Describe the preparation of insoluble salts by precipitation

Acids, Bases and Salts – Preparation of Insoluble Salts by Precipitation

Objective (AO1‑AO3)

• Write balanced molecular and net‑ionic equations for precipitation reactions.
• Predict which product will precipitate using the Cambridge solubility rules.
• Carry out a quantitative preparation of an insoluble salt, calculate theoretical yield, % yield and discuss factors that influence the amount of precipitate obtained.

1. Background – State of Matter & Diffusion

2. Recap of Acid‑Base Reactions (Syllabus 7.1)

These four characteristic reactions are required for the “acids & bases” part of the syllabus and are often used to generate the ions needed for precipitation:

1. Acid + metal → salt + hydrogen gas e.g. 2 HCl + Zn → ZnCl₂ + H₂↑
2. Acid + carbonate (or bicarbonate) → salt + carbon dioxide + water e.g. 2 H₂SO₄ + CaCO₃ → CaSO₄ + CO₂↑ + H₂O
3. Acid + base → salt + water (neutralisation) e.g. HNO₃ + NaOH → NaNO₃ + H₂O
4. Acid + ammonium salt → salt + ammonia + water e.g. HCl + NH₄Cl → NH₄Cl (no reaction – both soluble)

Corresponding indicator colour changes (litmus, phenolphthalein, methyl orange) are also listed in the syllabus but are not needed for the precipitation work.

3. Oxides (Syllabus 7.2)

• Acidic oxides – non‑metal oxides that form acids on reaction with water (e.g. SO₂, CO₂, NO₂).
• Basic oxides – metal oxides that form bases on reaction with water (e.g. CaO, Na₂O, CuO).
• Amphoteric oxides – oxides that can behave as either acid or base (e.g. Al₂O₃, ZnO).

These oxides are sometimes used to generate the required ions (e.g. CaO → Ca²⁺ + 2 OH⁻) before a precipitation step.

4. Ionic Nature & Lattice Structure of the Precipitate

• Reactants are soluble ionic salts that dissociate completely in water: AB(aq) → A⁺(aq) + B⁻(aq)
• The product that precipitates is an ionic solid with a giant lattice (e.g. BaSO₄ consists of a three‑dimensional array of ^Ba²⁺ and ^SO₄²⁻). The large lattice‑energy released makes the overall reaction energetically favourable.

5. Solubility Rules (Cambridge format)

Rule	Implication for Precipitation
All nitrates (NO₃⁻) are soluble.	Never give a precipitate.
All alkali‑metal (Group 1) salts and ammonium (NH₄⁺) salts are soluble.	Never give a precipitate.
All chlorides, bromides and iodides are soluble, except those of Ag⁺, Pb²⁺, Hg₂²⁺.	AgCl, AgBr, AgI, PbCl₂, PbI₂ are insoluble.
Most sulfates (SO₄²⁻) are soluble, except those of Ba²⁺, Pb²⁺, Ca²⁺, Sr²⁺.	BaSO₄, PbSO₄, CaSO₄, SrSO₄ precipitate.
All carbonates (CO₃²⁻), phosphates (PO₄³⁻) and hydroxides (OH⁻) are insoluble, except those of Group 1 and NH₄⁺.	CaCO₃, PbCO₃, Ag₂CO₃, etc. precipitate.

How to use the table: Identify the cation from one reactant and the anion from the other; if the pair appears in the “insoluble” column, a precipitate will form.

6. General Procedure for Preparing an Insoluble Salt (Syllabus 7.3 – Separation & Purification)

1. Write the balanced molecular equation for the reaction of the two soluble salts.
2. Identify the possible products and apply the solubility rules to decide which is insoluble.
3. Write the net‑ionic equation that shows the formation of the precipitate.
4. Prepare the required volumes of the reactant solutions (commonly 0.10 M) in clean, labelled beakers.
5. Mix the solutions while stirring continuously; observe the appearance of a solid.
6. Allow the mixture to stand (or centrifuge) until the precipitate settles.
7. Filtration: Set up a funnel with filter paper, pour the mixture through, and collect the solid on the paper.
8. Washing: Rinse the precipitate with a small amount of cold distilled water to remove adhering ions.
9. Drying: Air‑dry a stable solid; use a drying oven (~100 °C) for hygroscopic products or place in a desiccator.
10. Weigh the dried precipitate and calculate the experimental yield (see Section 7).

7. Yield Calculations (Mole‑Concept Application)

Theoretical mass (from the limiting reactant):

\[ m_{\text{theo}} = n_{\text{lim}} \times M_{\text{precipitate}} \]

where n = moles of limiting reactant (concentration × volume) and M = molar mass of the precipitate.

Percentage yield:

\[ \%\,\text{yield}= \frac{m_{\text{exp}}}{m_{\text{theo}}}\times 100 \]

Worked Example – Barium Sulphate (BaSO₄)

• 0.050 L of 0.10 M BaCl₂ mixed with 0.050 L of 0.10 M Na₂SO₄.
• Moles of BaCl₂ = 0.10 mol L⁻¹ × 0.050 L = 5.0 × 10⁻³ mol.
• Balanced equation: \(\displaystyle \text{BaCl}_{2}(aq) + \text{Na}_{2}\text{SO}_{4}(aq) \rightarrow \text{BaSO}_{4}(s) + 2\text{NaCl}(aq)\) 1 : 1 stoichiometry ⇒ BaCl₂ is the limiting reagent.
• Molar mass of BaSO₄ = 233.4 g mol⁻¹.
• Theoretical mass = 5.0 × 10⁻³ mol × 233.4 g mol⁻¹ = 1.17 g.
• If the dried precipitate weighs 1.05 g, then \(\displaystyle \%\,\text{yield}= \frac{1.05}{1.17}\times100 \approx 90\%.\)

8. Detailed Example 1 – Preparation of Barium Sulphate (BaSO₄)

Reaction: Barium chloride solution + sodium sulphate solution.

Balanced molecular equation:

\[ \text{BaCl}_{2}(aq) + \text{Na}_{2}\text{SO}_{4}(aq) \rightarrow \text{BaSO}_{4}(s) + 2\text{NaCl}(aq) \]

Net‑ionic equation:

\[ \text{Ba}^{2+}(aq) + \text{SO}_{4}^{2-}(aq) \rightarrow \text{BaSO}_{4}(s) \]

Procedure (condensed):

1. Prepare 0.10 M solutions of BaCl₂ and Na₂SO₄.
2. Measure 50 mL of BaCl₂ solution into a clean beaker.
3. Add 50 mL of Na₂SO₄ solution while stirring; a white precipitate forms instantly.
4. Allow the mixture to stand for 5 min, then filter, wash with cold distilled water and dry in a 100 °C oven.
5. Weigh the solid and calculate % yield using the method in Section 7.

9. Detailed Example 2 – Preparation of Silver Chloride (AgCl)

Reaction: Silver nitrate solution + sodium chloride solution.

Balanced molecular equation:

\[ \text{AgNO}_{3}(aq) + \text{NaCl}(aq) \rightarrow \text{AgCl}(s) + \text{NaNO}_{3}(aq) \]

Net‑ionic equation:

\[ \text{Ag}^{+}(aq) + \text{Cl}^{-}(aq) \rightarrow \text{AgCl}(s) \]

Procedure notes:

• Use 0.10 M AgNO₃ and 0.10 M NaCl; the precipitate is a white solid that is photosensitive – keep the beaker covered with aluminium foil.
• Follow the same filtration, washing and drying steps as in Example 1 (dry at 60 °C to avoid decomposition).

10. Factors Affecting the Amount of Precipitate (AO3)

• Concentration of reactants: More ions → larger amount of solid, provided the reaction is not limited by volume.
• Temperature: Most salts are less soluble at lower temperatures; cooling a saturated solution often increases precipitation. The formation of the lattice is usually exothermic, so the mixture may feel slightly warm.
• Common‑ion effect: Adding an ion already present in the solution reduces the solubility of the product (e.g., excess Na₂SO₄ drives BaSO₄ precipitation).
• Stirring / diffusion: Vigorous mixing shortens the diffusion distance between ions, accelerating precipitation.
• Solubility‑product constant (K_sp): Precipitation occurs when the ionic product (IP) exceeds K_sp. \[ \text{IP} = [\text{M}^{n+}][\text{X}^{m-}] \; > \; K_{sp} \;\; \Rightarrow \;\; \text{solid forms} \]

11. Energetics Note (AO1)

Precipitation reactions are generally exothermic because the lattice energy released on solid formation exceeds the energy required to break ion‑water interactions. This explains the slight rise in temperature often observed when the two solutions are mixed.

12. Redox Context (Link to Electrochemistry)

Most precipitation reactions, such as the formation of AgCl, are redox‑neutral: the oxidation states of all elements remain unchanged. This contrasts with electrolysis, where electrical energy forces oxidation‑reduction processes (e.g., 2 Ag⁺ + 2 e⁻ → 2 Ag(s) at the cathode).

13. Common Insoluble Salts (Precipitates) – Periodic Information

Insoluble Salt	Formula	Typical Colour	Metal Ion (Group/Period)
Silver chloride	AgCl	White	Ag⁺ – Group 11, Period 5
Silver bromide	AgBr	Pale yellow	Ag⁺ – Group 11, Period 5
Silver iodide	AgI	Yellow	Ag⁺ – Group 11, Period 5
Barium sulphate	BaSO₄	White	Ba²⁺ – Group 2, Period 6
Calcium carbonate	CaCO₃	White	Ca²⁺ – Group 2, Period 4
Lead(II) iodide	PbI₂	Bright yellow	Pb²⁺ – Group 14 (post‑transition), Period 6

14. Safety Considerations (Syllabus 7.3 – Hazard Awareness)

• Wear safety goggles, lab coat and nitrile gloves at all times.
• Silver nitrate, lead(II) nitrate, lead(II) iodide and other heavy‑metal salts are toxic and can stain skin – handle in a fume cupboard.
• Dispose of waste solutions according to school chemical‑waste procedures; never pour heavy‑metal solutions down the drain.
• Do not heat silver‑containing precipitates over an open flame (risk of photochemical decomposition).
• Keep photosensitive precipitates (AgCl, AgBr, AgI) protected from light.