O-Level Pure Chemistry: Acids, Bases and Salts — A Systematic Framework for Choosing the Right Salt Preparation Method

Why Acids, Bases and Salts Trips Up So Many Students

Ask a Secondary 3 or 4 student to list the four methods of salt preparation and most can recite them: titration, excess (insoluble reactant), precipitation, and direct combination. The problem isn't the list — it's that students memorise the methods as four unrelated recipes, then freeze when the exam gives them a salt they haven't seen before, like lead(II) chromate or ammonium sulfate. They start guessing based on which method "feels right" for the question, rather than reasoning it out.

This is exactly the kind of topic where memorisation fails and understanding wins. There is really only one decision tree behind all four methods, built on two questions: is the salt you want soluble or insoluble in water, and is the starting material you're given soluble or insoluble? Once a student can answer those two questions confidently, the "right" method falls out automatically — no memorising required, and it works for any salt the exam throws at them.

The Core Ideas You Actually Need

1. What Makes an Acid an Acid

At O-Level, an acid is a substance that releases H⁺ ions (technically, donates protons) when dissolved in water. This single fact explains almost every acid property you're asked to recall:

• Acids turn blue litmus red and have a pH below 7 because of the excess H⁺ ions in solution.
• Acids react with reactive metals to release hydrogen gas: metal + acid → salt + H₂
• Acids react with carbonates to release carbon dioxide: carbonate + acid → salt + H₂O + CO₂
• Acids neutralise bases: base + acid → salt + water (+ CO₂ if the base is a carbonate)

Strength and concentration are two different ideas students frequently mix up. Strength is about how completely the acid ionises in water (hydrochloric acid ionises fully — strong; ethanoic acid only partially ionises — weak). Concentration is simply how many moles of acid are dissolved per dm³ of solution. A dilute strong acid and a concentrated weak acid can have a similar pH, which is a classic distractor in MCQ questions.

2. Bases, Alkalis and the pH Scale

A base is any substance that reacts with an acid to form a salt and water. An alkali is simply a base that is soluble in water — so all alkalis are bases, but not all bases are alkalis (copper(II) oxide is a base but not an alkali, because it's insoluble). This distinction matters enormously for salt preparation, which is the next section.

The pH scale runs from 0 to 14. Below 7 is acidic, above 7 is alkaline, and exactly 7 is neutral. What examiners actually test is whether you understand that pH is a logarithmic scale — a change of one pH unit means a tenfold change in H⁺ concentration. So a solution of pH 2 is ten times more acidic than one of pH 3, not "slightly more acidic."

3. The Four Salt Preparation Methods — One Decision Tree

Instead of memorising four separate methods, walk through this sequence of questions every time:

1. Is the salt you need to make soluble or insoluble in water?
   • If insoluble → use precipitation (double decomposition). Mix two soluble solutions whose ions combine to form your insoluble salt, then filter, wash, and dry the residue.
   • If soluble → go to question 2.
2. Is the starting base/metal/carbonate soluble or insoluble in water?
   • If the starting material is insoluble (e.g. copper(II) oxide, a metal, or an insoluble carbonate) → use the excess method: add the insoluble solid in excess to the acid, so all the acid is used up, then filter off the unreacted excess solid and crystallise the filtrate.
   • If the starting material is soluble (e.g. sodium hydroxide, potassium hydroxide, ammonia solution) → you cannot add it in excess and simply filter, because there's nothing solid to filter off. Instead, use titration: neutralise the acid and alkali exactly using an indicator to find the correct volumes, then repeat without indicator and crystallise.

That's the entire topic. Two yes/no questions, and the method is decided. (The fourth "method" some syllabuses mention — direct combination of elements, e.g. iron + sulfur → iron(II) sulfide — is a special case reserved for a handful of non-oxysalts and is usually flagged explicitly in the question.)

4. Reading Solubility Rules Instead of Memorising Every Salt

You need a small set of solubility rules to answer question 1 above. Rather than memorising hundreds of individual salts, learn these patterns:

• All sodium, potassium and ammonium salts are soluble.
• All nitrates are soluble.
• All common chlorides are soluble, except silver chloride and lead(II) chloride.
• All common sulfates are soluble, except lead(II) sulfate, barium sulfate and calcium sulfate (calcium sulfate is only slightly soluble).
• All common carbonates are insoluble, except sodium, potassium and ammonium carbonate.

Notice the structure: sodium, potassium and ammonium salts are always soluble, and nitrates are always soluble — those two rules alone let you eliminate most confusion quickly.

Worked Examples: Applying the Decision Tree

Example 1: Preparing Copper(II) Sulfate Crystals

Question: Describe how you would prepare a pure, dry sample of copper(II) sulfate crystals from copper(II) oxide and dilute sulfuric acid.

Reasoning: Copper(II) sulfate is a common sulfate, so it's soluble → not precipitation. Copper(II) oxide is a metal oxide, which is insoluble in water → excess method.

Method:

1. Warm a fixed volume of dilute sulfuric acid in a beaker (warming speeds up the reaction, but do not boil, or acid will be lost as vapour).
2. Add copper(II) oxide powder gradually, stirring after each addition, until it is in excess — you will see unreacted black powder settle at the bottom, showing all the acid has reacted.
3. Filter the mixture to remove the excess unreacted copper(II) oxide, collecting the blue filtrate (copper(II) sulfate solution).
4. Heat the filtrate gently to evaporate off some water, until a hot saturated solution is obtained (test by dipping a glass rod — crystals should form at the tip on cooling).
5. Leave the solution to cool slowly and crystallise.
6. Filter off the crystals and pat dry between sheets of filter paper (do not use an oven, which may cause water of crystallisation to be lost).

Exam note: No indicator is used in this method, because the colour change of the excess solid tells you the reaction is complete — adding an indicator here is a very common wrong answer that costs marks.

Example 2: Preparing Sodium Sulfate Crystals (Titration)

Question: Sodium sulfate is soluble, and both sodium hydroxide and dilute sulfuric acid are soluble. Describe how sodium sulfate crystals are prepared.

Reasoning: The salt is soluble → not precipitation. Both starting materials are soluble, so there's no solid excess to filter → titration is required.

Method:

1. Using a pipette, measure a fixed volume of sodium hydroxide solution into a conical flask and add a few drops of a suitable indicator (e.g. methyl orange).
2. Fill a burette with dilute sulfuric acid. Run the acid into the alkali slowly, swirling constantly, until the indicator just changes colour — this is the end point. Record the volume of acid used.
3. Repeat the titration without indicator, adding the exact same recorded volume of acid to a fresh, identical volume of sodium hydroxide. This step is essential because indicators are organic dyes that would contaminate the final crystals.
4. Heat the resulting solution to evaporate it to the point of saturation, then cool to crystallise sodium sulfate.
5. Filter and dry the crystals as before.

Example 3: Preparing Lead(II) Iodide (Precipitation)

Question: Lead(II) iodide is insoluble. Describe how you would prepare a pure, dry sample from lead(II) nitrate and potassium iodide solutions.

Reasoning: The target salt is insoluble → skip straight to precipitation, no need to check solubility of the reactants.

Method:

1. Mix aqueous lead(II) nitrate with aqueous potassium iodide in a beaker. A bright yellow precipitate of lead(II) iodide forms immediately, alongside soluble potassium nitrate remaining in solution.
2. Filter the mixture to collect the yellow precipitate as residue.
3. Wash the residue with a small amount of distilled water to remove soluble impurities (potassium nitrate) — this is why washing appears in almost every precipitation mark scheme.
4. Dry the residue between sheets of filter paper, or leave in a warm place (not a hot oven, to avoid decomposition).

The Titration Calculation Skill

Once the practical method is understood, exams also test the numerical side: using titration results to find concentration or volume. The tool is always the same relationship.

moles = concentration (mol/dm³) × volume (dm³)

Worked example: 25.0 cm³ of sodium hydroxide solution of unknown concentration was exactly neutralised by 15.0 cm³ of 0.500 mol/dm³ hydrochloric acid. Find the concentration of the sodium hydroxide.

Step 1 — write the balanced equation: NaOH + HCl → NaCl + H₂O. The mole ratio of NaOH : HCl is 1 : 1.

Step 2 — find moles of HCl used: moles = 0.500 × (15.0 / 1000) = 0.00750 mol.

Step 3 — use the 1:1 ratio to find moles of NaOH: 0.00750 mol.

Step 4 — find concentration of NaOH: concentration = moles / volume = 0.00750 / (25.0 / 1000) = 0.300 mol/dm³.

The one place this trips students up is when the mole ratio is not 1:1 — for example, in reactions with sulfuric acid (a diprotic acid) or with metal hydroxides like calcium hydroxide. Always check the balanced equation for the ratio before dividing.

Common Exam Traps

• Using indicator in the excess method. The excess method never needs an indicator — the visible excess of solid is the evidence the reaction is complete. Adding indicator here is one of the most frequent wrong answers.
• Skipping the "repeat without indicator" step in titration. Many students describe titration correctly but forget that the final crystallisation must use a fresh, indicator-free mixture, since the dye would contaminate the pure salt.
• Trying to make a soluble salt of Na, K or NH₄⁺ using the excess method. These metals (or ammonium) are too reactive, or their hydroxides/carbonates are soluble, so there's no insoluble solid to add in excess and filter off — titration is the only option for their salts.
• Forgetting to wash the precipitate in precipitation questions. Omitting the wash step is a very common one-mark loss, because it leaves soluble by-product ions contaminating the final solid.
• Assuming all sulfates and carbonates follow the same solubility pattern. Remember the specific exceptions — lead(II), barium and calcium sulfates are insoluble; sodium, potassium and ammonium carbonates are the only soluble carbonates.
• Mixing up moles and concentration in titration calculations. Always convert cm³ to dm³ (divide by 1000) before using the moles = concentration × volume formula — a very common arithmetic slip under exam time pressure.

How to Practise This Topic Well

The fastest way to lock this topic in is not re-reading notes, but forcing yourself through the decision tree on unfamiliar salts until it becomes automatic:

• Take a list of ten random salts (e.g. zinc chloride, barium sulfate, ammonium nitrate, calcium carbonate) and, for each one, write down which of the four methods you'd use and why — before checking the answer.
• Practise writing out full preparation methods from memory, including the exact reasons for each step (why filter, why wash, why repeat without indicator) — O-Level mark schemes reward these justifications, not just the steps themselves.
• Work through past-year titration calculation questions with varying mole ratios (1:1, 1:2, 2:1) so the "check the equation first" habit becomes second nature.
• Use flashcards specifically for the solubility exceptions (Ag⁺, Pb²⁺, Ba²⁺, Ca²⁺) since these are the few facts in this topic that genuinely need to be memorised — everything else should be reasoned through the decision tree.

At Intuitional, we teach acids, bases and salts the same way we teach every calculation-heavy Chemistry topic: build the reasoning framework first, then drill it against real exam questions until the "aha, so that's why it's this method" moment sticks — because a framework a student understands will still work on a salt they've never seen before, while a memorised list won't.