Acids and Bases: Neutralisation and Salts
AdvancedNeutralisation is the exothermic reaction of an acid with a base giving salt plus water — essentially hydrogen ions joining hydroxide ions to form water. Titration uses a standard solution and an indicator to find an unknown concentration; salts are prepared by crystallisation or precipitation according to solubility.
- Neutralisation: acid + base → salt + water; H⁺(aq) + OH⁻(aq) → H₂O(l); exothermic
- Antacids carry mild bases (magnesium/aluminium hydroxide) against gastric acid; slaked lime fixes acidic soil; alkali treats acidic wastewater
- Titration: burette with standard solution, two or three drops of indicator, end point at the colour change; n = concentration × volume (dm³)
文體小知識：科學探究記敘——記述一次酸鹼滴定：由胃藥中和胃酸的原理，到用標準溶液和指示劑測定濃度，再到鹽的製備路線。
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1Gastric juice contains hydrochloric acid; excess causes burning pain. Bases in antacids, such as magnesium hydroxide, react away the excess — this is neutralisation: an acid and a base react to give a salt and water. Hydrochloric acid with sodium hydroxide yields sodium chloride and water. The essence is one line: the acid's hydrogen ions join the alkali's hydroxide ions to form water molecules, while the remaining ions stay in solution as the salt. Neutralisation is exothermic — mixing the solutions warms them slightly.
2Applications surround us: farmers scatter slaked lime over acidic soil to improve it; factories neutralise acidic wastewater with alkali before discharge to protect rivers; an ant bite dabbed with dilute ammonia or baking-soda paste neutralises formic acid to stop the itch. Choose the neutralising agent sparingly — too little leaves acid, too much brings new corrosiveness — which is why antacids use mild bases in calculated doses.
3In the laboratory, acid-alkali titration measures neutralisation precisely: fill a burette with a standard solution of known concentration, pipette a measured volume of the unknown solution into a conical flask, add two or three drops of indicator, then add drop by drop with swirling until the indicator just changes colour — the end point — and record the volume used to calculate the unknown concentration. Concentration is expressed as molar concentration, moles of solute per cubic decimetre of solution, the bridge of all such calculations.
4The products of neutralisation are salts, built from a metal ion (or ammonium ion) and an acid anion: sodium chloride, copper(II) sulfate, potassium nitrate — the names telling their origins: chloride from hydrochloric, sulfate from sulfuric, nitrate from nitric acid. To prepare a soluble salt, react the acid with an excess of an insoluble base, filter, then evaporate the filtrate to crystallisation; for an insoluble salt, mix two suitable solutions and let the salt precipitate, then filter. The wrong route wastes every effort.
1Summary 重點摘要
Neutralisation is the exothermic reaction of an acid with a base giving salt plus water — essentially hydrogen ions joining hydroxide ions to form water. Titration uses a standard solution and an indicator to find an unknown concentration; salts are prepared by crystallisation or precipitation according to solubility.
2Vocabulary cards Vocabulary
The reaction of an acid with a base giving salt and water — hydrogen ions joining hydroxide ions to form water; exothermic.
A compound of a metal ion (or ammonium ion) with an acid anion, e.g. sodium chloride, copper(II) sulfate.
A medicine containing mild bases (magnesium or aluminium hydroxide) that neutralises gastric acid and relieves pain.
Adding one solution dropwise to another until the indicator shows exact neutralisation, allowing concentrations to be calculated.
A solution of accurately known concentration, the reference of a titration.
The moment the indicator just changes colour, showing the two solutions have reacted exactly.
Moles of solute per cubic decimetre of solution, in mol dm⁻³.
A long graduated tube with a tap for delivering and reading accurate volumes, holding the standard solution in titration.
A salt that dissolves in water, e.g. sodium, potassium and nitrate salts, usually made from acid plus base then crystallised.
The solid formed when solutions are mixed; insoluble salts are made by precipitation and collected by filtering.
3Step by step Steps
- Step 1: The essence of neutralisationMix dilute hydrochloric acid with sodium hydroxide solution and HCl + NaOH → NaCl + H₂O takes place — salt and water. Behind the formula plays a single scene: the acid's hydrogen ions meet the alkali's hydroxide ions and combine into water; the ionic equation is H⁺(aq) + OH⁻(aq) → H₂O(l). Sodium and chloride ions merely watch, remaining in solution as sodium chloride. Every strong-acid-strong-alkali neutralisation shares this same ionic equation — the key of the topic, to be memorised word for word.
- Step 2: How antacids tame gastric acidWhen gastric juice turns too acidic, burning pain follows; antacids carry mild bases to mop up the excess hydrogen ions. For example, magnesium hydroxide reacts with gastric acid: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O — one mole of the hydroxide neutralises two moles of the acid, a 1:2 ratio. The base chosen is mild and insoluble, neutralising slowly and steadily so the stomach is not swung alkaline; drug companies dose each tablet with exactly the mole calculations behind titration — medicine meets chemistry here.
- Step 3: Neutralisation in daily lifeApplications run on three levels. On farms, slaked lime (calcium hydroxide) is scattered over acidic soil to neutralise it so crops thrive; in industry, acidic wastewater is neutralised with alkali to near-neutral before discharge, sparing the rivers; at home, an ant bite deposits formic acid, eased by dilute ammonia or baking-soda paste. For example, after perming, hairdressers rinse with a weak acid to neutralise residual alkali — same principle. Each case reminds us that neutralisation is written not only on exam papers but across farms, factories and pharmacies.
- Step 4: Titration — apparatus and procedureThe standard set-up: a burette holds the standard solution; a pipette delivers 25.0 cm³ of the unknown solution into a conical flask with two or three drops of indicator. Left hand works the tap, right hand swirls; near the end point add drop by drop, even half-drops, until the indicator just changes and holds for half a minute — record the burette readings, the difference being the volume used. For example, with phenolphthalein the alkali in the flask is pink; stop the instant the pink fades — one drop too many skews the calculation. Repeat two or three times and average for reliability.
- Step 5: Titration calculation in three movesThe calculation has three moves: moles from concentration times volume, convert by the equation's mole ratio, then back-calculate. Take the standard case: 25.0 cm³ of 0.1 M sodium hydroxide titrated with 0.1 M hydrochloric acid. Move one: moles of NaOH = 0.1 × 25.01000 = 2.5×10⁻³ mol. Move two: the ratio is 1:1, so HCl needed is likewise 2.5×10⁻³ mol. Move three: volume = 2.5×10⁻³ ÷ 0.1 = 0.025 dm³ = 25.0 cm³. Always convert volumes to dm³ before substituting — the most common slip in exams.
- Step 6: Naming salts and solubility rulesA salt's name pairs the metal (or ammonium) with the acid anion: hydrochloric acid gives chlorides, sulfuric gives sulfates, nitric gives nitrates, carbonic gives carbonates. Solubility follows rules: sodium, potassium, ammonium and nitrate salts all dissolve; chlorides dissolve except silver and lead chlorides; sulfates dissolve except barium, lead and calcium sulfates; most carbonates are insoluble. For example, making brine is effortless because sodium chloride dissolves, while silver chloride precipitates on sight — the very test for chloride ions. Master the rules and the preparation routes become obvious.
- Step 7: Two routes to a saltFirst ask: is the target salt soluble? Soluble salts go the crystallisation route: warm the dilute acid with an excess of an insoluble base so the acid is fully used up, filter off the leftover solid, evaporate part of the filtrate and cool to crystallise. For example, to make copper(II) sulfate crystals, stir copper(II) oxide into warm dilute sulfuric acid until in excess, filter, crystallise — blue crystals result. Insoluble salts go the precipitation route: mix two suitable solutions, let the salt precipitate, then filter, wash and dry. For example, mixing silver nitrate with sodium chloride solutions drops white silver chloride at once. The choice rests on one question: does the salt dissolve in water?
4Diagrams Diagrams
5Worked example Example
Experimental context: an antacid tablet is labelled as containing magnesium hydroxide. In the laboratory, a student pipettes 25.0 cm³ of 0.1 M sodium hydroxide into a conical flask, adds two drops of phenolphthalein, and titrates with 0.1 M hydrochloric acid from a burette; 25.0 cm³ of acid is used when the pink just fades. Answer: (1) the ionic equation of neutralisation; (2) the moles of sodium hydroxide in the flask; (3) verify the volume of acid required; (4) the equation for magnesium hydroxide neutralising gastric acid.
- Ionic equation: the essence of strong acid plus strong alkali is H⁺(aq) + OH⁻(aq) → H₂O(l)
- Moles: moles of NaOH = concentration × volume = 0.1 × 25.01000 = 2.5×10⁻³ mol
- Convert by ratio: HCl reacts 1:1 with NaOH, so 2.5×10⁻³ mol is needed; volume = 2.5×10⁻³ ÷ 0.1 = 0.025 dm³ = 25.0 cm³, matching the experiment
- Antacid equation: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O, a 1:2 ratio
- Answer: The ionic equation of neutralisation is H⁺(aq) + OH⁻(aq) → H₂O(l). Moles of sodium hydroxide = 0.1 mol dm⁻³ × 25.01000 dm³ = 2.5×10⁻³ mol. Hydrochloric acid reacts 1:1, so 2.5×10⁻³ mol is required, a volume of 2.5×10⁻³ ÷ 0.1 = 0.025 dm³ — 25.0 cm³, agreeing with the burette reading. In the antacid, magnesium hydroxide neutralises gastric acid by Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O; one mole of the hydroxide neutralises two moles of the acid.
6Common mistakes Mistakes
7Quick check Quick check
Q1Write the ionic equation for neutralisation.
Q2How many moles of sodium hydroxide are in 25.0 cm³ of 0.1 M solution?
Q3If the alkali above is titrated with 0.2 M hydrochloric acid instead, what volume is needed?
Q4Which base in antacids neutralises gastric acid? Give the equation.
Q5How do the preparations of soluble and insoluble salts differ?
8Recap Recap
- Neutralisation: acid + base → salt + water; H⁺(aq) + OH⁻(aq) → H₂O(l); exothermic
- Antacids carry mild bases (magnesium/aluminium hydroxide) against gastric acid; slaked lime fixes acidic soil; alkali treats acidic wastewater
- Titration: burette with standard solution, two or three drops of indicator, end point at the colour change; n = concentration × volume (dm³)
- Salts pair a metal (or ammonium) ion with an acid anion; Na/K/NH₄/nitrates soluble, silver chloride and barium sulfate not
- Preparation: soluble salts via acid + excess insoluble base then crystallisation; insoluble salts via precipitation and filtering