Chemistry (0620) | Preparation of salts

Chemistry (0620) Study Notes: 7.3 Preparation of Salts

Hello future chemists! This chapter is incredibly important because it brings together everything you've learned about acids, bases, and reactions. Salts are everywhere—from table salt (\(NaCl\)) to fertilisers, medicines, and antacids.
We are going to master the three main techniques used in the lab to make salts, depending on whether the salt is soluble (dissolves in water) or insoluble (does not dissolve in water).

Section 1: Solubility Rules – The Most Important Prerequisite!

Before you can prepare a salt, you must first know if it dissolves in water. This determines which preparation method you must use. Don't worry, the rules are simpler than they look!

Key Concept: General Solubility Rules

A salt is either soluble (forms a clear solution) or insoluble (forms a precipitate).

Rule 1: Always Soluble Salts (The "Always Wins" Group)

• All salts containing Sodium (\(Na^{+}\)) are soluble.
• All salts containing Potassium (\(K^{+}\)) are soluble.
• All salts containing Ammonium (\(NH_{4}^{+}\)) are soluble.
• All Nitrates (\(NO_{3}^{-}\)) are soluble.

Memory Aid (Mnemonic): Think of SNAP or NASH (Nitrates, Ammonium, Sodium, Hydroxides/Potassium). These are your soluble heroes!

Rule 2: Soluble with Exceptions

• All Chlorides (\(Cl^{-}\)) are soluble, except Lead(II) chloride (\(PbCl_{2}\)) and Silver chloride (\(AgCl\)).
• All Sulfates (\(SO_{4}^{2-}\)) are soluble, except Barium sulfate (\(BaSO_{4}\)), Calcium sulfate (\(CaSO_{4}\)), and Lead(II) sulfate (\(PbSO_{4}\)).

Rule 3: Mostly Insoluble Salts

• All Carbonates (\(CO_{3}^{2-}\)) are insoluble, except those of Sodium, Potassium, and Ammonium (Rule 1).
• Most Hydroxides are insoluble, except those of Sodium, Potassium, and Ammonium (Rule 1), and Calcium hydroxide (which is partially soluble).

Section 2: Preparation of Soluble Salts (Two Main Methods)

There are two methods for preparing soluble salts. The choice depends entirely on the nature of the reactants you are using.

We use these methods when the reactants are: Metal + Acid, Base + Acid, or Carbonate + Acid.

Method 2: Reaction of an Acid with Excess Insoluble Solid

This method is used when one reactant is soluble (the acid) and the other is insoluble (metal, base, or carbonate). By adding the insoluble solid in excess, you guarantee that all the acid is used up.

Example: Preparing Copper(II) sulfate (\(CuSO_{4}\) - soluble) using Copper(II) oxide (insoluble base) and Sulfuric acid.

Step 1: Reaction (Neutralisation)

Heat the dilute acid (Sulfuric acid) and slowly add the excess insoluble solid (Copper(II) oxide) while stirring. You know the acid is fully used up when the excess solid no longer dissolves and remains at the bottom of the beaker.
\(H_{2}SO_{4}(aq) + CuO(s) \rightarrow CuSO_{4}(aq) + H_{2}O(l)\)

Step 2: Separation (Filtration)

We need a pure salt solution. Use filtration to remove the undissolved excess Copper(II) oxide solid (the residue). The filtrate is the pure Copper(II) sulfate solution.

Step 3: Purification (Evaporation and Crystallisation)

1. Heat (Evaporation): Gently heat the filtrate (solution) to evaporate most of the water until the solution is saturated (meaning no more salt can dissolve). How do you test for saturation? Dip a cool glass rod into the solution—if tiny crystals form immediately on the rod, it's ready!
2. Cooling (Crystallisation): Leave the hot saturated solution to cool slowly, usually in an evaporating dish or beaker. Slow cooling allows large, pure crystals to form.
3. Drying: Filter the crystals to separate them from the remaining solution (mother liquor). Wash them briefly with a small amount of cold distilled water and dry them (e.g., between filter papers or in a desiccator).

Method 1: Titration (For Soluble Salt involving an Alkali)

This method is essential when both reactants are soluble (e.g., an acid reacting with a soluble base—an alkali). Since both reactants are colourless liquids, we cannot use the "excess solid" technique, as we wouldn't know when the reaction stops. We must measure the exact amounts needed.

Example: Preparing Sodium chloride (\(NaCl\) - soluble) using Hydrochloric acid and Sodium hydroxide (alkali).

Step 1: Standard Titration (Finding the End Point)

1. Measure a known volume of the alkali (e.g., \(25 cm^3\) Sodium hydroxide) using a volumetric pipette and place it in a conical flask.
2. Add a few drops of a suitable indicator (e.g., phenolphthalein or methyl orange).
3. Fill a burette with the acid (Hydrochloric acid) and record the initial reading.
4. Add the acid dropwise to the flask until the indicator changes colour permanently. This is the end-point. Record the final burette reading.
5. Calculate the exact volume of acid needed for complete neutralisation (called the titre).

Step 2: Preparation (Making the Pure Salt)

1. Repeat the experiment, but this time, do not add the indicator.
2. Mix the exact volumes of acid and alkali (using the titre volume you found in Step 1) to ensure perfect neutralisation.
3. The resulting solution is pure aqueous Sodium chloride.
4. Crystallisation: Use the evaporation and slow cooling process (as described in Method 2, Step 3) to obtain dry crystals.

Section 3: Hydrated Salts and Water of Crystallisation (Core & Supplement)

Many salts form crystals that incorporate water molecules into their structure. This water is called water of crystallisation.

• Hydrated substance: A substance that is chemically combined with water (it contains water of crystallisation).
• Anhydrous substance: A substance containing no water. (You make this by heating the hydrated salt).
• Water of Crystallisation: The water molecules present in hydrated crystals.

Example: Copper(II) sulfate.

The blue crystals you typically see are hydrated copper(II) sulfate, formula: \(CuSO_{4} \cdot 5H_{2}O\). The ' \(\cdot 5H_{2}O\) ' means five water molecules are attached to every one copper(II) sulfate molecule.

When heated strongly, this blue salt loses its water of crystallisation and becomes anhydrous copper(II) sulfate, which is a white powder:

\(CuSO_{4} \cdot 5H_{2}O(s) \rightarrow CuSO_{4}(s) + 5H_{2}O(g)\)

Section 4: Preparation of Insoluble Salts (Supplement Content)

How do we make a salt that doesn't dissolve in water? We use the precipitation method.

Method 3: Precipitation (Double Decomposition)

This method involves mixing two different soluble salts together in solution. The ions swap partners, and if one of the new combinations forms an insoluble salt (a precipitate), it crashes out of the solution as a solid.

Example: Preparing Lead(II) sulfate (\(PbSO_{4}\) - insoluble)

According to the solubility rules, Lead salts are insoluble (exception to the sulfate rule). We must start with two soluble compounds that contain the necessary ions: Lead ions (\(Pb^{2+}\)) and Sulfate ions (\(SO_{4}^{2-}\)).

We choose: Soluble Lead salt (e.g., Lead(II) nitrate) and Soluble Sulfate salt (e.g., Sodium sulfate).

\(Pb(NO_{3})_{2}(aq) + Na_{2}SO_{4}(aq) \rightarrow PbSO_{4}(s) + 2NaNO_{3}(aq)\)

The ionic equation shows only the ions involved in forming the precipitate:

\(Pb^{2+}(aq) + SO_{4}^{2-}(aq) \rightarrow PbSO_{4}(s)\)

Step-by-Step Practical Procedure

1. Mix: Dissolve the two soluble salts in separate beakers of distilled water. Mix the two solutions together. An immediate solid precipitate forms.
2. Separate: Use filtration to separate the insoluble precipitate (\(PbSO_{4}\)) from the solution containing the soluble by-product (\(NaNO_{3}\)) and any excess reactants.
3. Purify: Wash the precipitate remaining on the filter paper thoroughly with cold distilled water to remove all traces of the soluble contaminants.
4. Dry: Scrape the washed precipitate onto fresh filter paper or a watch glass and leave it to dry in a warm oven or desiccator.