Electrochemistry

Redox Fundamentals, Electrolytes, and Conduction

Every electrochemical reaction relies on oxidation-reduction chemistry. You need to know these four definitions word for word:

• Oxidation means the loss of electrons. In oxidation-number terms, it represents an increase in oxidation number.
• Reduction means the gain of electrons. In oxidation-number terms, it represents a decrease in oxidation number.
• An oxidising agent is a substance that brings about oxidation by accepting electrons (it is itself reduced).
• A reducing agent is a substance that brings about reduction by donating electrons (it is itself oxidised).

Oxidation and reduction always occur together; electrons released by one chemical species are taken up by another. In any cell, oxidation always occurs at the anode and reduction always occurs at the cathode (remember: An Ox, Red Cat, or the acronym CROA: Cathode Reduction, Oxidation Anode).

We describe an electrolyte as a substance that conducts electricity when molten or in aqueous solution and breaks down chemically during the process. Solid ionic compounds cannot conduct electricity because their ions are locked firmly in place within the crystal lattice by strong electrostatic attractions. Once an ionic compound is melted or dissolved in water, the lattice breaks apart and the ions become free to move, carrying electric charge towards the electrodes. A strong electrolyte, such as $\text{NaCl}$, $\text{HCl}$, or $\text{NaOH}$, dissociates fully in water and conducts strongly. A weak electrolyte, such as ethanoic acid or ammonia, ionises only partially and conducts weakly.

A clear demonstration of conduction in the molten state is the electrolysis of molten lead(II) bromide using inert carbon electrodes:

• Cathode (reduction): $\text{Pb}^{2+}(l) + 2e^- \rightarrow \text{Pb}(l)$ (a bead of molten grey lead forms).
• Anode (oxidation): $2\text{Br}^-(l) \rightarrow \text{Br}_2(g) + 2e^-$ (dense red-brown bromine vapour evolves).

Core Electrolysis Investigations with Inert Electrodes

Inert electrodes like platinum or carbon (graphite) conduct electricity and provide a physical surface for electron transfer without reacting chemically or dissolving. When you hook these electrodes up to a d.c. power supply, the external source acts like an electron pump. It pulls electrons away from the anode, leaving that terminal positive, and crowds electrons onto the cathode, making it the negative electrode.

1. Electrolysis of Acidified Water

• Electrolyte: Pure water with a small amount of dilute sulfuric acid added to supply mobile ions.
• Electrodes and apparatus: Platinum or carbon electrodes inside a Hofmann voltameter or positioned beneath inverted test tubes filled with electrolyte.
• Method: Submerge electrodes, connect to a low-voltage d.c. supply, and collect gases by downward displacement of liquid.
• Cathode observation and test: Rapid colourless bubbling. The collected gas pops sharply with a burning splint, confirming hydrogen:

• Anode observation and test: Colourless bubbling. The gas relights a glowing splint, confirming oxygen:

• Volume ratio: The volume of hydrogen collected is approximately twice that of oxygen. By Avogadro's law, equal volumes of gases at the same temperature and pressure contain equal numbers of molecules, so this 2:1 volume ratio directly reflects the 2:1 mole ratio in the overall decomposition: $2\text{H}_2\text{O}(l) \rightarrow 2\text{H}_2(g) + \text{O}_2(g)$.
• Sources of error: The measured hydrogen : oxygen ratio is typically slightly greater than 2 : 1 because oxygen is more soluble in water than hydrogen, meaning a small portion stays dissolved.

2. Electrolysing Copper(II) Chloride Solution

• Electrolyte: Aqueous copper(II) chloride ($\text{CuCl}_2$).
• Electrodes: Two inert carbon rods connected to a d.c. supply.
• Cathode reaction: $\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s)$. A pink-brown (red-brown) deposit of copper metal forms on the electrode.
• Anode reaction: $2\text{Cl}^-(aq) \rightarrow \text{Cl}_2(g) + 2e^-$. Pale green gas bubbles emerge with a sharp smell. Moist blue litmus paper held at the mouth of the tube turns red and is then bleached white, because chlorine dissolves to give an acidic solution and is a powerful oxidising agent.
• Safety: Because chlorine gas is toxic, make sure to carry out this experiment in a fume cupboard or a well-ventilated laboratory.

3. Electrolysing Potassium Iodide Solution

• Electrolyte: Colourless aqueous potassium iodide ($\text{KI}$) containing phenolphthalein indicator.
• Cathode reaction: $2\text{H}_2\text{O}(l) + 2e^- \rightarrow \text{H}_2(g) + 2\text{OH}^-(aq)$. Effervescence yields hydrogen gas (burns with a 'pop'). The solution turns bright pink around the cathode because accumulating hydroxide ions generate an alkaline environment.
• Anode reaction: $2\text{I}^-(aq) \rightarrow \text{I}_2(aq) + 2e^-$. A deep yellow-brown colour spreads around the anode as elemental iodine forms (giving a blue-black colour if starch is added).

Rules of Preferential Discharge

When several ions are present in aqueous solution, they compete at the electrode surfaces. You can predict which substance reacts using these standard rules:

At the Cathode (Reduction — competition between metal cations and water):

• Metals sitting low in the electrochemical series (such as $\text{Cu}^{2+}$, $\text{Ag}^+$, $\text{Pb}^{2+}$) are reduced in preference to water, depositing the pure metal on the electrode.
• Metals sitting high in the electrochemical series (such as $\text{K}^+$, $\text{Na}^+$, $\text{Ca}^{2+}$, $\text{Mg}^{2+}$, $\text{Al}^{3+}$) are not discharged from aqueous solution. Instead, water is more easily reduced, producing hydrogen gas and hydroxide ions:

At the Anode (Oxidation — competition between anions and water):

• Halide ions ($\text{Cl}^-$, $\text{Br}^-$, $\text{I}^-$) are oxidised in preference to water, liberating the free halogen.
• Oxoanions (such as $\text{SO}_4^{2-}$, $\text{NO}_3^-$, $\text{CO}_3^{2-}$) are extremely stable and are never discharged from aqueous solution. Water is oxidised instead, yielding oxygen gas and hydrogen ions:

• Concentration effects: Electrolysing concentrated sodium chloride solution (brine) discharges chloride ions at the anode, yielding chlorine. Very dilute sodium chloride yields oxygen instead because water oxidation becomes competitive.
• Electrolysis of sodium sulfate solution with universal indicator: Using inert carbon electrodes, water is reduced at the cathode forming $\text{OH}^-$, turning universal indicator blue/purple. Water is oxidised at the anode forming $\text{H}^+$, turning the indicator red. Overall, only water is decomposed.

Active Electrodes and Industrial Applications

When electrodes participate chemically in the reaction, they are termed active electrodes. The classic example is electrolysing copper(II) sulfate solution using copper metal strips:

• Anode (copper strip): $\text{Cu}(s) \rightarrow \text{Cu}^{2+}(aq) + 2e^-$. The anode dissolves and steadily loses mass.
• Cathode (copper strip): $\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s)$. Copper deposits onto the cathode, which gains mass.
• Solution appearance: The blue colour of the solution does not fade because $\text{Cu}^{2+}$ ions are replenished by the anode at the exact rate they are removed at the cathode.

Industrial Applications

• Purification (refining) of copper: Impure copper forms the anode; a thin sheet of pure copper forms the cathode in an acidified copper(II) sulfate bath. Copper dissolves from the anode and plates out purely onto the cathode. Valuable unreactive metals like silver and gold do not oxidise and drop to the tank floor as anode sludge.
• Electroplating: The setup examiners look for places the object you want to coat at the cathode, a bar of the plating metal at the anode, and an electrolyte containing ions of that same metal. For instance, silver plating uses a silver anode dipping into a silver nitrate electrolyte, giving an everyday metal piece a clean finish and protecting it from corrosion.
• Chlor-alkali industry (brine electrolysis): Electrolysing concentrated aqueous $\text{NaCl}$ produces chlorine gas at the anode, hydrogen gas at the cathode, and leaves sodium hydroxide ($\text{NaOH}$) solution behind in the cell.
• Extraction of aluminium from bauxite: Purified alumina ($\text{Al}_2\text{O}_3$) is dissolved in molten cryolite ($\text{Na}_3\text{AlF}_6$), which drastically lowers the required working temperature and conserves energy. Carbon linings act as the cathode, where $\text{Al}^{3+} + 3e^- \rightarrow \text{Al}(l)$ yields molten aluminium. Carbon blocks act as the anode ($2\text{O}^{2-} \rightarrow \text{O}_2 + 4e^-$). The oxygen reacts with the hot carbon anodes to produce $\text{CO}_2$, requiring the anodes to be replaced regularly.

The Electrochemical Series, Displacement, and Corrosion

The electrochemical series ranks elements in order of their tendency to lose electrons (ease of oxidation):

A metal higher in the electrochemical series (often referred to casually as the reactivity series) will displace a metal lower in the series from an aqueous solution of its ions:

• Placing a clean zinc strip into blue copper(II) sulfate produces a pink-brown coating of copper, the blue colour fades, and the beaker warms up:

• Placing copper into zinc sulfate produces no reaction, because copper sits lower in the series and cannot displace zinc.
• The halogen series follows the reverse trend for oxidising power: $\text{Cl}_2 > \text{Br}_2 > \text{I}_2$. Chlorine displaces iodine from potassium iodide solution: $\text{Cl}_2(aq) + 2\text{KI}(aq) \rightarrow 2\text{KCl}(aq) + \text{I}_2(aq)$, turning the solution brown.

Corrosion as an Electrochemical Phenomenon

Rusting is the electrochemical oxidation of iron requiring both oxygen and water. Iron is oxidised at anodic sites ($\text{Fe} \rightarrow \text{Fe}^{2+} + 2e^-$) while dissolved oxygen is reduced at cathodic sites ($\text{O}_2 + 2\text{H}_2\text{O} + 4e^- \rightarrow 4\text{OH}^-$), eventually forming hydrated iron(III) oxide. Rusting accelerates in salty or acidic conditions.

Methods of prevention include:

1. Barrier methods: Painting, greasing, plastic coating, or plating with tin.
2. Sacrificial protection: Attaching blocks of a metal higher in the electrochemical series (like zinc or magnesium) to ship hulls or buried pipelines. The more reactive metal is oxidised preferentially, sacrificing itself to preserve the iron.
3. Galvanising: Coating steel with a layer of zinc. Zinc serves as an initial barrier and provides sacrificial protection even if the coating becomes deeply scratched.

Galvanic Cells, Construction, and the Salt Bridge

Galvanic cells, also termed voltaic cells, harness a spontaneous redox reaction to produce an electrical current. When two metals with different oxidation tendencies are placed in solutions of their own ions and linked externally, electrons transfer through the wire from the more reactive metal to the less reactive one.

Cell Assembly and the Salt Bridge

In a standard zinc-copper cell, a zinc rod is placed in $1.0\text{ M }\text{ZnSO}_4$ solution and a copper rod in $1.0\text{ M }\text{CuSO}_4$ solution. A salt bridge—a strip of filter paper soaked in an unreactive electrolyte like $\text{KNO}_3$—connects the two beakers, with both ends dipping below the liquid surfaces.

As zinc oxidises ($\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$), positive zinc ions accumulate in the anode beaker. In the cathode beaker, copper ions discharge ($\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$), leaving behind an excess of negative sulfate ions. If this charge build-up were not relieved, the reaction would stop almost immediately and no current would flow. The salt bridge performs two vital roles: it completes the electrical circuit and maintains electrical neutrality by allowing ions to migrate between the half-cells. Potassium nitrate is chosen because neither $\text{K}^+$ nor $\text{NO}_3^-$ forms insoluble precipitates with the ions in either beaker.

Under standard conditions (1 M solutions at 25 °C) and with no current flowing, this cell gives a reading of about 1.1 V on a high-resistance voltmeter.

Laboratory Apparatus and Experimental Summary Table

How to Draw the Core Apparatus

• Electrolysis of acidified water: Draw a beaker containing dilute sulfuric acid. Sketch two carbon electrodes upright in the liquid, each enclosed by an inverted test tube filled with electrolyte to catch gases. Show wires running from the electrodes to a d.c. power pack. Label: dilute sulfuric acid, carbon electrodes, d.c. power supply, anode (+) showing oxygen, and cathode (−) showing hydrogen with roughly double the gas volume.
• Galvanic cell setup: Draw two beakers side by side. Label the left beaker zinc strip (anode, −) immersed in $1\text{ M }\text{ZnSO}_4$. Label the right beaker copper strip (cathode, +) immersed in $1\text{ M }\text{CuSO}_4$. Draw an inverted U-tube or strip connecting both solutions labelled salt bridge (KNO₃). Connect the metal strips through an external wire to a voltmeter, adding an arrow on the wire showing electrons flowing from zinc to copper.

Master Summary of Electrolyte Behaviours

Key terms

Oxidation

The loss of electrons from a chemical species, corresponding to an increase in oxidation number.

Reduction

The gain of electrons by a chemical species, corresponding to a decrease in oxidation number.

Oxidising Agent

A substance that brings about oxidation by accepting electrons, undergoing reduction itself in the process.

Reducing Agent

A substance that brings about reduction by donating electrons, undergoing oxidation itself in the process.

Electrolyte

A substance that conducts electricity when molten or dissolved in water and undergoes chemical decomposition during the process.

Electrolysis

The chemical decomposition of an electrolyte solution or molten compound caused by the passage of an electric current.

Inert Electrode

An unreactive electrical conductor, such as platinum or graphite, that provides a surface for electron transfer without taking part in the chemical reaction.

Active Electrode

An electrode that takes part chemically in an electrochemical reaction, oxidising and dissolving into solution or changing composition.

Cathode

The electrode at which reduction takes place and electrons are gained.

Anode

The electrode at which oxidation takes place and electrons are lost.

Salt Bridge

A strip of porous paper or gel tube soaked in an unreactive electrolyte (such as KNO₃) that completes the electrical circuit and maintains electrical neutrality between half-cells.

Electrochemical Series

A list of elements arranged in order of their standard ease of losing electrons and forming positive ions.

Displacement Reaction

A chemical reaction where a more reactive element replaces a less reactive element from an aqueous solution of its ions.

Sacrificial Protection

The prevention of metal corrosion by attaching a more reactive metal that oxidises preferentially.