Salts — preparation, properties & uses

Salts — preparation, properties & uses

{{TABLE: title=Classification of Salts at a Glance

What is a Salt?

In everyday language, "salt" usually means table salt, the white crystalline solid we add to our food. In chemistry, however, the term has a much broader meaning. A salt is an ionic compound that results from the neutralisation reaction between an acid and a base. It is made up of a positive ion (cation) from a base and a negative ion (anion) from an acid.

For example, when hydrochloric acid (HCl) reacts with sodium hydroxide (NaOH), they neutralise each other to form sodium chloride (NaCl) and water (H₂O). The Na⁺ is the cation from the base (NaOH), and the Cl⁻ is the anion from theacid (HCl). Together, they form the salt, NaCl. This is the fundamental principle behind the formation of all salts.

{{FORMULA: expr=Acid + Base → Salt + Water | symbols=This is the general equation for a neutralisation reaction, the primary way salts are formed.}}

{{KEY: type=definition | title=Salt | text=A salt is an ionic compound formed when the hydrogen ion (H⁺) of an acid is partially or completely replaced by a metal ion or an ammonium ion (NH₄⁺).}}

Preparation of Salts: A Method for Every Salt

The method used to prepare a specific salt depends critically on one key property: its solubility in water. We can't use the same method to make soluble table salt (NaCl) as we do to make the insoluble pigment lead(II) chromate (PbCrO₄). Therefore, we divide preparation methods into two main categories.

1. Preparation of Soluble Salts

Soluble salts can be prepared by methods that leave the salt dissolved in water, from which it can then be recovered by crystallisation. There are four main methods.

A) Acid + Reactive Metal

This method is suitable for preparing salts of moderately reactive metals like magnesium, zinc, iron, and aluminium. It is not suitable for very reactive metals (like sodium or potassium, which react explosively with acids) or unreactive metals (like copper, silver, or gold, which do not react with dilute acids).

1. Reaction: Add the reactive metal (e.g., zinc powder) in excess to a beaker of a suitable dilute acid (e.g., dilute sulphuric acid). The excess metal ensures all the acid is used up. Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
2. Observation: You will see bubbling or effervescence as hydrogen gas is produced. The reaction stops when all the acid has been consumed.
3. Filtration: Filter the mixture to remove the unreacted excess metal. The filtrate is a pure solution of the salt (zinc sulphate).
4. Crystallisation: Gently heat the filtrate in an evaporating dish to evaporate some of the water and create a saturated solution. To check if it's saturated, dip a cool glass rod into the solution; if small crystals form on it, the solution is ready.
5. Cooling: Allow the saturated solution to cool slowly at room temperature. As it cools, the solubility of the salt decreases, and beautiful, pure crystals will form.
6. Drying: Decant the remaining liquid and dry the crystals by pressing them gently between sheets of filter paper.

B) Acid + Insoluble Base or Carbonate

This is a very common and safe method. It uses an insoluble base (a metal oxide or hydroxide) or an insoluble carbonate. The process is very similar to the acid-metal reaction.

1. Reaction: Add the insoluble base (e.g., copper(II) oxide) or carbonate (e.g., calcium carbonate) in excess to a warm dilute acid (e.g., dilute sulphuric acid). Warming the acid speeds up the reaction. CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l) CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)
2. Filtration & Crystallisation: Follow the same steps of filtration, heating to saturation, cooling, and drying as described above to obtain pure crystals of the salt (e.g., copper(II) sulphate).

C) Acid + Alkali (Titration)

This method is essential when both the acid and the base are soluble in water. Since both are colourless solutions, we can't tell when the neutralisation is complete just by looking. We must use a technique called titration, which involves an indicator. This is used to make salts of sodium, potassium, and ammonium.

1. Setup: Pipette a known volume (e.g., 25 cm³) of the alkali (e.g., sodium hydroxide) into a conical flask. Add a few drops of a suitable indicator (e.g., methyl orange or phenolphthalein).
2. Titration: Fill a burette with the acid (e.g., hydrochloric acid) and record the initial reading. Slowly add the acid to the flask, swirling continuously, until the indicator changes colour permanently. This is the end point. Record the final burette reading to find the volume of acid used.
3. Pure Salt Preparation: Now that you know the exact volumes needed for neutralisation, repeat the experiment with the same volumes of acid and alkali but without the indicator. The indicator is an impurity we don't want in our final salt.
4. Crystallisation: The resulting solution is a pure solution of the salt (e.g., sodium chloride). Obtain the crystals by following the standard crystallisation procedure (heating, cooling, drying).

{{VISUAL: diagram: A labelled laboratory setup for acid-alkali titration, showing a burette containing acid dripping into a conical flask containing alkali and an indicator. The flask is being swirled.}}

{{KEY: type=exam | title=Why is Titration Necessary? | text=For reactions involving a soluble acid and a soluble base (alkali), titration is the only way to determine the exact volumes required for complete neutralisation. Simply mixing them would result in a solution contaminated with excess acid or alkali.}}

2. Preparation of Insoluble Salts

Insoluble salts cannot be recovered by crystallisation from a solution. Instead, they are prepared by a precipitation reaction, also known as double decomposition. This involves mixing two different soluble salt solutions.

1. Principle: When two soluble ionic compounds are mixed, their ions can 'swap partners'. If one of the new combinations of ions forms an insoluble compound, it will instantly appear as a solid precipitate. Soluble Salt A + Soluble Salt B → Insoluble Salt C (precipitate) + Soluble Salt D
2. Example: To prepare insoluble silver chloride (AgCl), we can mix solutions of two soluble salts: silver nitrate (AgNO₃) and sodium chloride (NaCl). AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)
3. Procedure:
   • Mix the two solutions in a beaker. An immediate white solid (precipitate) of silver chloride will form.
   • Filter the mixture to separate the solid precipitate from the solution.
   • Wash the precipitate on the filter paper with a small amount of distilled water to remove any remaining soluble impurities (like NaNO₃).
   • Dry the precipitate in a warm oven or a desiccator.

{{VISUAL: diagram: Two beakers, one with a clear solution labelled 'Silver Nitrate (aq)' and another with a clear solution labelled 'Sodium Chloride (aq)', are shown being poured into a larger beaker. In the larger beaker, a white solid precipitate is shown forming at the bottom, with an arrow pointing to it labelled 'Silver Chloride (s)'.}}

Properties of Salts

Salts exhibit a wide range of properties, which determine their behaviour and uses.

Physical Properties

1. State: Most salts are crystalline solids at room temperature due to the strong electrostatic forces of attraction between the positive and negative ions in their giant ionic lattice structure.
2. Melting and Boiling Points: They generally have high melting and boiling points because a large amount of thermal energy is required to overcome these strong electrostatic forces.
3. Electrical Conductivity:
   • In the solid state, salts do not conduct electricity because their ions are held in fixed positions and are not free to move.
   • In the molten (fused) state or when dissolved in water (aqueous solution), salts are good conductors of electricity. In these states, the ions are mobile and can move to carry an electric charge.
4. Solubility: The solubility of salts in water varies greatly. Some are highly soluble (like NaCl), while others are almost completely insoluble (like AgCl). There are general rules to predict solubility.

{{TABLE: title=General Solubility Rules for Salts in Water

Chemical Properties

1. Salt Hydrolysis

When some salts dissolve in water, they react with the water in a process called hydrolysis, which can make the resulting solution acidic, basic, or neutral. This depends on the strength of the acid and base from which the salt was formed.

• Salt of Strong Acid + Strong Base: (e.g., NaCl from HCl and NaOH). The solution is neutral (pH ≈ 7) because neither ion reacts significantly with water.
• Salt of Strong Acid + Weak Base: (e.g., NH₄Cl from HCl and NH₄OH). The solution is acidic (pH < 7) because the cation from the weak base (NH₄⁺) reacts with water to produce H⁺ ions.
• Salt of Weak Acid + Strong Base: (e.g., CH₃COONa from CH₃COOH and NaOH). The solution is basic/alkaline (pH > 7) because the anion from the weak acid (CH₃COO⁻) reacts with water to produce OH⁻ ions.
• Salt of Weak Acid + Weak Base: (e.g., NH₄CH₃COO). The solution can be acidic, basic, or neutral, depending on the relative strengths of the weak acid and weak base.

{{ZOOM: title=What is 'Strength' in Acids and Bases? | text=A 'strong' acid or base is one that fully ionises or dissociates in water, releasing all of its H⁺ or OH⁻ ions (e.g., HCl, NaOH). A 'weak' acid or base only partially ionises, meaning an equilibrium is established and only a small fraction of the ions are released (e.g., CH₃COOH, NH₄OH).}}

2. Effect of Heat (Thermal Decomposition)

Many salts decompose when heated strongly. The products of decomposition depend on the specific salt.

• Carbonates: Most metal carbonates decompose on heating to form a metal oxide and carbon dioxide. Carbonates of very reactive metals (Na, K) are thermally stable. CuCO₃(s) → CuO(s) + CO₂(g)
• Nitrates: The decomposition of nitrates is more complex and depends on the reactivity of the metal.
  • Nitrates of Na, K → Metal Nitrite + Oxygen
  • Nitrates of Ca, Mg, Zn, Fe, Cu → Metal Oxide + Nitrogen Dioxide + Oxygen
  • Nitrates of Ag, Hg → Metal + Nitrogen Dioxide + Oxygen
• Sulphates: Most sulphates are very stable, but some, like copper(II) sulphate, decompose at very high temperatures. Hydrated salts, like CuSO₄.5H₂O, first lose their water of crystallisation on gentle heating. CuSO₄.5H₂O(s) [blue] ⇌ CuSO₄(s) [white] + 5H₂O(g)

Important Uses of Common Salts

Salts are not just laboratory chemicals; they are vital components of industry, agriculture, medicine, and our daily lives.

{{TABLE: title=Everyday and Industrial Uses of Common Salts

From the salt on our tables to the minerals in our phones, salts are the quiet, essential compounds that build and run our world. Understanding their preparation and properties is fundamental to the study of chemistry.

{{FLASHCARD: q=What are the two main classes of methods for preparing salts, and what property determines which method to use? | a=The two main classes are methods for preparing soluble salts (e.g., acid + metal, titration) and methods for preparing insoluble salts (precipitation). The determining property is the salt's solubility in water.}}