Chemical reactions Revision Notes

Concentration vs. Strength (Higher Tier)

It is crucial to differentiate between an acid's concentration and its strength. Concentration describes the amount of acid solute dissolved per unit volume of solution. A concentrated acid has a high amount of acid dissolved, while a dilute acid has a low amount. This is independent of how the acid molecules behave in water. For example, a 1 M solution of hydrochloric acid is concentrated, and a 0.001 M solution is dilute. The term strength, however, refers to the extent of ionisation (dissociation) of the acid in water. A strong acid completely ionises to produce H⁺ ions, whereas a weak acid only partially ionises. These two properties are distinct: you can have a dilute solution of a strong acid or a concentrated solution of a weak acid.

Strong Acids (Higher Tier)

A strong acid is defined by its complete ionisation in an aqueous solution. When a strong acid dissolves in water, virtually all its molecules dissociate into hydrogen ions (H⁺) and their corresponding anion. This process is essentially irreversible and is represented by a single arrow in chemical equations. For example, hydrochloric acid (HCl), nitric acid (HNO₃), and sulfuric acid (H₂SO₄) are strong acids. Their complete dissociation means they readily contribute a high concentration of H⁺ ions to the solution, even in relatively dilute solutions. Because of this, strong acids are highly corrosive and reactive.

Weak Acids (Higher Tier)

In contrast to strong acids, a weak acid only partially ionises in an aqueous solution. This means that when a weak acid dissolves in water, only a small fraction of its molecules donate a proton (H⁺ ion), while the majority remain in their undissociated molecular form. The ionisation of a weak acid is a reversible process, establishing an equilibrium between the undissociated molecules and the ions. This is indicated by a reversible arrow (⇌) in chemical equations. Examples include ethanoic acid (CH₃COOH), carbonic acid (H₂CO₃), and citric acid. The limited production of H⁺ ions means that, for a given concentration, a weak acid solution will have a higher pH than a strong acid solution of the same concentration.

Impact on pH (Higher Tier)

The pH of an acid solution is directly related to the concentration of hydrogen ions [H⁺]. For strong acids, because they ionise completely, the [H⁺] is approximately equal to the initial concentration of the acid. For example, a 0.1 mol/dm³ solution of HCl will have a [H⁺] of 0.1 mol/dm³, resulting in a pH of 1. For weak acids, however, because they only partially ionise, the [H⁺] is significantly lower than the initial acid concentration. Consequently, a 0.1 mol/dm³ solution of ethanoic acid (a weak acid) will have a [H⁺] much less than 0.1 mol/dm³, leading to a pH value typically around 3, which is considerably higher than pH 1, despite having the same initial concentration. This illustrates how strength (degree of ionisation) profoundly impacts pH, independently of nominal concentration.

What is Neutralisation?

Neutralisation is a chemical reaction that occurs when an acid reacts with an alkali or a base. The key products of this reaction are always a salt and water. This process results in the acidic and alkaline properties of the reactants being cancelled out, bringing the pH of the solution closer to 7, which is neutral. For example, hydrochloric acid reacting with sodium hydroxide forms sodium chloride (a salt) and water.

Acids, Alkalis, and Bases in Neutralisation

Acids are substances that donate hydrogen ions (H⁺) when dissolved in water. Alkalis are soluble bases that produce hydroxide ions (OH⁻) when dissolved in water. Bases are substances that accept hydrogen ions; they can be soluble (alkalis) or insoluble (e.g., metal oxides, metal hydroxides). In a neutralisation reaction, the H⁺ ions from the acid combine with the OH⁻ ions from the alkali or base to form water (H₂O).

Example Reactions and Products

The general word equation for neutralisation is: Acid + Alkali/Base → Salt + Water. The specific salt formed depends on the acid and the alkali/base used. For example:

• Hydrochloric acid + Sodium hydroxide → Sodium chloride + Water
• Sulfuric acid + Potassium hydroxide → Potassium sulfate + Water
• Nitric acid + Calcium carbonate (a base) → Calcium nitrate + Water + Carbon dioxide (note: carbonate reactions also produce CO2).

The name of the salt comes from the metal in the alkali/base and the acid radical (e.g., chloride from hydrochloric acid, sulfate from sulfuric acid, nitrate from nitric acid).

The Role of Water in Neutralisation

The formation of water is the defining feature of all acid-base neutralisation reactions. Water is formed from the reaction of hydrogen ions (H⁺) from the acid and hydroxide ions (OH⁻) from the alkali or suitable base. The ionic equation H⁺(aq) + OH⁻(aq) → H₂O(l) represents this fundamental process, illustrating why neutralisation always produces water alongside a salt.

Chemical Symbols for Elements

Every element has a unique chemical symbol. This symbol consists of one or two letters. The first letter is always capitalised, and the second letter (if present) is always lowercase. For example, 'H' for Hydrogen, 'O' for Oxygen, 'Na' for Sodium, 'Cl' for Chlorine, and 'Fe' for Iron. These symbols are universally recognised in chemistry and are fundamental for writing chemical formulae. Knowing common element symbols is crucial for correctly representing substances.

Formulae of Elements

Some elements exist as single atoms (monatomic) in their stable state, and their formula is simply their chemical symbol, e.g., He (Helium), Ne (Neon), Ar (Argon), Fe (Iron), Cu (Copper). Other elements exist as molecules made of two or more atoms of the same element bonded together. For example, hydrogen, oxygen, nitrogen, fluorine, chlorine, bromine, and iodine all exist as diatomic molecules (two atoms per molecule) and their formulae are H₂, O₂, N₂, F₂, Cl₂, Br₂, and I₂ respectively. Sulfur commonly exists as S₈ molecules.

Formulae of Simple Covalent Compounds

In simple covalent compounds, atoms share electrons to form molecules. The chemical formula represents the exact number of atoms of each element present in one molecule. To write these formulae, you need to know the symbols of the constituent elements and the number of atoms of each present. Subscripts are used to indicate the number of atoms. For example, water is H₂O (two hydrogen, one oxygen); carbon dioxide is CO₂ (one carbon, two oxygen); methane is CH₄ (one carbon, four hydrogen); ammonia is NH₃ (one nitrogen, three hydrogen).

Formulae of Simple Ionic Compounds

Ionic compounds are formed when atoms transfer electrons to form ions, which then attract electrostatically. The formula for an ionic compound represents the simplest whole-number ratio of positive and negative ions required to achieve overall electrical neutrality. This means the total positive charge must equal the total negative charge. For example, in sodium chloride (NaCl), one Na⁺ ion combines with one Cl⁻ ion. In magnesium chloride (MgCl₂), the Mg²⁺ ion needs two Cl⁻ ions to balance the charge. In aluminium oxide (Al₂O₃), two Al³⁺ ions (total +6) combine with three O²⁻ ions (total -6). Polyatomic ions (like SO₄²⁻) retain their identity; if more than one is needed, enclose it in brackets, e.g., Ca(OH)₂.

What are Ionic Equations? (Higher)

An ionic equation shows only the ions and molecules directly involved in a chemical reaction, particularly those occurring in aqueous solutions. Unlike a full symbol equation, it omits 'spectator ions' – ions that do not participate in the chemical change and remain in solution before and after the reaction. This makes ionic equations more concise and highlights the essential chemical transformation. They are crucial for representing reactions like precipitation, acid-base neutralisation, and some redox processes.

Steps to Construct a Balanced Ionic Equation (Higher)

1. Write the balanced symbol equation for the reaction, including state symbols (s, l, g, aq).
2. Dissociate all soluble ionic compounds (marked 'aq') into their constituent free ions. For example, NaCl(aq) becomes Na⁺(aq) + Cl⁻(aq).
3. Keep insoluble precipitates (s), liquids (l), and gases (g) in their molecular or empirical formula form. For instance, H₂O(l) remains H₂O(l) and a precipitate like AgCl(s) remains AgCl(s).
4. Identify and cancel out spectator ions. These are ions that appear on both sides of the equation with the same state, charge, and number.
5. Write the remaining species as the balanced ionic equation.
6. Ensure the final equation is balanced for both atoms and total charge on both sides. The sum of charges on the reactant side must equal the sum of charges on the product side.

Identifying Spectator Ions (Higher)

Spectator ions are present in the reaction mixture but do not undergo any chemical change. They are 'spectators' to the reaction, simply watching the other ions react. In an ionic equation, spectator ions are aqueous ions that appear identically on both the reactant and product sides of the equation before cancellation. Common spectator ions often include alkali metal ions (Na⁺, K⁺), nitrate ions (NO₃⁻), and halide ions (Cl⁻, Br⁻, I⁻) unless they are involved in forming a precipitate.

Example: Precipitation of Silver Chloride (Higher)

Consider the reaction between silver nitrate (AgNO₃) and sodium chloride (NaCl) to form silver chloride (AgCl) precipitate and sodium nitrate (NaNO₃).

1. Balanced symbol equation: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)
2. Dissociate aqueous compounds: Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq)
3. Identify spectator ions: Na⁺(aq) and NO₃⁻(aq) appear on both sides.
4. Cancel spectator ions: Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

This is the balanced ionic equation. Notice that the charges (+1 + -1 = 0) on the left balance the neutral compound (0) on the right.