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Cambridge IGCSE • Year 11 • Chemistry
Acids, Bases and Salts
Acid-base chemistry, salts, pH and qualitative reactions.
Chapter 7
Verified Curriculum Topic
What is Acids, Bases and Salts?
Acid-base chemistry, salts, pH and qualitative reactions.
Acids, Bases and Salts matters because it links chemical ideas, reactions, and reasoning patterns that recur throughout the syllabus. At Year 11 level, students are often expected to define terms accurately, explain processes clearly, and connect theory to reactions, observations, or applications.
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Summary
The One Thing
Acidic and alkaline behaviour is determined by the ions present in aqueous solution: acids produce hydrogen ions, H+, whereas alkalis produce hydroxide ions, OH-. Their reactions, pH values, salt-preparation methods, and qualitative tests can therefore be interpreted in terms of these ions and the products formed.
Reactions, Processes and Experiments
| What happens | Equation or process | What you observe | Type |
|---|---|---|---|
| An acid ionises in water to produce hydrogen ions and an anion. | acid(aq) ⇌ H+(aq) + anion(aq) | — | Ionisation |
| Hydrogen ions react with hydroxide ions to form water. | H+(aq) + OH-(aq) → H2O(l) | — | Neutralisation; ionic equation |
| Hydrochloric acid reacts with magnesium to form magnesium chloride and hydrogen. | 2HCl + Mg → MgCl2 + H2 | Hydrogen gas is produced. | Acid plus metal |
| Nitric acid reacts with copper(II) oxide to form copper(II) nitrate and water. | 2HNO3 + CuO → Cu(NO3)2 + H2O | — | Acid plus metal oxide |
| Sulfuric acid reacts with sodium hydroxide to form sodium sulfate and water. | H2SO4 + 2NaOH → Na2SO4 + 2H2O | — | Acid plus metal hydroxide; neutralisation |
| Hydrochloric acid reacts with calcium carbonate to form calcium chloride, water, and carbon dioxide. | 2HCl + CaCO3 → CaCl2 + H2O + CO2 | Carbon dioxide is released. | Acid plus carbonate |
| An acid reacts with a hydrogencarbonate to form a salt, water, and carbon dioxide. | acid + hydrogencarbonate → salt + water + carbon dioxide | Carbon dioxide is released. | Acid plus hydrogencarbonate |
| A soluble salt is prepared from two soluble reactants. The solutions are mixed in the exact reacting volumes, identified using an indicator, and the resulting solution is evaporated and crystallised. | Titration is used when both reactants are soluble, because excess reactant cannot be removed by filtration. | An indicator changes colour at the end point; crystals form during crystallisation. | Salt preparation; titration |
| An insoluble metal oxide, hydroxide, or carbonate is added to a warm acid until no more reacts. The excess solid is filtered off, and the solution is evaporated and crystallised. | Add the insoluble solid to warm acid until no more reacts; filter off excess solid; evaporate and crystallise the solution. | Excess solid remains and is removed by filtration; crystals form from the solution. | Salt preparation |
| Two soluble solutions react to form an insoluble salt. The solid is filtered, washed, and dried. | Mix two soluble solutions to form an insoluble precipitate, then filter, wash, and dry the solid. | An insoluble precipitate forms. | Precipitation; salt preparation |
| Lithium ions are identified by heating the sample in a flame. | Flame test for lithium ions. | Crimson-red flame. | Qualitative analysis |
| Sodium ions are identified by heating the sample in a flame. | Flame test for sodium ions. | Yellow flame. | Qualitative analysis |
| Potassium ions are identified by heating the sample in a flame. | Flame test for potassium ions. | Lilac flame. | Qualitative analysis |
| Calcium ions are identified by heating the sample in a flame. | Flame test for calcium ions. | Brick-red flame. | Qualitative analysis |
| Copper(II) ions are identified by heating the sample in a flame. | Flame test for copper(II) ions. | Blue-green flame. | Qualitative analysis |
| A sample is warmed with sodium hydroxide to test for ammonium ions. | Warm the sample with sodium hydroxide. | Ammonia gas is released and turns damp red litmus paper blue. | Test for ammonium ions |
| A dilute acid is added to test for carbonate ions. | Add a dilute acid, then test the gas with limewater. | Carbon dioxide is released and turns limewater milky. | Test for carbonate ions |
| A sample is acidified with dilute nitric acid and aqueous silver nitrate is added to test for chloride ions. | Acidify with dilute nitric acid, then add aqueous silver nitrate. | A white precipitate of silver chloride forms. | Test for chloride ions |
| A sample is acidified with dilute nitric acid and aqueous silver nitrate is added to test for bromide ions. | Acidify with dilute nitric acid, then add aqueous silver nitrate. | A cream precipitate of silver bromide forms. | Test for bromide ions |
| A sample is acidified with dilute nitric acid and aqueous silver nitrate is added to test for iodide ions. | Acidify with dilute nitric acid, then add aqueous silver nitrate. | A yellow precipitate of silver iodide forms. | Test for iodide ions |
| A sample is acidified with dilute hydrochloric acid and barium nitrate or barium chloride is added to test for sulfate ions. | Acidify with dilute hydrochloric acid, then add barium nitrate or barium chloride. | A white precipitate of barium sulfate forms. | Test for sulfate ions |
| A sample is treated with aqueous sodium hydroxide to test for metal ions. | Add aqueous sodium hydroxide. | Copper(II) gives a light blue precipitate; iron(II) gives a green precipitate; iron(III) gives a red-brown precipitate; aluminium gives a white precipitate that dissolves in excess sodium hydroxide. | Test for metal ions |
| Litmus is used to distinguish acidic and alkaline solutions. | Add litmus indicator. | Red in acid and blue in alkali. | Indicator test |
| Methyl orange is used to distinguish acidic and alkaline solutions. | Add methyl orange indicator. | Red in acid and yellow in alkali. | Indicator test |
| Phenolphthalein is used to distinguish acidic and alkaline solutions. | Add phenolphthalein indicator. | Colourless in acid and pink in alkali. | Indicator test |
| Universal indicator is used to estimate the pH of a solution. | Use a mixture of indicators that gives a range of colours across the pH scale. | A colour corresponding to the solution’s approximate pH is produced. | Indicator test |
| The pH of a solution is measured on a scale commonly ranging from 0 to 14. | pH 0-6 is acidic, pH 7 is neutral, and pH 8-14 is alkaline. | Lower pH values indicate greater acidity; higher pH values indicate greater alkalinity. | pH measurement |
| A decrease of one pH unit changes the hydrogen ion concentration. | A decrease of one pH unit represents approximately a tenfold increase in hydrogen ion concentration. | — | Quantitative relationship |
| The concentration of a solution is calculated from amount of substance and volume. | concentration = amount of substance / volume | — | Concentration calculation |
| The amount of substance is calculated from concentration and volume. | amount of substance = concentration × volume | — | Amount-of-substance calculation |
| In a 1:1 acid-base reaction, the reacting amounts in moles are equal. | Compare directly the amount in moles of acid with the amount in moles of alkali. | — | Stoichiometric calculation |
| In reactions with ratios other than 1:1, the balanced equation determines the reacting mole ratio. | Use the coefficients in the balanced equation. | — | Stoichiometric calculation |
| Spectator ions are removed when writing an ionic equation. | Omit ions that do not change during the reaction. | — | Ionic equation |
| Solubility rules are used to determine whether a salt dissolves or forms a precipitate. | All nitrates and ammonium salts are soluble; most chlorides are soluble except silver and lead chlorides; most sulfates are soluble except barium, lead, and calcium sulfate; most carbonates and hydroxides are insoluble except those of group 1 metals and ammonium, with some calcium compounds being sparingly soluble. | A precipitate forms when an insoluble compound is produced. | Solubility analysis |
Key Terms
- Acid: A substance that produces hydrogen ions, H+, when dissolved in water.
- Base: A substance that reacts with an acid to neutralise it; bases often contain metal oxides, hydroxides, or carbonates.
- Alkali: A soluble base that produces hydroxide ions, OH-, in aqueous solution.
- Hydrogen ion: The ion, H+, associated with acidic properties in aqueous solutions.
- Hydroxide ion: The ion, OH-, associated with alkaline properties.
- Neutralisation: A reaction in which an acid reacts with a base to form a salt and water.
- Salt: An ionic compound formed when the hydrogen ion of an acid is replaced by a metal ion or an ammonium ion.
- pH: A numerical measure of acidity or alkalinity; lower values are more acidic, pH 7 is neutral, and higher values are more alkaline.
- Indicator: A substance that changes colour over a particular pH range, helping to show whether a solution is acidic or alkaline.
- Universal indicator: A mixture of indicators that gives a range of colours across the pH scale.
- Strong acid: An acid that ionises almost completely in water, such as hydrochloric acid, nitric acid, or sulfuric acid.
- Weak acid: An acid that only partially ionises in water, such as ethanoic acid.
- Concentration: The amount of solute dissolved in a given volume of solution, commonly measured in mol/dm3.
- Solubility: The maximum amount of a substance that dissolves in a specified amount of solvent at a particular temperature.
- Precipitate: An insoluble solid formed when two solutions react.
- Qualitative analysis: Testing used to identify the ions or substances present in a sample.
- Titration: A quantitative method used to find the volume of one solution needed to react exactly with a measured volume of another solution.
Easily Confused
- Strong acid and concentrated acid: Strength describes how fully an acid ionises in water; concentration describes the amount of solute per unit volume. A strong acid is not necessarily concentrated.
- Base and alkali: A base neutralises an acid, whereas an alkali is a base that dissolves in water and produces OH- ions.
- Neutralisation and precipitation: Neutralisation forms a salt and water through
H+(aq) + OH-(aq) → H2O(l); precipitation forms an insoluble solid from two solutions. - Titration and filtration: Titration is required when both reactants are soluble; filtration removes an insoluble excess or an insoluble precipitate.
- Chloride, bromide, and iodide tests: With acidified nitric acid followed by aqueous silver nitrate, chloride gives a white precipitate, bromide a cream precipitate, and iodide a yellow precipitate.
- Metal-ion precipitates: Aqueous sodium hydroxide gives copper(II) a light blue precipitate, iron(II) a green precipitate, iron(III) a red-brown precipitate, and aluminium a white precipitate that dissolves in excess sodium hydroxide.
- Litmus, methyl orange, and phenolphthalein: Litmus is red in acid and blue in alkali; methyl orange is red in acid and yellow in alkali; phenolphthalein is colourless in acid and pink in alkali.
- Acid plus metal and acid plus carbonate: An acid plus a metal produces hydrogen, whereas an acid plus a carbonate produces carbon dioxide, water, and a salt.
- Acid plus metal oxide or hydroxide and acid plus carbonate: Metal oxides and hydroxides produce a salt and water; carbonates additionally produce carbon dioxide.
What Gets Asked
- Explain acid and alkaline behaviour in terms of hydrogen ions, H+, and hydroxide ions, OH-. A common mark-losing error is to describe a base as necessarily soluble; only soluble bases are alkalis.
- Complete and classify equations for acid reactions with metals, metal oxides, hydroxides, carbonates, and hydrogencarbonates. The specific gas must be identified: hydrogen with metals and carbon dioxide with carbonates and hydrogencarbonates.
- Interpret pH values and indicators. Questions may test that pH 7 is neutral, pH 0-6 is acidic, pH 8-14 is alkaline, and a decrease of one pH unit represents approximately a tenfold increase in hydrogen ion concentration.
- Distinguish acid strength from concentration. Strong acids such as hydrochloric acid, nitric acid, and sulfuric acid ionise almost completely, whereas weak ethanoic acid only partially ionises; neither description alone specifies concentration.
- Select a salt-preparation method from solubility information. Titration is used when both reactants are soluble, while an insoluble excess is filtered off when an insoluble metal oxide, hydroxide, or carbonate reacts with a warm acid.
- Identify ions from qualitative-test observations. Marks depend on linking the stated colour, gas, or precipitate to the correct ion, such as a white silver chloride precipitate for chloride or a white barium sulfate precipitate for sulfate.
Flashcards
Quick quiz
Which ion is responsible for acidic properties in aqueous solutions?
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Learning objectives
- 7.1Describe the characteristic properties of acids and bases, including their effect on litmus and reaction with metals.
- 7.2Define pH and relate pH values to acidity, neutrality, and alkalinity using the pH scale and indicators.
- 7.3Describe the general reactions of acids with metals, metal oxides, metal hydroxides, and carbonates, including word equations.
- 7.4Define a salt and describe methods of preparing soluble salts, including titration and excess-solid methods.extended
- 7.5Explain neutralisation in terms of the reaction between H+ ions and OH− ions to form water.extended
- 7.6Describe a simple test to identify carbonate, sulfate, and chloride ions in an unknown salt.extended
Syllabus-verified
Practice questions
Q1. A solution has a pH of 3. This solution is:1 mark · core
- A. Strongly alkaline
- B. Neutral
- C. Weakly alkaline
- D. Acidic
Answer: D
- • 1 mark for selecting D
Any pH value below 7 indicates an acidic solution; pH 3 is strongly acidic.
Q2. Describe how a pure, dry sample of the soluble salt copper sulfate could be prepared from copper oxide (an insoluble base) and dilute sulfuric acid.4 marks · extended
Answer: Add excess copper oxide powder to warm dilute sulfuric acid and stir until no more dissolves, to ensure all the acid has reacted. Filter the mixture to remove the unreacted excess copper oxide, leaving a solution of copper sulfate. Heat the filtrate gently to evaporate some water and produce a saturated solution, then leave it to cool and crystallise. Filter off the crystals and dry them between sheets of filter paper.
- • 1 mark: add excess (insoluble) copper oxide to the acid to ensure all acid reacts
- • 1 mark: filter to remove unreacted excess copper oxide
- • 1 mark: heat/evaporate the filtrate to produce a saturated solution
- • 1 mark: cool to crystallise, then filter and dry the crystals
Q3. Write a word equation for the reaction between magnesium metal and dilute hydrochloric acid.2 marks · core
Answer: Magnesium + hydrochloric acid → magnesium chloride + hydrogen.
- • 1 mark: correct reactants named (magnesium and hydrochloric acid)
- • 1 mark: correct products named (magnesium chloride and hydrogen)
Q4. Write the ionic equation for the neutralisation reaction between an acid and an alkali, and explain what it shows.2 marks · extended
Answer: H+(aq) + OH−(aq) → H2O(l). It shows that neutralisation is fundamentally the reaction between hydrogen ions (from the acid) and hydroxide ions (from the alkali) to form water, regardless of which specific acid or alkali is used.
- • 1 mark: correct ionic equation H+ + OH− → H2O with correct state symbols
- • 1 mark: correct explanation that neutralisation is the combination of H+ and OH− ions to form water, independent of the specific acid/alkali
Key ideas to master
- Learn the precise terms, laws, and reaction patterns associated with Acids, Bases and Salts.
- Understand why each step or change happens instead of memorising the result only.
- Practise writing balanced equations, comparisons, or structured explanations where relevant.
- Revise common exceptions, observations, and applications that examiners often test.
Common exam prompts
- Define the main idea in Acids, Bases and Salts using correct chemical terminology.
- Write or interpret the reactions, observations, or comparisons that belong to this topic.
- Explain why a process happens, not just what happens.
- Summarise the high-yield facts and exceptions examiners often choose from this chapter.
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Quick answers students usually need
What is Acids, Bases and Salts in Cambridge IGCSE Year 11 Chemistry?
Acid-base chemistry, salts, pH and qualitative reactions.
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