Chemical reactions Revision Notes

The Law of Conservation of Mass

The law of conservation of mass is a fundamental principle in chemistry. It states that during any chemical reaction, atoms are neither created nor destroyed. This means that the total mass of the reactants (the substances that react) must be equal to the total mass of the products (the substances formed). This law is essential for understanding stoichiometry and balancing chemical equations, as it implies that the number of atoms of each element remains constant throughout a reaction, only their arrangement changes.

Apparent Mass Changes in Open Systems

In some chemical reactions carried out in an open system, the observed mass may appear to change. This is often due to the involvement of gases. For example, if a gas is produced during a reaction (e.g., combustion or the reaction of an acid with a carbonate) and allowed to escape, the measured mass of the remaining substances will decrease. Conversely, if a gas from the surroundings, such as oxygen, combines with a solid reactant (e.g., in oxidation or burning) and is incorporated into the products, the measured mass of the solid products will increase. These are not true violations of the law of conservation of mass, but rather a reflection of substances entering or leaving the system.

Closed Systems and True Conservation (Higher)

To demonstrate the law of conservation of mass unambiguously, chemical reactions should ideally be carried out in a closed system. A closed system prevents any reactants or products from escaping or entering the reaction vessel. If a gas is produced, it remains within the sealed container, and if a gas from the surroundings is consumed, it must have been present within the sealed container initially. In such a system, precise measurements would show that the total mass of the system before the reaction is exactly the same as the total mass after the reaction, confirming that mass is conserved.

Calculations Involving Mass Changes (Higher)

When a mass change is observed in an open system, it indicates that a gaseous reactant or product has been involved. For example, if magnesium burns in air to form magnesium oxide, the mass of the solid increases because oxygen from the air has combined with the magnesium. The mass increase directly corresponds to the mass of oxygen that reacted. Similarly, if calcium carbonate decomposes to form calcium oxide and carbon dioxide, the mass of the solid decreases; this decrease is equal to the mass of carbon dioxide gas that escaped. These changes allow for calculations of the mass of the gas involved, based on the conservation of mass principle.

Calculating Moles from Mass (Higher)

To work with stoichiometry from masses, the first crucial step is to convert the given masses of reactants and products into moles. The formula for this conversion is: moles = mass (in grams) / molar mass (in g/mol). The molar mass (Mᵣ) of a compound is calculated by summing the relative atomic masses (Aᵣ) of all the atoms in its chemical formula. For an element, the molar mass is numerically equal to its relative atomic mass. For instance, to find the moles of 10 g of H₂O, you would first calculate its molar mass: (2 × 1.0) + 16.0 = 18.0 g/mol. Then, moles = 10 g / 18.0 g/mol ≈ 0.556 mol.

Deducing Stoichiometry from Reacting Masses (Higher)

Once you have converted the masses of reactants and products into moles, you can deduce the stoichiometry of the chemical equation. The stoichiometric coefficients represent the simplest whole-number ratio of moles of reactants and products. To find this ratio, divide the number of moles of each substance by the smallest number of moles calculated. If the resulting numbers are not whole numbers, you may need to multiply all the ratios by a common factor to obtain the simplest whole-number ratio. This ratio then gives you the coefficients for your balanced chemical equation. For example, if moles are 0.5 mol A, 1.5 mol B, and 1.0 mol C, dividing by 0.5 gives a ratio of 1:3:2.

Identifying Limiting Reactants (Higher)

A limiting reactant is the substance that is completely consumed during a chemical reaction, thereby dictating the maximum amount of product that can be formed. The other reactant(s) are in excess. To identify the limiting reactant, you need to compare the actual mole ratio of reactants given with the theoretical mole ratio required by the balanced chemical equation. Calculate the moles of each reactant. Then, divide the moles of each reactant by its stoichiometric coefficient from the balanced equation. The reactant with the smallest resulting value is the limiting one. This method helps determine which reactant will run out first.

Impact of Limiting Reactants on Yield (Higher)

The amount of product formed in a reaction is directly determined by the quantity of the limiting reactant. Even if a large excess of another reactant is present, the reaction stops once the limiting reactant is used up. Therefore, to calculate the theoretical yield of a product, you must base your calculations on the moles of the limiting reactant. In industrial chemistry, controlling limiting reactants is vital for economic efficiency; ensuring the more expensive or harder-to-obtain reactant is limiting minimises waste and maximises the yield of desired products for a given investment in raw materials.

Exothermic Reactions: Releasing Heat

Exothermic reactions are chemical processes that release energy into their surroundings, typically as heat. This results in an increase in the temperature of the surroundings. The word 'exo' means 'out' or 'exit', indicating energy moving out of the system. Examples include burning fuels (combustion), respiration, and many neutralisation reactions. When you feel warmth from a reaction, it is exothermic. In terms of energy, the products of an exothermic reaction have less chemical potential energy than the reactants, with the difference being released as heat.

Endothermic Reactions: Absorbing Heat

Endothermic reactions are chemical processes that absorb energy from their surroundings, typically as heat. This causes a decrease in the temperature of the surroundings. The word 'endo' means 'in' or 'into', indicating energy moving into the system. Examples include instant cold packs, photosynthesis, and the thermal decomposition of calcium carbonate. When a reaction feels cold to the touch, it is endothermic. In terms of energy, the products of an endothermic reaction have more chemical potential energy than the reactants, as energy from the surroundings has been absorbed and stored.

Measuring Temperature Changes (PAG C8)

To distinguish between endothermic and exothermic reactions experimentally, we measure the temperature change of the reaction's surroundings. A thermometer is used to record the initial temperature of the reactants and the final temperature after the reaction has occurred. For accuracy, reactions should ideally be carried out in an insulated container like a polystyrene cup to minimise heat exchange with the external environment. A rise in temperature indicates an exothermic reaction, while a fall in temperature indicates an endothermic reaction. The independent variable is the specific reaction being investigated, the dependent variable is the temperature change, and control variables include initial temperature, volume/mass of reactants, and insulation.

Key Characteristics and Examples

Exothermic reactions: Temperature of surroundings increases. Energy is released (e.g., heat, light). Examples: combustion, respiration, neutralisation. Products are more stable (lower energy) than reactants. Feel warm/hot. Endothermic reactions: Temperature of surroundings decreases. Energy is absorbed (e.g., heat). Examples: instant cold packs, photosynthesis, thermal decomposition. Products are less stable (higher energy) than reactants. Have a cooling effect/feel cold.

What are Acids and Their Ions?

Acids are chemical substances that, when dissolved in water, produce hydrogen ions (H⁺). These hydrogen ions are responsible for the acidic properties of the solution. For instance, hydrochloric acid (HCl) in water dissociates into H⁺ and Cl⁻ ions. The higher the concentration of H⁺ ions, the stronger the acid solution. Acids typically have a pH value less than 7.

Understanding Alkalis and Their Ions

Alkalis are soluble bases that, when dissolved in water, release hydroxide ions (OH⁻). These hydroxide ions are the key to the alkaline properties of the solution. For example, sodium hydroxide (NaOH) dissolves in water to form Na⁺ and OH⁻ ions. A higher concentration of OH⁻ ions indicates a stronger alkali solution. Alkalis typically have a pH value greater than 7.

The Role of Water in Ion Formation

The process of dissolving in water is crucial for both acids and alkalis to exhibit their characteristic properties. Water molecules help to separate the acid or alkali compound into its respective ions. Without water, many acids and alkalis would not fully dissociate, and their acidic or alkaline properties would not be as apparent or potent.