Volumetric Analysis: Titrations of Every Kind

Most students meet titration as a way of neutralising an acid with an alkali. But acid–base titration is just one member of a large family of techniques called volumetric analysis, in which the amount of a substance is found by measuring the volume of a solution of known concentration needed to react with it completely. The same basic method measures vitamin C in orange juice, iron in tablets, calcium in hard water and chloride in seawater.

This article is a map of that family.

The common principle

Every titration has the same ingredients:

1. A standard solution (the titrant) of accurately known concentration, usually in a burette.
2. A measured amount of the analyte, usually pipetted into a conical flask.
3. A reaction between them that is fast, complete and of known stoichiometry.
4. A way of detecting the end point, when the reaction is just complete: an indicator colour change, or an instrumental signal.

The calculation is always the same three steps:

1. Moles of titrant = concentration × volume.
2. Moles of analyte from the balanced equation’s mole ratio.
3. Concentration, mass or percentage of analyte.

See titration calculations for the full method.

The four main types

Type	Reaction	Typical titrant	Typical end point detection	Example use
Acid–base	H⁺ + OH⁻ → H₂O	NaOH, HCl	pH indicators (phenolphthalein, methyl orange) or pH meter	acid in vinegar, antacid capacity
Redox	electron transfer	KMnO₄, K₂Cr₂O₇, I₂, Na₂S₂O₃	self-indicating (permanganate), starch, redox indicators	iron in tablets, vitamin C, chlorine in bleach
Complexometric	metal ion + ligand → complex	EDTA	metal ion indicators (Eriochrome Black T)	water hardness, calcium and magnesium
Precipitation	ions → insoluble solid	AgNO₃	chromate, adsorption indicators, or Fe³⁺/SCN⁻	chloride in water or food

1. Acid–base titrations

The most familiar type. The reaction is neutralisation, and the end point is detected by an indicator that changes colour in the steep part of the pH curve.

Example: 25.00 cm³ of vinegar (diluted tenfold) titrated with 0.100 mol/dm³ NaOH using phenolphthalein.

More in acid–base titration, choosing an indicator and finding the acetic acid content of vinegar.

2. Redox titrations

The analyte is oxidised or reduced by the titrant. Common systems:

• Potassium manganate(VII) (permanganate), KMnO₄: a strong oxidising agent, used in acidic solution. It’s self-indicating: its deep purple MnO₄⁻ ions are reduced to almost colourless Mn²⁺, so the end point is the first permanent pale pink from a slight excess of permanganate. MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺
• Potassium dichromate(VI), K₂Cr₂O₇: orange Cr₂O₇²⁻ is reduced to green Cr³⁺. A redox indicator marks the end point.
• Iodine–thiosulfate: iodine is released by an oxidising agent and then titrated with sodium thiosulfate. Starch is added near the end point; the blue-black starch–iodine colour disappears sharply when all the iodine has reacted. I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻

Example: measuring iron(II) in an iron supplement tablet with potassium permanganate. The 1 : 5 mole ratio (MnO₄⁻ : Fe²⁺) is crucial in the calculation.

3. Complexometric titrations

A metal ion reacts with a ligand to form a soluble complex. The most important titrant is EDTA (ethylenediaminetetraacetic acid), which forms very stable 1 : 1 complexes with most metal ions.

Example: measuring water hardness by titrating calcium and magnesium ions with EDTA at pH 10, using Eriochrome Black T indicator (wine red to blue).

More in EDTA titrations and water hardness.

4. Precipitation titrations

The titrant and analyte form an insoluble precipitate. The main example is argentometry: titration with silver nitrate to measure halide ions.

Ag⁺ + Cl⁻ → AgCl(s)

Example: measuring the chloride content of drinking water or seawater by Mohr’s method, using potassium chromate as the indicator.

More in precipitation titrations.

Direct vs back titrations

• In a direct titration, the titrant reacts directly with the analyte in the flask.
• In a back titration, a known excess of reagent is added to the analyte, and the unreacted excess is then titrated. The amount that reacted with the analyte is found by subtraction.

Back titrations are used when:

• the analyte is an insoluble solid (such as an antacid tablet or limestone)
• the reaction with the analyte is slow
• a direct end point is hard to see
• the analyte is volatile (such as ammonia)

See back titration and the antacid titration practical.

Primary standards and standardisation

A titration is only as accurate as the concentration of the titrant. The most reliable titrants are made from a primary standard, a substance that:

• is available in a very pure form
• is stable in air (doesn’t absorb water or CO₂, and doesn’t oxidise)
• has a known, fixed formula
• is soluble in water
• ideally has a high molar mass (so weighing errors are small)

Titrant	Can it be a primary standard?	Standardised against
Sodium hydroxide	no (absorbs water and CO₂)	potassium hydrogenphthalate (KHP)
Hydrochloric acid	no (concentration of stock varies)	anhydrous sodium carbonate
Potassium permanganate	no (reacts slowly with impurities; light-sensitive)	sodium ethanedioate (sodium oxalate)
Sodium thiosulfate	no (slowly decomposes)	potassium iodate
Potassium dichromate	yes
EDTA (disodium salt)	often used directly, or standardised against CaCO₃ or zinc

Standardisation means titrating a solution against a primary standard to find its exact concentration before using it to analyse unknowns.

Instrumental end points

Not every titration has a convenient colour change. The end point can also be found by:

• pH meter (potentiometric titration): steepest point on a pH curve. See plotting a pH curve.
• Electrode potential: for redox and precipitation titrations.
• Conductivity: the minimum or change in slope of a conductivity graph.
• Temperature: the peak of a thermometric titration.

These methods work for coloured or cloudy solutions and remove the subjectivity of judging a colour.

Choosing the right type

If the analyte is…	Consider
an acid or base	acid–base titration
a species that can be oxidised or reduced (Fe²⁺, H₂O₂, vitamin C, I₂)	redox titration
a metal ion that forms complexes (Ca²⁺, Mg²⁺, Zn²⁺)	EDTA titration
a halide or other ion forming an insoluble silver salt	precipitation titration
an insoluble solid or slow-reacting substance	back titration

A short history

Volumetric analysis grew up in the late 18th and 19th centuries alongside industrial chemistry. The French chemist François-Antoine-Henri Descroizilles designed an early burette in the 1790s to test the strength of bleach, and Joseph Louis Gay-Lussac refined the burette and the method in the 1820s and 1830s. Karl Friedrich Mohr added the tap-controlled burette and the volumetric pipette in the mid-1800s, much like the ones used today. Titration became the workhorse of industrial quality control because it was fast, cheap and accurate.

Key takeaways

• Volumetric analysis measures an amount from the volume of a standard solution needed for complete reaction.
• The four main types are acid–base, redox, complexometric (EDTA) and precipitation titrations.
• Back titrations use a known excess and subtraction, for insoluble or slow-reacting analytes.
• Titrants are standardised against primary standards: pure, stable substances of known formula.
• End points can be found with indicators or instruments (pH, potential, conductivity, temperature).