Ethylene \(\left(\mathrm{C}_{2} \mathrm{H}_{4}\right),\) an important industrial organic chemical, can be prepared by heating hexane \(\left(\mathrm{C}_{6} \mathrm{H}_{14}\right)\) at \(800^{\circ} \mathrm{C}:$$$\mathrm{C}_{6} \mathrm{H}_{14} \longrightarrow \mathrm{C}_{2} \mathrm{H}_{4}+\text { other products }$$ If the yield of ethylene production is 42.5 percent, what mass of hexane must be reacted to produce \)481 \mathrm{~g}$ of ethylene?

Stoichiometry is like a recipe for chemical reactions. It involves calculations that help us understand the relationships between the reactants and products in a reaction. In essence, stoichiometry provides the proportions needed to make sure that all reactants turn into products effectively, without any waste.

Consider a simple analogy: Baking a cake requires a specific ratio of eggs to flour. Too many or too few eggs, and the cake won't turn out right. Stoichiometry ensures that in a chemical reaction, we combine elements in the right 'ratios' based on their molecular weights and the coefficients from the balanced chemical equation.

While working through a problem, clear identification of the needed moles of reactants and products is crucial. The balanced equation gives us the 'recipe', and stoichiometry is the method by which we apply it to compute the quantities required for the reaction.

The concept of molar mass is a cornerstone of chemical calculations. Think of it as the 'weight' of one mole of a substance. One mole contains Avogadro's number of particles, which is roughly 6.022 x 10²³ particles. The molar mass tells us how much one mole of a substance weighs in grams.

To calculate molar mass, we add up the atomic masses of every element that makes up a chemical compound. For example, water (H₂O) has a molar mass equal to twice the atomic mass of hydrogen plus the atomic mass of oxygen. When solving problems, knowing the molar mass allows us to move between the mass of a substance and the number of moles, aiding us in many types of yield calculations.

The theoretical yield is the amount of product that should be produced in a perfect chemical reaction, where every molecule of reactants fully converts to products. It is a 'best case' scenario, where nothing is lost or wasted.

To calculate it, we must first know the balanced chemical reaction and the molar amounts of reactants we plan to use. The reactant that runs out first, known as the limiting reactant, determines the maximum amount of product we can make. We then use stoichiometry to find the theoretical yield starting from that limiting reactant. Understanding theoretical yield helps us set a benchmark for what we could achieve under ideal conditions in laboratory or industrial processes.

In a real-world laboratory setting, reactions often yield less product than the theoretical calculations predict. This could be due to side reactions, loss of product during transfer, or incomplete reactions. To understand how efficient a reaction is, we use the concept of percent yield.

The percent yield is found by taking the actual yield (what was actually produced in the lab) and dividing it by the theoretical yield, then multiplying by 100 to get a percentage. This tells us what fraction of the theoretical yield we actually obtained. It is a practical measure of a reaction's efficiency and is crucial for evaluating the success of laboratory procedures and industrial processes.

A chemical reaction is a process where substances, known as reactants, transform into new substances called products. It's like cooking, where ingredients come together to create a new dish.

The reaction can be represented symbolically using a chemical equation that shows the starting materials and the end products, along with their respective quantities. For example, burning propane in a grill is a chemical reaction where propane and oxygen react to form carbon dioxide and water.

In every chemical reaction, conservation of mass applies, meaning the mass of the products equals the mass of the reactants. In our stoichiometry problem, the chemical equation helped us understand what happens when hexane is converted to ethylene and other products.