Chemical reactions Revision Notes

The Mole as a Unit of Amount (Higher Tier)

The mole (symbol: mol) is the SI unit for the amount of substance. It provides a convenient way to count the enormous number of atoms, molecules, or ions in a given sample. One mole of any substance contains a specific number of particles, which is always the same, regardless of the substance. This concept is fundamental for converting between macroscopic masses and microscopic particle numbers.

The Avogadro Constant (Higher Tier)

The Avogadro constant, N₀ or L, represents the number of elementary entities (atoms, molecules, ions, etc.) in one mole of a substance. Its value is 6.022 × 10²³ mol⁻¹. This means that if you have 1 mol of carbon, you have 6.022 × 10²³ carbon atoms. If you have 1 mol of water molecules, you have 6.022 × 10²³ water molecules. It links the molar mass of a substance to the mass of a single particle.

Relating Moles, Particles, and Molar Mass (Higher Tier)

The number of particles in a sample can be calculated by multiplying the number of moles by the Avogadro constant. Conversely, the number of moles can be found by dividing the number of particles by the Avogadro constant. The molar mass (M) of a substance, measured in g/mol, is numerically equal to its relative atomic mass (Ar) or relative formula mass (Mr). This relationship is vital for converting between mass and moles in chemical calculations.

Practical Application (Higher Tier)

Understanding the mole and Avogadro constant is essential for stoichiometry. For instance, to find the number of atoms in a specific mass of an element, first convert the mass to moles using the molar mass, then multiply by the Avogadro constant. This allows chemists to predict the quantities of reactants and products in chemical reactions, ensuring efficient and accurate experimental work.

Understanding Ion Charges

Ions are atoms or groups of atoms that have gained or lost electrons, resulting in an electrical charge. Cations are positive ions (lost electrons), and anions are negative ions (gained electrons). For example, a sodium atom loses one electron to become Naⁱ⁷, and a chlorine atom gains one electron to become Clⁱ⁻. The charge of an ion is crucial for determining how it combines with other ions to form neutral compounds. Common charges to remember include Group 1 elements forming +1 ions, Group 2 forming +2 ions, Group 13 forming +3 ions, Group 15 forming -3 ions, Group 16 forming -2 ions, and Group 17 forming -1 ions.

The Principle of Charge Neutrality

When ions combine to form an ionic compound, the compound must be electrically neutral. This means the total positive charge from all cations must exactly cancel out the total negative charge from all anions. For instance, if you combine a cation with a +2 charge and an anion with a -1 charge, you will need two of the anions to balance the charge of one cation. This ensures that the resulting compound has no overall net charge, which is a fundamental rule for stable ionic compounds.

The 'Criss-Cross' Method for Deducing Formulas

The 'criss-cross' method is a simple technique to deduce chemical formulas. Write down the symbols of the ions with their charges. Then, take the numerical value of the charge of the cation and make it the subscript for the anion, and take the numerical value of the charge of the anion and make it the subscript for the cation. For example, for aluminium (Al³⁷) and oxide (O²⁻), the 3 from Al³⁷ becomes the subscript for O, and the 2 from O²⁻ becomes the subscript for Al, giving Al₂O₃. If the subscripts can be simplified (e.g., Mg²O² simplifies to MgO), always do so. Remember to use parentheses for polyatomic ions when a subscript is needed, like Ca(OH)₂.

Examples of Common Ion Formulas

Knowing common ion charges is key. Examples include: Sodium ion (Naⁱ⁷), Potassium ion (Kⁱ⁷), Magnesium ion (Mg²⁷), Calcium ion (Ca²⁷), Aluminium ion (Al³⁷), Chloride ion (Clⁱ⁻), Bromide ion (Brⁱ⁻), Oxide ion (O²⁻), Sulfide ion (S²⁻), Nitride ion (N³⁻). Polyatomic ions are groups of atoms that carry a charge: Sulfate (SO₄²⁻), Nitrate (NO₃ⁱ⁻), Carbonate (CO₃²⁻), Hydroxide (OHⁱ⁻), and Ammonium (NH₄ⁱ⁷). Practising combining these ions will help you confidently deduce compound formulas.

Electrolysis Basics: Electrodes and Ion Movement

Electrolysis is the process of using electrical energy to break down ionic compounds. It involves two electrodes connected to a power supply. Cations (positive ions) are attracted to the negatively charged electrode, called the cathode. Anions (negative ions) are attracted to the positively charged electrode, called the anode. Inert electrodes, such as graphite or platinum, are commonly used because they do not participate in the chemical reactions themselves but provide a surface for electron transfer.

Products Formed at the Cathode

At the cathode (negative electrode), reduction occurs as cations gain electrons. If the electrolyte contains metal ions, these ions will be discharged and deposited as solid metal if they are less reactive than hydrogen. If the metal is more reactive than hydrogen, or if the electrolyte is an aqueous solution, then hydrogen gas (H₂) is often produced instead, as hydrogen ions (H⁺) from water are discharged. For example, copper metal forms at the cathode during electrolysis of copper(II) sulfate solution, while hydrogen gas forms during the electrolysis of aqueous sodium chloride.

Products Formed at the Anode

At the anode (positive electrode), oxidation occurs as anions lose electrons. Non-metals are typically formed at the anode. For example, if the electrolyte contains halide ions (Cl⁻, Br⁻, I⁻), the corresponding halogen gas (chlorine, bromine, iodine) will be produced. If no halide ions are present, and the electrolyte is an aqueous solution, then oxygen gas (O₂) is usually formed from the hydroxide ions (OH⁻) present in water, or from the oxidation of water itself. For instance, in the electrolysis of molten lead bromide, bromine gas forms at the anode.

Understanding Stoichiometry from Masses (Higher)

Stoichiometry is about the quantitative relationships between reactants and products in a chemical reaction. It's built upon the Law of Conservation of Mass, meaning the total mass of reactants must equal the total mass of products. To use stoichiometry with masses, you need a balanced chemical equation. The coefficients in the balanced equation represent the molar ratios. For example, in 2Na + Cl₂ → 2NaCl, two moles of sodium react with one mole of chlorine to produce two moles of sodium chloride. To calculate masses: convert given masses to moles using (moles = mass / Mr), use the molar ratios from the balanced equation to find moles of the desired substance, then convert these moles back to mass (mass = moles × Mr).

Identifying Limiting Reactants (Higher)

A limiting reactant is the substance that gets completely used up first in a chemical reaction. It determines the maximum amount of product that can be formed, as the reaction stops once it's gone. Any other reactant still remaining is called the excess reactant. To identify the limiting reactant: first, balance the chemical equation. Then, calculate the moles of each reactant available using their given masses and molar masses. Next, compare the actual mole ratio of reactants to the required mole ratio from the balanced equation. The reactant that would be fully consumed first, given the available amounts and reaction stoichiometry, is the limiting reactant.

Calculations Involving Limiting Reactants (Higher)

Once the limiting reactant is identified, all subsequent calculations for product yield or amount of excess reactant consumed must be based on the amount of the limiting reactant. The theoretical yield of a product is the maximum amount that can be formed from the given reactants, assuming perfect conditions and 100% reaction completion. To calculate theoretical yield: determine the moles of the limiting reactant, use the molar ratio from the balanced equation to find the moles of the product, and then convert these moles back to mass. The amount of excess reactant consumed can also be calculated in a similar way, based on the limiting reactant, to find out how much of the excess reactant is left over.