How it Works
This simulator models a potentiometric redox titration: an analyte solution (a reduced species such as an Fe²⁺-like ion) sits in a beaker while an oxidizing titrant is added dropwise from a burette. An inert electrode tracks the solution's potential E as titrant is added. Before the equivalence point, E follows the Nernst equation for the analyte's own redox couple as the ratio of its oxidized to reduced forms changes; after the equivalence point, E follows the Nernst equation for the titrant couple, now present in excess.
Right at the equivalence point, for a simple 1-electron-transfer stoichiometry, E jumps steeply through the average of the two standard potentials E°. The size of that jump is controlled by ΔE°, the separation between the analyte and titrant standard potentials — a bigger ΔE° gives a sharper, more easily detected jump, which is why real titrations pair couples with well-separated E° values. Near the endpoint, the simulated beaker liquid shifts color, mimicking a redox indicator's color change or, for self-indicating titrants like permanganate, the first persistent tint of excess oxidant.
After equivalence: E = E°₂ + (0.0592/n)·log([Ox]/[Red])titrant
At equivalence (n₁=n₂=1): E ≈ (E°₁ + E°₂) / 2
Equivalence volume: Veq = Cₐ·Vₐ / Cₜ
Frequently Asked Questions
What does a redox titration measure?
A redox titration determines the concentration of a reducible or oxidizable analyte by reacting it with a titrant of known concentration that undergoes a complementary electron-transfer reaction. Unlike acid-base titrations, the reaction transfers electrons rather than protons, and progress is followed by watching the solution's electrode potential rather than its pH.
Why is potentiometric (electrode potential) monitoring used instead of just pH?
pH tracks proton concentration, but a redox reaction doesn't necessarily change pH at all — it changes the ratio of oxidized to reduced forms of the species in solution. An inert electrode immersed in the solution develops a potential that reflects that ratio via the Nernst equation, so measuring E against a reference electrode is the direct way to follow a redox titration's progress.
What is the Nernst equation and how is it applied here?
The Nernst equation relates a half-cell's electrode potential to its standard potential E° and the ratio of oxidized to reduced species: E = E° + (0.0592/n)·log([Ox]/[Red]) at 25°C. Before the equivalence point, the solution potential is governed by the analyte's own redox couple; after the equivalence point, it's governed by the titrant couple, now present in excess.
Why does the potential jump sharply right at the equivalence point?
As the titrated fraction approaches 1, the ratio of oxidized to reduced analyte swings from very small to very large over just a tiny addition of titrant, and the logarithm in the Nernst equation amplifies that swing into a steep, near-vertical rise in E. Exactly at equivalence, the potential passes through the average of the two couples' standard potentials.
What determines the size of the potential jump at equivalence?
The size of the jump is set mainly by ΔE°, the difference between the titrant's and analyte's standard potentials. A larger ΔE° produces a bigger, sharper, easier-to-detect jump, which is exactly why redox titrations are designed around couples with well-separated E° values.
What is a redox indicator?
A redox indicator is a dye whose own oxidized and reduced forms have different colors and whose standard potential lies close to the expected equivalence-point potential, so it flips color right as the sharp potential jump occurs, for example ferroin, used in cerium(IV) titrations of iron(II).
How does self-indicating permanganate work?
Potassium permanganate is intensely purple, while its reduction product Mn²⁺ is nearly colorless. During the titration the purple color is consumed as fast as it is added, so the solution stays clear; once past the equivalence point, the first slight excess of permanganate produces a faint, persistent pink that signals the endpoint — no separate indicator is needed.
What's the difference between the equivalence point and the endpoint?
The equivalence point is the exact stoichiometric point, calculated from the reaction's known mole ratio. The endpoint is the point an experimenter actually observes, usually a color change, and it should closely coincide with the equivalence point if the indicator or self-indicating couple is well chosen.
What real-world applications use redox titrations?
Permanganate titrations determine iron content in ores and other samples; iodometric titrations measure the concentration of oxidizing agents in industrial and water-quality analysis; and the Winkler method, a classic iodometric procedure, determines dissolved oxygen concentration in water samples.
How does this connect to galvanic cells and the Nernst equation?
A titration curve is really just a half-cell potential measurement, exactly like the electrode potentials that combine to give a galvanic cell's EMF. The difference is that here the concentrations are continuously changing as titrant is added, so the Nernst equation has to be re-evaluated point by point rather than for one fixed pair of concentrations.