Stoichiometric calculations, in their basic form, assume ideal conditions where reactions go to completion, reactants are pure, and there are no side reactions. However, in real-world scenarios, these conditions are seldom met. Factors such as reaction conditions (temperature, pressure, catalysts) and the purity of reactants significantly impact the outcome of a reaction. For instance, impurities in reactants can reduce the yield, and non-ideal conditions can shift the equilibrium of reversible reactions, affecting the amount of product formed. Moreover, side reactions can consume some reactants, leading to lower yields of the desired product. To account for these real-world factors, chemists use the concept of percentage yield, which compares the actual yield obtained in an experiment to the theoretical yield predicted by stoichiometry. This adjustment provides a more realistic view of the reaction’s efficiency. In industrial processes, understanding these variations is crucial for process optimisation, cost estimation, and environmental impact assessment. Chemists must often perform additional calculations or experiments to determine the actual efficiency and yields under specific conditions.

Balancing equations for combustion reactions, particularly those involving hydrocarbons (compounds consisting of hydrogen and carbon), follows a systematic approach. First, write the unbalanced equation with the hydrocarbon and oxygen on the reactant side, and carbon dioxide and water on the product side. Start by balancing the carbon atoms, followed by hydrogen atoms, and finally, adjust the oxygen atoms. An essential tip is to balance the oxygen atoms last since they are in both CO₂ and H₂O. For instance, in the combustion of methane (CH₄), the unbalanced equation is CH₄ + O₂ → CO₂ + H₂O. Balance the carbons (1 C in CH₄ and 1 C in CO₂), then hydrogens (4 H in CH₄ and 2 H in H₂O, so put a 2 before H₂O), and finally oxygens (2 O in CO₂ and 4 O in 2H₂O, totalling 6 O atoms, thus 3 O₂ molecules). The balanced equation becomes CH₄ + 2O₂ → CO₂ + 2H₂O. This systematic approach ensures accuracy, especially in complex hydrocarbon combustion reactions.

The stoichiometric coefficient in a balanced chemical equation represents the relative number of moles of each reactant and product involved in the reaction. These coefficients are essential for understanding the proportions in which different substances react and are formed. For instance, in the equation 2H₂ + O₂ → 2H₂O, the coefficients 2, 1, and 2 indicate that two moles of hydrogen react with one mole of oxygen to produce two moles of water. In reaction calculations, these coefficients are used to determine the molar ratios of reactants and products, which are crucial for calculating the amounts of substances involved. For example, if you know the amount of one reactant, you can use the stoichiometric coefficients to find the amount of another reactant required or the amount of product that will be formed. This concept is fundamental in stoichiometry and is applied in a wide range of calculations including determining the limiting reactant, calculating theoretical yields, and even in more complex scenarios such as gas law calculations where volumes are directly proportional to moles.

In a chemical reaction, the limiting reactant is the substance that is completely consumed first, thus determining the amount of product formed. The excess reactant, on the other hand, is the substance that remains after the reaction has completed. When performing stoichiometric calculations, it is crucial to identify the limiting reactant as it dictates the maximum amount of product that can be formed. For example, if a reaction involves two reactants A and B in a 1:1 ratio, and there is more of reactant B present than required, reactant A becomes the limiting reactant. The amount of product formed is calculated based on the moles of reactant A. In practical scenarios, this concept is essential for optimising the efficiency and cost-effectiveness of chemical processes. For instance, in industrial settings, identifying and using the correct amount of limiting reactant can significantly reduce waste and increase yield, impacting both economic and environmental aspects of the process.

Stoichiometry plays a crucial role in environmental chemistry, especially in the field of pollution control. It is used to understand and quantify the reactions involved in the formation and breakdown of pollutants. For instance, stoichiometry is essential in calculating the amounts of reactants needed to neutralize pollutants in wastewater treatment or to reduce emissions in industrial processes. In the case of combustion reactions, stoichiometric calculations help in determining the optimal fuel-to-air ratio to minimise the production of pollutants like nitrogen oxides (NOx) and sulphur oxides (SOx). Additionally, these calculations are instrumental in designing processes for the removal of pollutants from the environment, such as in the scrubbing of flue gases to remove acidic gases. By understanding the exact stoichiometry of reactions involved in the formation and mitigation of pollutants, chemists can design more efficient and effective methods for pollution control, contributing significantly to environmental protection and sustainability. This application of stoichiometry demonstrates its importance not just in theoretical chemistry but also in practical, real-world scenarios where it can have a significant impact on environmental health.