved in molten cryolite (Na₃AlF₆). Cryolite lowers the melting point from ~2050 °C to ~950 °C, reducing energy costs significantly and also improving conductivity.
3. Electrolysis of the molten mixture is carried out in large steel tanks lined with carbon (cathode), with carbon anodes suspended in the melt.

The carbon anodes gradually burn away as the oxygen produced reacts with them: $\text{C(s)} + \text{O}_2\text{(g)} \rightarrow \text{CO}_2\text{(g)}$

Anodes must be replaced periodically.

Electroplating¶

Electroplating coats an object with a thin layer of another metal using electrolysis. Purposes include improving appearance, preventing corrosion, and reducing cost (using a cheaper base metal).

Setup for silver plating an object:

• Object to be plated = cathode
• Silver metal = anode
• Electrolyte = silver nitrate solution (AgNO₃)

The silver concentration in the solution stays constant as the anode dissolves at the same rate as silver deposits on the cathode.

Anodising¶

Anodising increases the thickness of the protective oxide layer on aluminium. The aluminium object is made the anode in dilute sulfuric acid. Oxygen produced at the anode reacts with the aluminium surface to build up a thicker, harder Al₂O₃ layer. This layer can be dyed to produce coloured finishes.

Purification of Copper¶

Impure copper (from smelting) is refined using electrolysis in copper sulfate solution. Impure copper is the anode, pure copper is the cathode. Copper dissolves from the anode and deposits as pure copper at the cathode. Insoluble impurities fall to the bottom as anode sludge.

Rusting and Corrosion¶

Corrosion is the gradual destruction of a metal by chemical reaction with its environment. Rusting is specifically the corrosion of iron, forming hydrated iron(III) oxide:

$4\text{Fe} + 3\text{O}_2 + x\text{H}_2\text{O} \rightarrow 2\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$

Both oxygen and water are required — neither alone causes rusting. Salt water accelerates rusting because it improves conductivity and speeds up the electrochemical process involved.

Methods of preventing rusting:

In sacrificial protection, the more reactive metal is oxidised in preference to iron because it loses electrons more readily. As long as the sacrificial metal is present and in electrical contact with the iron, the iron is protected — even if its surface is scratched.