Molar Volume, Concentration & Stoichiometry

The molar volume of any gas at Room Temperature and Pressure (RTP — 20°C and 1 atm) is 24 dm³/mol. This value applies to all ideal gases regardless of their identity.

It allows chemists to convert between the volume of a gas and the number of moles without needing to measure mass directly. This is especially useful in stoichiometric calculations for gaseous reactions, where volume is easier to measure than mass.

Avogadro’s Law states that equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules. Molar volume is the direct application of this law — it is the volume occupied by exactly one mole (6.02 × 10²³ molecules) of any gas at a given temperature and pressure.

Because empirical formulas represent only the simplest whole-number ratio of atoms. Two compounds can share the same ratio but have different actual numbers of atoms in their molecules. For example, ethene (C₂H₄) and propene (C₃H₆) both have the empirical formula CH₂, yet they are entirely different compounds.

Percentage yield measures the efficiency of a chemical reaction. It helps compare actual production with the theoretically possible amount, identify losses in a process, and evaluate cost-effectiveness. In industrial chemistry, even a small improvement in yield can result in enormous economic savings.

Molar mass of H₂SO₄ = 2(1) + 32 + 4(16) = 98 g/mol. Moles = 49 ÷ 98 = 0.5 mol. Concentration = 0.5 mol ÷ 1 dm³ = 0.5 mol/dm³.

First, write the balanced chemical equation for the reaction. Then calculate the number of moles of each reactant using the formula: moles = mass ÷ molar mass. Divide the moles of each reactant by its stoichiometric coefficient from the balanced equation. The reactant that gives the smallest quotient is the limiting reactant — it will be completely consumed first and determines the maximum yield of product.

The purity of a medicinal compound is critical because impurities can cause toxic side effects, trigger allergic reactions, or interact negatively with other drugs. Additionally, the dosage prescribed by a doctor assumes a certain purity — if the active ingredient is diluted by impurities, the patient receives less medicine than intended, reducing effectiveness. Regulatory bodies such as WHO set strict purity standards for pharmaceutical compounds to ensure patient safety.

Yes, the concept of a limiting reactant still applies to reversible reactions at the start of the reaction. However, because the reaction does not go to completion — it reaches a state of dynamic equilibrium — both reactants and products remain present at equilibrium. Therefore, while one reactant is initially “limiting,” neither is entirely consumed. The limiting reactant concept applies strictly in the context of irreversible or complete reactions.

To check whether a formula is empirical, find the greatest common divisor (GCD) of all the subscripts in the formula. If the GCD is 1 (i.e., no subscript can be divided further), the formula is empirical. If the GCD is greater than 1, the formula is molecular and can be simplified — for example, C₆H₁₂O₆ has GCD 6, so its empirical formula is CH₂O.

Mass spectrometry is a direct method — the compound is vaporised, ionised, and the mass-to-charge ratio of the molecular ion peak (M⁺) gives the molar mass. Alternatively, colligative property methods such as freezing point depression or boiling point elevation can be used: by dissolving a known mass of the compound in a solvent and measuring the change in freezing or boiling point, the molar mass can be calculated using the formula: M = (mass of solute × K) ÷ (change in temperature × mass of solvent).