What Is Titration? A Simple Guide to Measuring Concentration

In a typical acid base titration, one solution is measured into a flask and another is added slowly from a burette. The burette lets you measure how much solution has been delivered. If the concentration of the burette solution is known, the volume needed to react with the sample gives information about how many moles were present in the unknown solution.

The central idea in titration is stoichiometry. The balanced equation tells you the reacting ratio between the substances. For example, if one mole of an acid reacts with one mole of a base, reaching the correct reaction point means the moles added from the burette match the moles required by that ratio. Different reactions can have different mole ratios.

Titration can be used for more than acids and bases. Other versions rely on oxidation and reduction reactions, complex formation, or precipitation. The exact chemicals and endpoint method change, but the reasoning stays similar: add a measured amount of a known reagent until it has reacted in a predictable proportion with the substance being measured.

Many titrations need a visible or measurable sign that enough reagent has been added. An acid base indicator changes colour over a particular pH range. You choose an indicator whose colour change occurs close to the equivalence point for the reaction being studied. Electronic pH meters can also track the change and help locate the endpoint more precisely.

When practising titration, it helps to separate two related terms. The equivalence point is the theoretical point where reactants have been mixed in the exact stoichiometric amounts required by the equation. The endpoint is the observed signal, such as a colour change. A well-chosen method makes those two points very close, but they are not identical definitions.

A reliable titration is easier when every measurement is recorded clearly. A common school laboratory routine is:

• •Rinse and fill the burette with the solution of known concentration, then record the starting reading.
• •Measure a known volume of the unknown solution into a flask and add a suitable indicator if one is needed.
• •Add the burette solution while swirling, then slow to dropwise addition as the endpoint approaches.
• •Record the final burette reading and calculate the volume delivered.
• •Repeat until you have closely agreeing results rather than relying on one rushed trial.

After the experiment, the calculation usually follows a steady path. Use concentration and volume to find moles of the known solution, apply the mole ratio from the balanced equation, then use the sample volume to calculate the unknown concentration. When learning about titration, writing the units at every stage is one of the best error checks. Litres, moles, and moles per litre should connect cleanly.

Here is a typical calculation. Suppose 25.0 cubic centimetres of sodium hydroxide solution needs 20.0 cubic centimetres of hydrochloric acid, concentration 0.100 moles per cubic decimetre, to reach the endpoint. First convert the acid volume to cubic decimetres, giving 0.0200. Moles of acid equal 0.100 times 0.0200, which is 0.00200. The equation shows one mole of acid reacting with one mole of alkali, so there are also 0.00200 moles of sodium hydroxide. Dividing by its volume, 0.0250 cubic decimetres, gives a concentration of 0.0800 moles per cubic decimetre.

Indicator choice depends on the pH at the equivalence point, which depends on the strengths of the acid and base. A strong acid with a strong base gives a solution close to pH 7, and the pH jumps steeply near the endpoint, so several indicators work. A weak acid such as ethanoic acid titrated with sodium hydroxide has an equivalence point above 7, so phenolphthalein, which changes colour at roughly pH 8 to 10, is suitable, while methyl orange, changing at roughly pH 3 to 4.5, would signal far too early. Plotting pH against volume produces a titration curve showing this.

Most titration errors come from reading and handling the burette. Read the scale at the bottom of the curved liquid surface, called the meniscus, with your eye level with it, because looking from above or below shifts the reading. Before starting, open the tap briefly so the narrow tip below it is full of liquid, since an air bubble there will empty out during the titration and add a false volume. Removing the funnel used for filling also matters, as a drop clinging to it can fall in later.

A few habits make the endpoint sharper. Place the flask on a white tile so a faint colour change is easier to see, and rinse droplets from the inside walls of the flask with distilled water near the end. Adding water in this way does not change the result, because the number of moles in the flask stays the same. Many students run a quick rough titration first to find the approximate volume, then perform careful runs and use only results that agree closely, often within 0.10 cubic centimetres of each other.

Some samples cannot be titrated directly. Eggshell, chalk and many indigestion tablets contain calcium carbonate, which does not dissolve in water and reacts too slowly with acid to give a sharp endpoint. Chemists get round this with a back titration. You add a known excess of acid, let the solid react completely, then titrate the acid that is left over with an alkali. The difference between the acid you added and the acid remaining tells you how much the sample used up, and from that how much carbonate it contained.

Suppose 1.60 grams of crushed eggshell reacts with 50.0 cubic centimetres of hydrochloric acid of concentration 1.00 moles per cubic decimetre, which is 0.0500 moles of acid. The leftover acid needs 20.0 cubic centimetres of 1.00 moles per cubic decimetre sodium hydroxide, so 0.0200 moles remained and 0.0300 moles reacted with the shell. One mole of calcium carbonate reacts with two moles of acid, giving 0.0150 moles of carbonate. With a molar mass of about 100 grams per mole, that is 1.50 grams, so the shell is roughly 94 percent calcium carbonate.

Some titrations signal their own endpoint. Potassium manganate(VII), often called potassium permanganate, forms an intensely purple solution, but when it reacts with iron(II) ions in acid it becomes manganese(II), which is almost colourless. During the titration each drop loses its colour as it reacts. Once all the iron(II) has been used up, the next drop leaves a faint pink tint that does not fade, and that is the endpoint. The reaction is done in dilute sulfuric acid rather than hydrochloric acid, because manganate(VII) would also oxidise chloride ions, using up extra solution and giving too large a reading.

This method is used to check the iron content of supplement tablets. The balanced equation shows that one manganate(VII) ion reacts with five iron(II) ions. Suppose a dissolved tablet needs 23.30 cubic centimetres of 0.0100 moles per cubic decimetre potassium manganate(VII). That is 0.0100 multiplied by 0.02330, or 0.000233 moles of manganate(VII). Multiplying by five gives 0.001165 moles of iron(II). Iron has a molar mass of about 55.8 grams per mole, so the tablet contains about 0.0650 grams, or 65 milligrams, of iron, a typical amount for a common iron supplement.

Titration is a quantitative method that uses a known solution, a measured reaction volume, and stoichiometry to work out an unknown concentration. Focus on three pieces: accurate volume readings, the correct reaction ratio, and a dependable endpoint. Once those are clear, the calculation becomes a logical extension of the experiment rather than a formula to memorise.