Acids, Bases and Salts
Indicators
Indicators are substances that change colour depending on whether a solution is acidic or alkaline.
| Indicator | Acidic solution | Neutral | Alkaline solution |
|---|---|---|---|
| Litmus | Red | Purple | Blue |
| Phenolphthalein | Colourless | Colourless | Pink/magenta |
| Methyl orange | Red | Orange | Yellow |
Universal indicator gives a range of colours corresponding to different pH values, but is not suitable for titrations because it does not give a sharp, definite colour change at the endpoint.
The pH Scale
The pH scale runs from 0 to 14 and classifies solutions by their acidity or alkalinity:
| pH range | Classification |
|---|---|
| 0 to 3 | Strongly acidic |
| 4 to 6 | Weakly acidic |
| 7 | Neutral |
| 8 to 10 | Weakly alkaline |
| 11 to 14 | Strongly alkaline |
Universal indicator can be used to measure the approximate pH of an aqueous solution by comparing the colour of the indicator to a colour chart.
Acids and Alkalis
Acids in aqueous solution are a source of hydrogen ions, H+.
Alkalis in aqueous solution are a source of hydroxide ions, OH-.
Neutralisation is the reaction between hydrogen ions and hydroxide ions to form water:
H+(aq) + OH-(aq) → H2O(l)
Proton Transfer
Acids and bases can be understood in terms of proton transfer:
- An acid is a proton donor: it donates H+ ions (protons) to another substance.
- A base is a proton acceptor: it accepts H+ ions from an acid.
Example: When hydrochloric acid dissolves in water: HCl(aq) → H+(aq) + Cl-(aq)
The HCl donates a proton (H+), so it acts as an acid (proton donor).
Example: When sodium hydroxide neutralises an acid: NaOH(s) → Na+(aq) + OH-(aq)
The OH- accepts a proton from the acid, so the hydroxide ion acts as a base (proton acceptor).
Bases and Alkalis
Bases are substances that neutralise acids to form a salt and water. Bases include:
- Metal oxides (e.g. copper(II) oxide, CuO)
- Metal hydroxides (e.g. sodium hydroxide, NaOH)
- Ammonia (NH3)
An alkali is a base that is soluble in water. All alkalis are bases, but not all bases are alkalis.
Alkalis produce hydroxide ions (OH-) in aqueous solution:
| Alkali | Formula | Ions in solution |
|---|---|---|
| Sodium hydroxide | NaOH | Na+(aq) + OH-(aq) |
| Potassium hydroxide | KOH | K+(aq) + OH-(aq) |
| Aqueous ammonia | NH3(aq) | NH4+(aq) + OH-(aq) |
Reactions of Acids
With Metals
Only metals above hydrogen in the reactivity series react with dilute acids. The products are a salt and hydrogen gas:
metal + acid → salt + hydrogen
| Metal | With H2SO4 | With HCl |
|---|---|---|
| Magnesium | Mg(s) + H2SO4(aq) → MgSO4(aq) + H2(g) | Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g) |
| Zinc | Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g) | Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g) |
| Iron | Fe(s) + H2SO4(aq) → FeSO4(aq) + H2(g) | Fe(s) + 2HCl(aq) → FeCl2(aq) + H2(g) |
Note: Nitric acid is not used with metals in reactions covered by this specification, as it behaves differently.
With Bases (Metal Oxides and Hydroxides)
Acids react with metal oxides and metal hydroxides in neutralisation reactions to form a salt and water:
acid + base → salt + water
Examples:
- 2HCl(aq) + CuO(s) → CuCl2(aq) + H2O(l)
- H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l)
- HNO3(aq) + KOH(aq) → KNO3(aq) + H2O(l)
With Metal Carbonates
Acids react with metal carbonates to form a salt, water and carbon dioxide:
acid + metal carbonate → salt + water + carbon dioxide
Examples:
- 2HCl(aq) + Na2CO3(aq) → 2NaCl(aq) + H2O(l) + CO2(g)
- H2SO4(aq) + CaCO3(s) → CaSO4(aq) + H2O(l) + CO2(g)
Observation: Effervescence (fizzing) as carbon dioxide gas is produced.
With Ammonia
Ammonia acts as a base and neutralises acids to form ammonium salts:
HCl(g) + NH3(g) → NH4Cl(s)
Salt Names
The name of the salt produced depends on the acid used and the metal (or ammonium) ion in the base:
| Acid | Salt formed | Example |
|---|---|---|
| Hydrochloric acid (HCl) | Chloride | Sodium chloride, NaCl |
| Sulfuric acid (H2SO4) | Sulfate | Copper(II) sulfate, CuSO4 |
| Nitric acid (HNO3) | Nitrate | Potassium nitrate, KNO3 |
Solubility Rules
A knowledge of solubility helps determine the correct method for preparing a salt.
| Type of compound | Soluble | Insoluble |
|---|---|---|
| Sodium, potassium and ammonium compounds | All | None |
| Nitrates | All | None |
| Chlorides | Most | Silver chloride (AgCl), lead(II) chloride (PbCl2) |
| Sulfates | Most | Barium sulfate (BaSO4), calcium sulfate (CaSO4), lead(II) sulfate (PbSO4) |
| Carbonates | Sodium, potassium and ammonium carbonates | Most others |
| Hydroxides | Sodium, potassium hydroxides; calcium hydroxide is slightly soluble | Most others |
Exam Tip
Calcium hydroxide solution is also called limewater and is used to test for carbon dioxide.
Preparing a Soluble Salt from an Insoluble Reactant
When one reactant is insoluble (a metal oxide, metal carbonate or insoluble metal), this method is used.
Method (spec 2.39):
- Warm the acid in a beaker using a Bunsen burner.
- Add the insoluble base (in excess) to the acid and stir. Continue adding until no more dissolves and some solid remains at the bottom (this ensures all the acid has reacted).
- Filter the mixture to remove excess unreacted solid.
- Transfer the filtrate (salt solution) to an evaporating basin.
- Heat gently to evaporate some of the water until small crystals begin to appear (check by dipping a cold glass rod into the hot solution).
- Leave the solution to cool slowly so crystals form.
- Filter to collect the crystals, wash with a little distilled water, then dry on filter paper or in a warm oven.
Why is the base added in excess? To ensure all the acid has reacted. Unreacted acid would become dangerously concentrated during evaporation.
Practical (spec 2.42): Preparing copper(II) sulfate from copper(II) oxide:
CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l)
Copper(II) oxide (black solid) is added in excess to warm dilute sulfuric acid and stirred. The mixture is filtered to remove excess CuO. The blue filtrate is heated to evaporate water until saturated, then left to cool and crystallise. The resulting blue crystals of hydrated copper(II) sulfate (CuSO4.5H2O) are filtered, washed and dried.
Preparing a Soluble Salt from an Acid and an Alkali (Titration)
Higher Tier
This method (spec 2.40C) applies when both reactants are soluble, so filtration cannot be used to separate an excess.
Method:
- Use a pipette to measure a fixed volume of the alkali into a conical flask and add a few drops of indicator (phenolphthalein or methyl orange).
- Add the acid from a burette slowly, until the indicator just changes colour (the endpoint). Note the volume of acid used.
- Repeat the titration without indicator, using the exact volumes of acid and alkali determined. This gives a pure salt solution with no indicator contamination.
- Heat the solution to partially evaporate the water until saturated.
- Leave to cool slowly to crystallise, filter, and dry the crystals.
Apparatus:
- Pipette: for measuring a fixed, accurate volume of alkali
- Burette: for adding a variable volume of acid and measuring the exact amount used
- Conical flask: holds the reaction mixture
Why is universal indicator not used in titrations? It gives a gradual range of colours, so the endpoint is unclear and cannot be identified precisely.
Indicator colours:
- Phenolphthalein: pink/magenta in alkali, colourless in acid
- Methyl orange: yellow in alkali, red in acid
Preparing an Insoluble Salt by Precipitation
Higher Tier
This method (spec 2.41C) is used when the required salt is insoluble but two soluble reactants containing the needed ions are available.
Principle: Mix two solutions, each containing one of the required ions. The insoluble salt precipitates out of solution.
soluble salt 1 + soluble salt 2 → insoluble salt (precipitate) + soluble salt 3
Method:
- Dissolve each solid in water to make solutions.
- Mix the two solutions together and stir.
- Filter the mixture to collect the precipitate on the filter paper.
- Wash the precipitate with distilled water to remove traces of the soluble reactants.
- Dry on filter paper or in a warm oven.
Do not use direct heat to dry, as this may decompose the salt.
Practical (spec 2.43C): Preparing lead(II) sulfate from lead(II) nitrate and potassium sulfate:
Pb(NO3)2(aq) + K2SO4(aq) → PbSO4(s) + 2KNO3(aq)
Lead(II) nitrate is used (rather than lead(II) carbonate) because all nitrates are soluble, allowing a solution to be made. Lead(II) carbonate is insoluble, so it cannot form a solution for mixing.
Exam Tip
Take care when handling lead salts as they are toxic. Use appropriate safety precautions.
Ionic Equations for Precipitation
The overall ionic equation shows only the ions that form the precipitate (the spectator ions are omitted):
Example: Preparation of barium sulfate: Ba2+(aq) + SO42-(aq) → BaSO4(s)
Example: Preparation of silver chloride: Ag+(aq) + Cl-(aq) → AgCl(s)
Exam Questions
Q: What is the difference between a base and an alkali?
A base is any substance that neutralises an acid to form a salt and water. An alkali is a base that is soluble in water, producing hydroxide ions in solution. All alkalis are bases, but not all bases are alkalis.
Q: Explain, in terms of proton transfer, what happens when hydrochloric acid reacts with sodium hydroxide.
Hydrochloric acid donates protons (H+ ions) to the hydroxide ions from sodium hydroxide. The hydroxide ions accept these protons, forming water. Acid is a proton donor; base is a proton acceptor.
Q: Describe how you would prepare pure, dry crystals of copper(II) sulfate from copper(II) oxide and dilute sulfuric acid.
Warm the sulfuric acid in a beaker. Add copper(II) oxide in excess and stir until no more dissolves. Filter to remove excess copper(II) oxide. Heat the filtrate (copper(II) sulfate solution) gently until crystals begin to form. Leave to cool so crystals form. Filter off the crystals, wash with distilled water and dry on filter paper.
Q: Explain why the base is added in excess when preparing a soluble salt from an insoluble base.
The base is added in excess to ensure that all the acid has completely reacted. If unreacted acid remained, it would become very concentrated and dangerously corrosive during the evaporation stage.
Q: State the solubility of: (a) all nitrates (b) lead(II) chloride (c) sodium carbonate.
(a) All nitrates are soluble. (b) Lead(II) chloride is insoluble (most chlorides are soluble, except silver and lead(II)). (c) Sodium carbonate is soluble (sodium, potassium and ammonium carbonates are soluble).
Q: Describe how you would prepare a pure, dry sample of lead(II) sulfate.
Dissolve lead(II) nitrate and potassium sulfate separately in water. Mix the two solutions together; a white precipitate of lead(II) sulfate forms. Filter the mixture to collect the precipitate. Wash with distilled water. Leave to dry in a warm oven or on filter paper. (Do not heat directly as this may decompose the salt.)
Exam Tip
The method chosen to prepare a salt depends on the solubility of the reactants and the product: if one reactant is insoluble, use the excess solid method and filtration/crystallisation; if both reactants are soluble, use a titration; if the salt is insoluble, use a precipitation reaction from two solutions.