Redox Reactions: Balancing Equations and Electrode Potential

The oxidation number method balances a redox equation in five clear steps. First, write correct formulas for every reactant and product. Second, assign oxidation numbers to every atom and identify exactly which ones change. Third, calculate how much each atom's oxidation number increases or decreases, then multiply by whatever coefficients make the total increase exactly equal the total decrease, since electrons lost must always equal electrons gained. Fourth, if the reaction happens in solution, add H+ ions (acidic medium) or OH- ions (basic medium) until the total charge matches on both sides. Fifth, balance hydrogen atoms by adding water molecules, then check that oxygen atoms also balance as a final confirmation. Balancing dichromate oxidising sulphite, for instance, means recognising chromium drops by 3 per atom while sulphur rises by 2 per atom, so multiplying chromium's side by 2 and sulphur's side by 3 makes both changes equal at 6 electrons total, before adding H+ and water to finish the job.

The half-reaction (ion-electron) method takes a different, often more intuitive route. Write the unbalanced ionic equation, then split it into two separate half-reactions, one oxidation, one reduction. Balance every atom except oxygen and hydrogen first, then in acidic medium, add water to balance oxygen and H+ to balance hydrogen. Next, balance each half-reaction's own electric charge by adding electrons to whichever side needs them, and multiply each half-reaction by whatever factor makes the electron count identical in both, since electrons can't simply vanish between one half-reaction and the other. Finally, add the two half-reactions together and cancel the now-matching electrons on each side, leaving one clean, fully balanced overall equation. Balancing iron(II) oxidised by dichromate this way means writing Fe2+ to Fe3+ + e- alongside Cr2O72- to 2Cr3+ (needing 14H+ and 7H2O to balance), then multiplying the iron half-reaction by 6 to match dichromate's own 6-electron requirement before combining.

Just as acid-base titrations rely on a colour-changing indicator, redox titrations need some way to visibly mark the exact endpoint, and three genuinely different strategies exist. When the titrant itself is intensely coloured, as with deep purple permanganate, it acts as its own self-indicator: the very first faint, persistent pink tinge signals that every reducing species present has finally been consumed. When no such dramatic colour exists, an external indicator does the job instead; dichromate itself is barely coloured, but it oxidises diphenylamine immediately after the equivalence point, producing a sudden, unmistakable intense blue. A third, cleverer method relies on iodine's own specific chemistry: certain oxidants convert iodide to iodine, which gives an intense blue colour with starch and reacts specifically and stoichiometrically with thiosulphate ions, so the blue colour's sudden disappearance, rather than appearance, marks the exact endpoint.

Oxidation number was always a useful bookkeeping device rather than a literal description of reality, and the concept itself has kept evolving. More recently, chemists have found it more accurate to describe oxidation as a decrease in electron density around an atom, and reduction as an increase in electron density, rather than framing it strictly around whole electrons being lost or gained. This shift matters because most real bonds aren't purely ionic or purely covalent; electron density genuinely does shift, continuously and partially, rather than jumping in discrete, whole-electron steps the older language implies. This evolving understanding is a good reminder that even a chapter's central organising concept, oxidation number here, is itself a model built for convenience and predictive power, refined over time as chemists' understanding of bonding itself has deepened.

Dip zinc directly into copper sulphate solution, and electrons transfer instantly at the point of contact, releasing their energy purely as heat. Separate the zinc and copper into two beakers instead, zinc metal in zinc sulphate solution, copper metal in copper sulphate solution, connect the two metal rods with a wire, and connect the two solutions with a salt bridge (a tube of gelled electrolyte allowing ion migration without letting the solutions mix directly), and something genuinely different happens: the exact same overall reaction proceeds, but now the electrons are forced to travel through the external wire to get from zinc to copper, generating a measurable, usable electric current instead of simply releasing heat at a single point of contact. This setup is called a Daniell cell, and each redox couple, the oxidised and reduced forms of one species together, like Zn2+/Zn or Cu2+/Cu, is written with the oxidised form first, separated from the reduced form by a slash representing the interface between them.

Every electrode in a cell carries its own electrode potential, the tendency of the species at that electrode to gain or hold onto electrons, and this can be measured precisely under standard conditions: unit concentration for dissolved species, 1 atmosphere pressure for any gas, and a temperature of 298 K. Since only relative measurements between two electrodes are actually possible, chemists needed one universal reference point, and by convention, the hydrogen electrode (H+/H2) is assigned a standard electrode potential of exactly 0.00 volts, letting every other electrode's potential be measured relative to it. The sign carries real, specific meaning: a negative standard electrode potential (E°) means that redox couple is a stronger reducing agent than the H+/H2 couple, more eager to lose electrons, while a positive E° means it's a weaker reducing agent, holding onto its electrons more tightly than hydrogen does.

A table of standard electrode potentials, arranged from the most positive at the top (strongest oxidising agents, like fluorine at +2.87 V) to the most negative at the bottom (strongest reducing agents, like lithium at -3.05 V), is genuinely a prediction tool, not just a reference list. A spontaneous reaction always pairs a species higher on the table (a stronger oxidising agent) with a species lower on the table (a stronger reducing agent); a species will spontaneously oxidise anything positioned below it. This is exactly why Fe3+ (E° = +0.77 V) can oxidise I- (E° for I2/I- = +0.54 V), since Fe3+ sits higher, but Fe3+ cannot oxidise Cl- (E° for Cl2/Cl- = +1.36 V), since chlorine sits higher still, meaning the reverse reaction, chlorine oxidising Fe2+, is what actually happens instead. This single table, memorised or looked up, lets a chemist predict the outcome of an enormous number of possible redox reactions without running a single experiment.

Fluorine's position at the very top of the electrode potential table, +2.87 V, the single most positive value there, means it's the strongest oxidising agent among all common elements, capable of stripping electrons away from almost anything, including water itself, which is why displacement reactions among the other halogens are never carried out using fluorine in aqueous solution. There's genuinely no chemical oxidant strong enough to convert F- back into F2; the only way to achieve that conversion is electrolytically, forcing the reaction with an external electric current rather than relying on another chemical species. At the opposite extreme, hydroiodic acid (HI) is the most easily oxidised, and therefore strongest reducing, hydrohalic acid, since iodide's own electrode potential sits comparatively low on the table, readily giving up its electron to almost any oxidising agent it encounters, exactly the behaviour that makes it so useful in iodometric titrations relying on iodine's specific, trackable chemistry.