Preparing a Soluble Salt by Titration

Introduction

Sodium chloride, potassium nitrate and ammonium sulfate are all soluble salts, and so are the bases that contain their metal or ammonium ion: sodium hydroxide, potassium hydroxide and ammonia solution all dissolve in water. That creates a problem. We cannot add an excess of a soluble alkali and filter it off, because the excess would stay dissolved. The answer is titration : measuring the exact volumes of acid and alkali that just neutralise each other, then mixing those volumes again to give a solution of pure salt.

Core explanation

The problem with soluble reactants. When we made salts from insoluble bases, we added excess base and filtered it away. With a soluble alkali, any excess dissolves and cannot be separated from the salt solution. Equally, excess acid would remain in solution. So we need to mix acid and alkali in exactly the right proportions — no more of either than is needed.

The reaction. An acid and an alkali react to form a salt and water. For sodium chloride:

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

The underlying ionic reaction is always H⁺(aq) + OH⁻(aq) → H₂O(l).

Stage one: find the reacting volumes. A fixed, measured volume of alkali is placed in a conical flask using a pipette, and a few drops of indicator are added. Acid is added from a burette, slowly and with swirling, until the indicator just changes colour. This is the end point . The volume of acid added is read from the burette. The titration is repeated until two results agree closely, giving a reliable volume.

Stage two: make the salt without indicator. The indicator is a coloured impurity that would contaminate the crystals. So the same volume of alkali is measured into a clean flask, and exactly the volume of acid found in stage one is added — this time with no indicator. The resulting solution contains only the salt and water. (An alternative is to add activated charcoal to absorb the indicator and filter it off, but repeating without indicator is simpler.)

Stage three: crystallise. As with other soluble salts, the solution is warmed to evaporate some of the water until it is saturated, then left to cool so that crystals form. The crystals are separated and dried.

Which salts use this method? Salts of sodium, potassium and ammonium, because their bases (hydroxides, carbonates or ammonia) are soluble. For example:

- KOH(aq) + HNO₃(aq) → KNO₃(aq) + H₂O(l) - 2NH₃(aq) + H₂SO₄(aq) → (NH₄)₂SO₄(aq) - Na₂CO₃(aq) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g)

Accuracy matters. Near the end point a single drop can change the colour, so acid is added drop by drop and the flask is swirled. Reading the burette at eye level, at the bottom of the meniscus, gives accurate volumes.

Step-by-step reasoning

To choose the right preparation method for a soluble salt:

1. Is the salt soluble? If not, use precipitation instead. 2. Is the base containing the metal soluble (sodium, potassium or ammonium)? 3. If yes, use titration to find exact volumes, then repeat without indicator. 4. If no, add excess insoluble base and filter.

Visual explanation

Imagine a tall burette of acid clamped above a conical flask of pink alkali containing phenolphthalein on a white tile. Drop by drop the acid falls; the pink flashes colourless where each drop lands, then returns on swirling — until one drop turns the whole flask colourless for good.

Real-world analogy

Titration is like adding exactly enough milk to cool a cup of coffee to drinking temperature: first you test carefully, noting how much you used; next time you pour that measured amount straight in, without needing to keep checking.

Real-world example

Ammonium sulfate and potassium nitrate are important fertilisers. On an industrial scale they are made by neutralising sulfuric acid with ammonia and nitric acid with a potassium base, with the proportions controlled by continuous pH monitoring — the large-scale version of reaching an end point.

Why?

Why is the titration repeated without indicator? The indicator is itself a chemical. If it were left in the flask it would be trapped in or on the crystals, making the salt impure and often coloured. Using the known volumes lets us reach neutrality without it.

Common misconception

"The end point means the solution is always exactly pH 7." The indicator changes colour at the equivalence point for a strong acid and strong alkali, which is close to pH 7, but for weak acids or weak bases the salt solution at neutralisation may be slightly acidic or alkaline.

Worked example

Question: In a titration, 25.0 cm³ of potassium hydroxide solution needed 21.40 cm³ of dilute nitric acid. How would you now make pure potassium nitrate crystals?

Reasoning: The volumes react exactly. Mixing them again without indicator gives only KNO₃ and water: KOH(aq) + HNO₃(aq) → KNO₃(aq) + H₂O(l).

Answer: Pipette 25.0 cm³ of the alkali into a clean flask, add exactly 21.40 cm³ of the acid with no indicator, then concentrate the solution by warming and leave it to cool and crystallise.

Quick check

1. Why can excess sodium hydroxide not be used and filtered off when making sodium chloride? Answer: Sodium hydroxide is soluble, so the excess would stay dissolved and could not be removed by filtration.

Exam focus

Be ready to explain why titration is used for sodium, potassium and ammonium salts, to name the burette, pipette and indicator, and to state that the titration is repeated without indicator before crystallising. Give ionic and full equations.

Advanced insight

The end point (when the indicator changes) and the equivalence point (when the amounts of acid and alkali match exactly from the equation) are not quite the same thing. A well-chosen indicator changes colour in the steep part of the pH curve, so the two points differ by less than a drop. A pH meter can locate the equivalence point directly without any indicator.

Summary

When both acid and alkali are soluble, excess cannot be filtered off, so titration is used. A first titration with indicator finds the exact reacting volumes; these volumes are then mixed without indicator to give a pure salt solution, which is concentrated and cooled to crystallise. This method makes sodium, potassium and ammonium salts.

Practice questions

1. Name two pieces of apparatus used to measure volumes accurately in a titration. Answer: A burette and a pipette. 2. Which salts are usually made by titration, and why? Answer: Sodium, potassium and ammonium salts, because their bases are soluble and excess could not be filtered off. 3. Write the ionic equation for any acid–alkali neutralisation. Answer: H⁺(aq) + OH⁻(aq) → H₂O(l) 4. Write a balanced equation for making ammonium nitrate from ammonia solution and nitric acid. Answer: NH₃(aq) + HNO₃(aq) → NH₄NO₃(aq) 5. Why should acid be added drop by drop near the end point? Answer: A single drop can change the indicator colour, so adding slowly avoids overshooting the exact reacting volume.