Chemistry Study Notes: Acids, Bases, and Salt Preparations
Welcome to one of the most fundamental and important chapters in Inorganic Chemistry! Acids, bases, and salts are everywhere—from the citric acid in lemons to the soap you use every day. Understanding how these chemicals interact is essential for your examinations.
Let us break down the definitions, reactions, and salt preparations into simple, easy-to-follow steps!
Section I: Defining Acids, Bases, and Alkalis
What are Acids and Bases?
In the Pearson Edexcel IGCSE specification, acid-base reactions are defined in terms of proton transfer:
- An acid is a proton donor. In aqueous solution, acids dissociate to produce hydrogen ions (\(H^+\), which are protons).
- A base is a proton acceptor. Bases react with acids to neutralise them. Metal oxides, metal hydroxides, and ammonia (\(NH_3\)) act as bases.
Key Examples of Acids:
- Hydrochloric Acid (\(HCl\)) — forms chloride salts
- Sulfuric Acid (\(H_2SO_4\)) — forms sulfate salts
- Nitric Acid (\(HNO_3\)) — forms nitrate salts
What is an Alkali?
An alkali is a water-soluble base:
- Alkalis dissolve in water to produce hydroxide ions (\(OH^-\)).
- Common alkalis include sodium hydroxide (\(NaOH\)), potassium hydroxide (\(KOH\)), and aqueous ammonia (\(NH_3(aq)\)).
The pH Scale and Indicators
The pH scale runs from 0 to 14 and classifies solutions as:
- 0–3: Strongly acidic
- 4–6: Weakly acidic
- 7: Neutral
- 8–10: Weakly alkaline
- 11–14: Strongly alkaline
Important Indicators to Know:
- Universal Indicator: Used to measure approximate pH across the full range (red in strong acid, green at neutral, purple/violet in strong alkali).
- Litmus: Red in acid, blue in alkali.
- Methyl Orange: Red in acid, yellow in alkali.
- Phenolphthalein: Colourless in acid, pink in alkali.
Key Takeaway: An acid is a proton (\(H^+\)) donor, and a base is a proton acceptor. Alkalis are water-soluble bases that release \(OH^-\) ions.
Section II: Key Reactions of Acids
Acids react with metals, bases, and metal carbonates to form salts.
1. Reaction with Bases (Neutralisation)
Acid + Base (or Alkali) \(\rightarrow\) Salt + Water
Example:
\(HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)\)
Ionic Equation for Neutralisation:
\(H^+(aq) + OH^-(aq) \rightarrow H_2O(l)\)
2. Reaction with Metals
Dilute hydrochloric acid and sulfuric acid react with reactive metals (e.g. magnesium, zinc, iron) to form a salt and hydrogen gas. (Note: reactions of metals with nitric acid are excluded from your syllabus).
Acid + Metal \(\rightarrow\) Salt + Hydrogen
Example:
\(Zn(s) + H_2SO_4(aq) \rightarrow ZnSO_4(aq) + H_2(g)\)
3. Reaction with Metal Carbonates
Acid + Metal Carbonate \(\rightarrow\) Salt + Water + Carbon Dioxide
Example:
\(2HCl(aq) + CaCO_3(s) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)\)
Testing for Carbon Dioxide: Bubble the gas through limewater (aqueous calcium hydroxide); it turns cloudy/milky due to the formation of a white precipitate of calcium carbonate.
Section III: Preparation of Salts
A salt is a compound formed when the hydrogen ion (\(H^+\)) of an acid is replaced by a metal ion or an ammonium ion (\(NH_4^+\)).
Solubility Rules (Must Be Memorised)
- Soluble:
- All common compounds of sodium, potassium, and ammonium
- All nitrates
- Common chlorides, except silver chloride and lead(II) chloride
- Common sulfates, except barium sulfate, calcium sulfate, and lead(II) sulfate
- Insoluble:
- Common carbonates, except sodium carbonate, potassium carbonate, and ammonium carbonate
- Common hydroxides, except sodium hydroxide, potassium hydroxide, and ammonium hydroxide (calcium hydroxide is slightly soluble)
Method 1: Preparing a Soluble Salt from an Insoluble Reactant
Used when making a soluble salt starting from an insoluble base (metal oxide/hydroxide) or metal carbonate (e.g. preparing hydrated copper(II) sulfate crystals from copper(II) oxide and sulfuric acid).
- Measure a set volume of acid into a beaker and warm it gently.
- Add the insoluble reactant in excess, stirring until no more dissolves (this ensures all acid is neutralised).
- Filter the mixture to remove the unreacted excess insoluble solid. The filtrate contains only the aqueous salt and water.
- Heat the filtrate gently in an evaporating basin to evaporate water until the crystallisation point is reached (crystals form on a cold glass rod).
- Leave the saturated solution to cool and crystallise slowly.
- Filter off the crystals and dry them (e.g. by patting between pieces of filter paper or in a warm desiccator).
Method 2: Preparing a Soluble Salt from an Acid and an Alkali (Titration)
Used when both starting reactants are soluble (e.g. making sodium, potassium, or ammonium salts):
- Use a pipette to measure an exact volume of alkali into a conical flask and add a few drops of indicator (e.g. methyl orange or phenolphthalein).
- Add acid from a burette slowly until the indicator changes colour (the end-point). Record the volume of acid added (the titre).
- Repeat the experiment using the exact same volumes of acid and alkali, but without indicator, to obtain an uncontaminated salt solution.
- Heat the solution to the crystallisation point, allow crystals to form, filter, and dry.
Method 3: Preparing an Insoluble Salt from Two Soluble Reactants (Precipitation)
Used to make an insoluble salt (e.g. pure, dry lead(II) sulfate or barium sulfate):
- Mix two aqueous solutions containing the required ions (e.g. lead(II) nitrate solution + sodium sulfate solution).
- A precipitate forms immediately: \(Pb(NO_3)_2(aq) + Na_2SO_4(aq) \rightarrow PbSO_4(s) + 2NaNO_3(aq)\).
- Filter the mixture to collect the insoluble salt precipitate on the filter paper.
- Wash the precipitate thoroughly with distilled water to remove traces of the soluble salt solution.
- Dry the precipitate in a warm oven or warm place.
Key Takeaway: Choose the preparation method based on solubility: use excess insoluble base for standard soluble salts, titration for soluble salts containing \(Na^+\), \(K^+\), or \(NH_4^+\), and precipitation for insoluble salts.