The Harber process is used making ammonia from nitrogen and hydrogen according to the equation shown below. The yield of the reaction, however, is not \(100 \%\). a. Suppose you end up with \(6.2\) moles of ammonia, but the reaction stoichiometry predicts that you should have \(170.0\) grams of ammonia. What is the percent yield for this reaction? b. If you started with \(6.2\) grams of nitrogen and you produce \(6.2\) grams of ammonia what would be the percent yield?

Stoichiometry is a branch of chemistry that deals with the quantitative relationships between the reactants and products in a chemical reaction. In essence, it's about balancing equations and using these balances to predict the amount of substance required or produced in a reaction.

For instance, when examining a reaction such as the Haber Process, stoichiometry informs us of the proportional relationship between nitrogen (N₂), hydrogen (H₂), and ammonia (NH₃). This ratio is crucial when trying to determine the theoretical and experimental yields of a reaction. Knowing how to perform stoichiometric calculations enables students to grasp how much of each reactant is required to produce a certain amount of product.

The molar mass of a substance is the mass of one mole of that substance. It is a physical property that is of particular interest in the study of stoichiometry, as it provides a link between the mass of a substance and the number of particles or moles present. Molar mass is usually expressed in grams per mole (g/mol).

In the provided exercise, the molar mass of ammonia (NH₃) is 17 g/mol. This means that one mole of ammonia weighs 17 grams. For nitrogen (N₂), the molar mass is 28 g/mol. These molar masses are essential in converting between moles and grams, which is a common step in stoichiometry when working out percentages or predicting yields.

The theoretical yield is the maximum amount of product that can be produced in a chemical reaction based on the amount of limiting reactant, assuming complete conversion and no side reactions. This yield is called 'theoretical' because it's what we would expect if everything happens perfectly, which is rarely the case in practical settings.

The steps in calculating the theoretical yield involve using the balanced chemical equation to figure out the mol-to-mol relationship between reactants and products. In the example of the Haber process, the theoretical yield of ammonia is determined by using the stoichiometric ratios between nitrogen, hydrogen, and ammonia.

The experimental yield is the actual amount of product obtained from a chemical reaction, which is often less than the theoretical yield due to various practical factors such as incomplete reactions, side reactions, or loss of product during the process.

Calculating the percentage yield requires both the experimental yield and the theoretical yield. It's a measure of the efficiency of the reaction, telling us how much of the theoretical yield was actually achieved. The percent yield formula, shown in the exercise solution, illustrates how to compare the experimental yield (i.e., the actual, observed result) to the theoretical prediction.

The Haber process is an industrial method for the synthesis of ammonia from nitrogen and hydrogen gas. This reaction plays a significant role in modern agriculture, as ammonia is a key precursor to fertilizers. The balanced chemical equation for the Haber process is N₂ + 3H₂ ↔ 2NH₃. Despite its apparent simplicity, the Haber process is a dynamic equilibrium, with several factors affecting the yield of ammonia.

When calculating yields for reactions like the Haber process, it's vital to understand that the actual (experimental) yield will often be lower than the one predicted stoichiometrically (theoretical yield). The percent yield calculation allows us to measure the efficiency of the reaction in converting reactants to desired products under real-world conditions.