A Different Kind of Reaction to Titrate
So far our reactions have passed acid and base back and forth. But a titration works with any reaction that is fast, complete, and has a recognisable finish. A huge family fits that bill: reactions where one substance gives up electrons and another takes them on. Titrations built on this electron-trading are called redox titrations, and they let us measure things acids never could — the iron in an ore, the chlorine in a pool, the vitamin C in juice.
The bench routine is exactly what you already know: deliver a standardized titrant from a burette into the sample, watch for the end point, read the volume, do the arithmetic. What changes is the chemistry doing the counting — and, delightfully, sometimes the way we spot the finish.
Permanganate: The Reagent That Is Its Own Indicator
One of the most beautiful redox titrants is potassium permanganate, used in a permanganate titration. The solution is an intense, almost violent purple. When you drip it into a sample it can oxidise, it reacts and the purple instantly vanishes, leaving a colourless or faintly pink mixture. This continues drop after drop — each drop disappears as fast as it lands.
Then comes the magic. The instant all the analyte has been consumed, the very next drop of purple has nothing left to react with — so the purple no longer disappears. A faint but persistent pink colour suddenly stays. That lingering pink is your end point, and it needs no added indicator at all: the titrant signals its own completion. We call such a reagent self-indicating.
Iodometry: Counting an Oxidant Through Iodine
Some substances you want to measure are oxidisers — bleach, dissolved oxygen, copper in an alloy. A clever indirect method handles them all: iodometry. The idea has two beats. First, let your oxidising sample react with an excess of iodide; the sample converts a portion of that iodide into iodine, and crucially it makes exactly as much iodine as there was sample. The unknown has been faithfully translated into an amount of iodine.
Second, you titrate that liberated iodine with a standardized titrant that reacts with it, a reducing solution. As the iodine is consumed, its tea-brown colour fades toward pale yellow, then to colourless. Right near the end, a little starch added to the flask turns an inky blue-black with the last traces of iodine — and the disappearance of that blue is a sharp, unmistakable end point. Count the titrant, work backward through both beats, and you have the original oxidiser.
When the Direct Route Fails: Back-Titration
Sometimes you cannot titrate an analyte directly. Maybe its reaction with the titrant is hopelessly slow, or it is a stubborn solid that dissolves only grudgingly, or there is no good end point to watch. For these awkward cases there is a graceful workaround called a back titration. The trick is to attack the analyte not with just enough titrant, but with a known, deliberate excess of a reagent — more than enough to react with all of it.
- Add a precisely known, deliberate excess of a reagent to the analyte, so you are sure all the analyte has reacted.
- Give the slow or difficult reaction plenty of time to go to completion — heating or waiting as needed.
- Now titrate the leftover, unreacted excess with a second standardized titrant, to a clear end point.
- Subtract the leftover from the total you added — the difference is the amount that reacted with the analyte.
The logic is simple bookkeeping: total added minus leftover equals what the analyte consumed. You never had to watch the slow reaction reach an end point directly — you only needed to measure how much excess survived it. Back-titration is a quiet reminder that titrimetry is flexible: when a problem will not yield to the obvious approach, a little indirection often opens the door.