The characteristic properties of acids and bases

Acids are substances that form hydrogen ions (H⁺) when dissolved in water. Common laboratory acids include hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and nitric acid (HNO₃). They are corrosive, taste sour, and have a pH less than 7. Alkalis are bases that are soluble in water. They form hydroxide ions (OH⁻) in solution. Common alkalis include sodium hydroxide (NaOH), potassium hydroxide (KOH), and aqueous ammonia (NH₃). They feel soapy, are corrosive, and have a pH greater than 7. Indicators are substances that change colour depending on the pH. Litmus paper turns red in acid and blue in alkali. Methyl orange turns red in acid and yellow in alkali. The pH scale, from 0 to 14, measures acidity and alkalinity. A pH of 7 is neutral (like pure water), values below 7 are acidic (the lower the number, the more acidic), and values above 7 are alkaline (the higher the number, the more alkaline). Universal indicator has a range of colours corresponding to the entire pH scale, from red for strong acids to dark purple for strong alkalis.

Acids take part in several characteristic reactions, typically producing a salt. A salt is an ionic compound formed when the hydrogen of an acid is replaced by a metal or ammonium ion. The name of the salt depends on the acid used: hydrochloric acid produces chlorides, sulfuric acid produces sulfates, and nitric acid produces nitrates. There are three key general reactions to learn.

The strength of an acid or alkali is not the same as its concentration. Strength refers to the degree of ionisation (or dissociation) in water. A strong acid, like hydrochloric acid (HCl), fully ionises in water. Every HCl molecule splits into H⁺ and Cl⁻ ions. A weak acid, like ethanoic acid (CH₃COOH), only partially ionises. Only a small fraction of the molecules split into ions, so the solution contains mostly un-ionised molecules and a low concentration of H⁺ ions. This is shown with a reversible arrow: CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq). Similarly, a strong alkali like sodium hydroxide (NaOH) fully dissociates into Na⁺ and OH⁻ ions. A weak alkali like ammonia (NH₃) only reacts partially with water to form a low concentration of OH⁻ ions: NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq). Because strong acids/alkalis produce a higher concentration of mobile ions, their solutions have higher electrical conductivity and a more extreme pH (lower for acids, higher for alkalis) than weak acids/alkalis of the same molar concentration.

Oxides are compounds containing oxygen and at least one other element. They can be classified into four types based on their acidic and basic properties.

1. Basic Oxides: These are typically oxides of metals. They react with acids to form a salt and water. Most are insoluble (e.g., copper(II) oxide, CuO), but those of Group 1 and 2 metals are soluble and form alkaline solutions (e.g., sodium oxide, Na₂O).
2. Acidic Oxides: These are typically oxides of non-metals. They react with bases/alkalis to form a salt and water. Examples include carbon dioxide (CO₂) and sulfur dioxide (SO₂).
3. Amphoteric Oxides (Extended): These are special metal oxides that show both acidic and basic properties. They react with both acids and strong alkalis to form a salt and water. Common examples are zinc oxide (ZnO) and aluminium oxide (Al₂O₃). For example: ZnO + 2HCl → ZnCl₂ + H₂O (acting as a base) and ZnO + 2NaOH → Na₂ZnO₂ + H₂O (acting as an acid).
4. Neutral Oxides (Extended): These are non-metal oxides that do not react with either acids or bases. Examples include carbon monoxide (CO), nitrogen(I) oxide (N₂O), and water (H₂O).

Neutralisation is the reaction between an acid and a base to form a salt and water. The essential reaction involves hydrogen ions from the acid reacting with hydroxide ions from the alkali to form neutral water. The ionic equation for this is H⁺(aq) + OH⁻(aq) → H₂O(l). This principle is used to prepare soluble salts. There are two main methods. Method 1: Reacting an Acid with an Insoluble Solid. This is used for making salts from an acid and an insoluble metal, base, or carbonate. Step 1: Add the insoluble solid in excess (a little at a time) to a warm acid until no more reacts. This ensures all the acid is used up. Step 2: Filter the mixture to remove the unreacted excess solid. The filtrate is the salt solution. Step 3: Gently heat the filtrate to evaporate some of the water and form a saturated solution. Step 4: Leave the solution to cool, allowing crystals of the salt to form. Step 5: Filter off the crystals, wash them with a small amount of cold distilled water, and dry them. Method 2: Titration. This is used when both the acid and the base (an alkali) are soluble. Step 1: Use a pipette to measure a known volume of alkali into a conical flask and add a few drops of indicator. Step 2: Add acid from a burette, swirling the flask, until the indicator just changes colour (the end point). Record the volume of acid added. Step 3: Repeat the experiment using the exact same volumes of acid and alkali but without the indicator to get a pure salt solution. Step 4: Crystallise the salt from the pure solution as described in Method 1.

Insoluble salts cannot be made by crystallisation because they do not stay dissolved. Instead, they are made by a precipitation reaction. This involves mixing two different soluble salt solutions. One solution provides the required positive ion (cation) and the other provides the required negative ion (anion). When the two solutions are mixed, the cation and anion combine to form the insoluble salt, which appears as a solid precipitate. The other ions remain dissolved in the solution; these are called spectator ions because they do not take part in the reaction. The method is simple:

1. Mix the two chosen soluble salt solutions.
2. Filter the mixture to separate the solid precipitate from the solution.
3. Wash the precipitate on the filter paper with distilled water to remove any soluble impurities.
4. Dry the precipitate in a warm oven or by leaving it in a warm place.

Titration is a precise experimental technique used to determine the unknown concentration of a solution (the analyte) by reacting it with a solution of known concentration (the standard solution). The key is to find the exact volume of the standard solution needed to completely neutralise the analyte. This is called the equivalence point, which is observed using an indicator. The calculation involves three main steps:

1. Write a balanced chemical equation for the reaction to find the mole ratio (n) of acid to alkali.
2. Calculate the moles of the solution with the known concentration and volume using the formula: Moles = Concentration × Volume (in dm³). Remember to convert volumes in cm³ to dm³ by dividing by 1000.
3. Use the mole ratio from the balanced equation to determine the moles of the unknown solution. Then, calculate its concentration using the formula: Concentration = Moles / Volume (in dm³). Alternatively, the combined formula C₁V₁/n₁ = C₂V₂/n₂ can be used, where C is concentration, V is volume, and n is the mole ratio for substance 1 (acid) and substance 2 (alkali).