Salt Preparation from Insoluble Reactants

Step 1: Adding the Insoluble Reactant in Excess

• A measured volume of acid is poured into a beaker and gently warmed (e.g., to $50^\circ\text{C}$) to increase the rate of reaction.
• The insoluble reactant is added one spatula at a time and stirred.
• You must continue adding the solid until it stops dissolving and a visible layer of unreacted powder remains at the bottom of the beaker. This means the reactant is in excess.
• Why add an excess? This ensures that all of the acid has completely reacted and undergone neutralisation.
• Safety rationale: If the acid is not fully neutralised, any unreacted acid left in the mixture would become dangerously concentrated during the later heating stage.

For example, reacting black copper(II) oxide with dilute sulfuric acid produces a blue solution of copper(II) sulfate and water:

$$\text{CuO(s)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{CuSO}_4\text{(aq)} + \text{H}_2\text{O(l)}$$

Step 2: Removing the Excess Reactant via Filtration

• Once the reaction is complete, the unreacted solid must be removed to leave a pure salt solution. We do this using filtration.
• The mixture is poured through a filter funnel lined with filter paper, usually into a conical flask to prevent splashing.
• How it works: Filtration separates mixtures based on particle size. The dissolved salt and water particles are small enough to pass through the microscopic pores in the filter paper, collecting in the flask as a clear filtrate.
• The solid particles of the excess reactant are too large to pass through and remain trapped on the filter paper as the residue.

Step 3: Isolating the Salt via Crystallisation

• The filtrate is poured into an evaporating basin and heated gently, preferably over a water bath rather than a direct Bunsen flame. A water bath provides gentler heating and stops the hot salt solution from spitting.
• The solution is heated to evaporate some of the water, which increases the concentration of the salt.
• Heating must stop when the crystallisation point is reached, meaning the solution has become a saturated solution.
• The Glass Rod Test: To check for saturation, dip a cold glass rod into the hot solution. If small crystals quickly form on the rod when removed, the solution is saturated.
• The basin is then taken off the heat and left in a warm place to cool slowly.
• Why cool slowly? The solubility of the salt decreases as the temperature drops, forcing the solid to crystallise out of the solution. Cooling slowly gives the particles more time to arrange themselves into a regular, ordered lattice structure, resulting in much larger, well-defined crystals.
• Finally, the remaining liquid is filtered off, and the crystals are gently patted dry between sheets of filter paper.

Why Not Heat to Dryness?

You must never evaporate all the water off with direct heat. Stopping at the crystallisation point is essential because heating to completely dry crystals can cause:

1. Thermal decomposition of the salt (breaking it down into dangerous gases).
2. A safety hazard due to the hot solid spitting out of the basin.
3. The formation of a powdery, anhydrous salt instead of the desired hydrated crystals (e.g., turning blue copper(II) sulfate crystals into a white powder).