Chapter 7: Electrochemistry

Electrochemistry studies the link between chemical reactions and electrical energy. In electrochemical cells, oxidation and reduction take place at electrodes.

Oxidation means loss of electrons or an increase in oxidation number. In many traditional examples it also means gain of oxygen or loss of hydrogen.

Reduction means gain of electrons or a decrease in oxidation number. It may also involve loss of oxygen or gain of hydrogen.

Cu²⁺ + 2e⁻ → Cu (reduction; oxidation number +2 → 0).

Oxidation Is Loss. Reduction Is Gain. Think of electrons. When one substance loses electrons, another gains them. Oxidation and reduction happen together.

Accepts electrons from another species and causes that species to be oxidized. The oxidizing agent itself is reduced.

Donates electrons to another species and causes that species to be reduced. The reducing agent itself is oxidized.

Zn goes from oxidation number 0 to +2. It loses electrons, so it is oxidized and is the reducing agent. Chlorine goes from 0 to −1. It gains electrons, so it is reduced and is the oxidizing agent.

Hydrogen goes from 0 to +1: it is the reducing agent. Chlorine goes from 0 to −1: it is the oxidizing agent.

A substance’s role depends on the reaction. For example, concentrated H₂SO₄ and HNO₃ can act as oxidizing agents in suitable conditions.

Common reducing agents: Al, Na, H₂S and Zn. Common oxidizing agents: KMnO₄ and, in appropriate conditions, H₂SO₄.

A device in which oxidation and reduction occur at electrodes, allowing chemical and electrical energy to be converted.

Always remember: oxidation happens at the anode. Reduction happens at the cathode. This is true in both types of cell.

Electrons move through the metal wires outside the cell. Ions carry charge inside the electrolyte. A bridge or separator that lets ions move helps complete a galvanic cell’s circuit.

During electrolysis, positive ions move towards the cathode. Negative ions move towards the anode. Which ions actually react depends on the substances and conditions.

A substance whose molten state or aqueous solution contains mobile ions and conducts electric current. Not every electrolyte is conductive as a solid.

Forms ions almost completely when dissolved in water, e.g. HCl(aq), HNO₃(aq), NaOH(aq), KOH(aq) and soluble salts such as NaCl(aq).

Forms ions only partly in water, e.g. CH₃COOH(aq), H₂CO₃(aq), NH₃(aq) and H₂S(aq).

Produces almost no mobile ions when dissolved in water. Examples are glucose and sucrose solutions, which conduct very poorly.

A salt that dissolves only slightly is not automatically a weak electrolyte. The small amount that dissolves may separate fully into ions. Strength tells us how fully the dissolved substance forms ions.

Ionization makes ions from neutral particles. Dissociation separates ions already present in a compound.

Ammonia in water is NH₃(aq). Some books write it as NH₄OH. Only part reacts with water to make NH₄⁺ and OH⁻. It is a weak base.

Cathode (reduction): Mⁿ⁺ + ne⁻ → M, if the metal ion is preferentially reduced.

Anode (oxidation): 2X⁻ → X₂ + 2e⁻ for halide ions under suitable conditions.

Cathode: K⁺ + e⁻ → K. Anode: 2Cl⁻ → Cl₂ + 2e⁻. Overall: 2KCl(l) → 2K + Cl₂.

Molten potassium chloride electrolysis involves extremely reactive potassium and hazardous chlorine gas. It is a theoretical illustration, not an activity to attempt.

Uses include extracting metals, purifying copper and coating objects with metal. Industry also uses electrolysis of brine to make chlorine and sodium hydroxide.

The mass of substance deposited or liberated at an electrode is directly proportional to charge passed, for the same electrode reaction, assuming all the charge is used for that reaction.

For equal amounts of electric charge, the masses of different substances produced are proportional to their chemical equivalent masses.

m = (M × I × t)/(n × F), where M is molar mass (g/mol), n is electrons required per ion, and F ≈ 96,485 C per mole of electrons.

Chemical equivalent mass is E = M/n. If all the charge forms the stated product, one Faraday deposits one equivalent mass.

With 100% current efficiency, 96,485 C can deposit about 108 g Ag, 31.75 g Cu or 9 g Al from the stated ions. “100% current efficiency” means all the charge is used to make that product.

One Faraday is the charge of 1 mole of electrons. It is not a unit of current.

A primary cell is designed mainly for one discharge cycle and is not normally recharged. A zinc–carbon dry cell gives approximately 1.5 V.

The zinc container is the negative anode. Zinc loses electrons here.

Manganese dioxide mixed with carbon is the active cathode material. It takes part in reduction.

The graphite rod collects current. Graphite itself is not the active cathode reactant.

A damp paste containing NH₄Cl and/or ZnCl₂ lets ions move. It is the electrolyte.

Zn → Zn²⁺ + 2e⁻.

2MnO₂ + 2NH₄⁺ + 2e⁻ → Mn₂O₃ + 2NH₃ + H₂O. Real chemistry can be more complex.

A battery can contain one or more cells. The cells do not always have to be connected in series.

A secondary battery can be recharged by reversing its discharge chemistry with external electrical energy.

Negative electrode: lead (Pb).

Positive electrode: lead dioxide (PbO₂).

Electrolyte: aqueous sulfuric acid (H₂SO₄).

Pb + PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O.

Charging uses an outside supply to drive the reaction backwards. Each lead–acid cell gives about 2 V. Six cells connected in series give about 12 V, as in many car batteries.

Lead and sulfuric acid are harmful. Used batteries need proper recycling. Do not put them in ordinary rubbish.

Bronze is mainly copper and tin. Duralumin is mainly aluminium with copper, magnesium and manganese. Different types use different proportions.

24-karat gold is almost pure gold. 18-karat gold is 18/24, or 75%, gold by mass. White gold is gold mixed with other metals.

Some solders and amalgams contain harmful metals. Modern uses may require safer substitutes.

Corrosion is the gradual damage to a metal as it reacts chemically or electrochemically with its surroundings.

Rusting is electrochemical corrosion of iron in the presence of oxygen and water. Rust consists mainly of hydrated iron(III) oxides/oxyhydroxides.

At a microscopic anode: Fe → Fe²⁺ + 2e⁻.

At a nearby cathode in aerated neutral water: O₂ + 2H₂O + 4e⁻ → 4OH⁻.

Fe²⁺ joins OH⁻. Further reactions with oxygen form hydrated iron(III) compounds. These make up rust.

Rust has pores and cracks. Water and oxygen can still reach fresh iron, so rusting can continue.

Paint, grease and plastic coatings keep water and oxygen away from metal. If a coating breaks, corrosion may start at the exposed part.

Stainless steel contains chromium. Chromium forms a thin oxide film that slows further reaction. This protective film is called a passive layer. Some stainless steels also contain nickel.

Tin, zinc, nickel and chromium coatings can protect steel. Zinc also protects scratched areas by corroding in place of iron.

Connect steel to a more reactive metal, such as Mg or Zn. That metal loses electrons and corrodes first. It supplies electrons to the steel and helps protect it. The metal used up is the sacrificial anode.

Another method uses an external power supply to keep steel cathodic. This is impressed-current cathodic protection.

Chromium makes a protective oxide film. Zinc can corrode in place of iron. Both protect metal, but they work in different ways.

Cathode: Sn²⁺ + 2e⁻ → Sn(s). With a soluble tin anode: Sn(s) → Sn²⁺ + 2e⁻.

Zinc electroplating uses Zn²⁺ + 2e⁻ → Zn at the steel cathode. Hot-dip galvanizing is different: steel is dipped into molten zinc.

Cathode: Ag⁺ + e⁻ → Ag.

Soluble anode: Ag → Ag⁺ + e⁻.

AgCl dissolves very little in water, so it cannot supply many Ag⁺ ions in a simple aqueous bath. Real silver-plating solutions need carefully controlled chemistry.

Chromium plating needs controlled industrial solutions. Cr³⁺ + 3e⁻ → Cr shows the electron balance. It is not a full set of practical instructions.

Plating chemicals and fumes can be harmful. These examples explain the chemistry; practical work needs trained supervision.

Al, Na and Zn typically act as reducing agents because they lose electrons. H₂S is a reducing agent in reactions such as Br₂ + H₂S → 2HBr + S. KMnO₄ is generally a strong oxidizing agent. Concentrated H₂SO₄ may act as an oxidizing agent under suitable conditions.

H₂ is the reducing agent and gets oxidized (H: 0 to +1); Cl₂ is the oxidizing agent and gets reduced (Cl: 0 to −1).

Strong: HCl(aq), KI(aq), NaOH(aq). Weak: H₂S(aq), CH₃COOH(aq), aqueous ammonia NH₃(aq), often represented NH₄OH(aq).

At the negative cathode: K⁺ + e⁻ → K (reduction). At the positive anode: 2Cl⁻ → Cl₂ + 2e⁻ (oxidation). This is not a safe classroom practical.

Protect metal by a barrier paint or polymer coating, corrosion-resistant alloy, electroplating/metal coating, or cathodic protection. Cathodic protection uses a more reactive sacrificial metal such as magnesium or zinc, or an impressed current.

Approximately 96,485 C, the charge of one mole of electrons.

The cathode is always the site of reduction, both in galvanic and electrolytic cells.

Cu + Zn = brass. Cu + Sn = bronze or bell metal, depending on tin content. A typical duralumin alloy contains Al, Cu, Mg and Mn; its exact composition depends on the alloy grade.

Their ions can move freely in molten form or suitable aqueous solution, carrying electrical charge. In an ionic solid, ions are fixed in a lattice and cannot carry current through the solid.

An electrolytic cell uses external electrical energy to drive a reaction that needs an external supply of electrical energy. It includes a power supply, two electrodes and an electrolyte. Oxidation occurs at the positive anode; reduction occurs at the negative cathode.

An oxidizing agent accepts electrons and is reduced (example Cl₂). A reducing agent loses electrons and is oxidized (example Zn).

Zn: 0 to +2, oxidized, reducing agent. Cl₂: 0 to −1, reduced, oxidizing agent.

Br₂: 0 to −1, reduced, oxidizing agent. Sulfur in H₂S: −2 to 0, oxidized, so H₂S is the reducing agent.

Ca: 0 to +2, oxidized, reducing agent. O₂: 0 to −2, reduced, oxidizing agent.

Li: 0 to +1, oxidized, reducing agent. S: 0 to −2, reduced, oxidizing agent.

Cu–Zn is brass. Cu–Zn–Sn could describe a ternary copper alloy, but conventional bronze is principally Cu–Sn. Duralumin is mainly Al–Cu–Mg–Mn.

A common zinc–carbon dry cell has a zinc anode (negative container), a cathode mix based on MnO₂, a graphite current collector and NH₄Cl/ZnCl₂ paste. Zinc oxidizes to Zn²⁺; MnO₂ is reduced. The cell supplies about 1.5 V and is normally non-rechargeable.

A battery has one or more interconnected electrochemical cells. In a lead–acid rechargeable cell, Pb is the negative plate, PbO₂ the positive plate and H₂SO₄(aq) the electrolyte. Discharge: Pb + PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O. The reaction is driven backward during charging.

A direct-current (DC) supply provides energy to make a redox reaction occur in an electrolyte. Cations move toward the negative cathode and may gain electrons (reduction); anions move toward the positive anode and may lose electrons (oxidation). Example in molten KCl: K⁺ + e⁻ → K; 2Cl⁻ → Cl₂ + 2e⁻.

Rusting is corrosion of iron with oxygen and water, forming hydrated iron(III) oxides/oxyhydroxides. Iron oxidizes at anodic regions: Fe → Fe²⁺ + 2e⁻. Oxygen is reduced at cathodic regions: O₂ + 2H₂O + 4e⁻ → 4OH⁻. Use paint, galvanizing, stainless steel or cathodic protection to reduce corrosion.

The steel object is connected to the negative cathode. Tin: Sn²⁺ + 2e⁻ → Sn. Zinc: Zn²⁺ + 2e⁻ → Zn. Silver: Ag⁺ + e⁻ → Ag. A suitable metal anode may dissolve to replace the ions used up. Proper commercial electrolytes and safety controls are necessary; silver chloride is poorly soluble, and electro-zinc plating differs from hot-dip galvanizing.

First law: mass deposited is proportional to electric charge, m ∝ Q = It. Second law: equal charge deposits masses proportional to their equivalent masses. Combined: m = MIt/(nF), with F ≈ 96,485 C/mol e⁻.

One faraday theoretically deposits 108 g Ag from Ag⁺, 31.75 g Cu from Cu²⁺, or 9 g Al from Al³⁺, assuming complete deposition and the stated rounded molar masses.

Oxidation = loss of electrons; reduction = gain of electrons.

Anode = oxidation; cathode = reduction (always).

Galvanic anode is negative; electrolytic anode is positive.

Mobile ions conduct electricity through an electrolyte; electrons move through the external metal circuit.

Faraday law: Q = It and m = MIt/(nF).

An oxidizing agent gains electrons itself; a reducing agent loses electrons itself.

Graphite is a current collector in the zinc–carbon dry cell, not the cathode active substance.

Rusting needs oxygen and water in ordinary conditions. Sacrificial anodes oxidize to protect steel.

AgCl is sparingly soluble; it is not a good simple soluble Ag⁺ electrolyte for plating.

Copper–tin is typical bronze; copper–zinc is brass. Zinc electroplating is not the same process as hot-dip galvanizing.