Consider the reaction between calcium oxide and carbon dioxide: $$ \mathrm{CaO}(s)+\mathrm{CO}_{2}(g) \longrightarrow \mathrm{CaCO}_{3}(s) $$ A chemist allows \(14.4 \mathrm{~g}\) of \(\mathrm{CaO}\) and \(13.8 \mathrm{~g}\) of \(\mathrm{CO}_{2}\) to react. When the reaction is finished, the chemist collects \(19.4 \mathrm{~g}\) of \(\mathrm{CaCO}_{3}\). Determine the limiting reactant, theoretical yield, and percent yield for the reaction.

Understanding the concept of the limiting reactant is crucial when performing chemical reactions. It is the substance in a chemical reaction that runs out first, thereby preventing more product from being formed. In the given exercise, the chemist had both calcium oxide (CaO) and carbon dioxide (CO2) to start the reaction.

To determine which one is the limiting reactant, a calculation of the number of moles of each reactant based on their respective masses and molar masses must be done. Then, using the stoichiometric coefficients from the balanced chemical equation, you compare the mole ratios to see which reactant would be completely consumed first when reacting with the available amount of the other substance. This reactant is identified as the limiting reactant because it limits the extent of the reaction and thus the amount of product formed.

The theoretical yield is the maximum amount of product that can be generated from a chemical reaction if all the limiting reactant is used up entirely with no side reactions or losses. It is based purely on the stoichiometry of the balanced equation and the amount of the limiting reactant.

In our exercise, after identifying the limiting reactant, it is imperative to calculate the moles of product that can be formed from it. The mole ratio from the balanced equation is used to convert the moles of the limiting reactant to the moles of the desired product. Then, this number is multiplied by the molar mass of the product to obtain the theoretical yield in grams. It is important to note that the theoretical yield is an ideal maximum value that assumes perfect conversion and no experimental errors.

The percent yield is a measure of the efficiency of a chemical reaction and is calculated by comparing the actual yield (the amount of product actually obtained from the experiment) to the theoretical yield.

To calculate the percent yield, the formula is: \[\text{percent yield} = \left(\frac{\text{actual yield}}{\text{theoretical yield}}\right) \times 100%\]. The actual yield will often be less than the theoretical yield due to experimental errors, incomplete reactions, impurities, or other side reactions. In the exercise, the student would use the mass of the CaCO3 collected (19.4g) as the actual yield and the previously calculated theoretical yield to determine the percent yield of the reaction. This value is important for chemists to understand the effectiveness and practicality of the chemical reaction.

A chemical reaction involves the transformation of one or more substances, called reactants, into one or more new substances, known as products.

In the provided exercise, calcium oxide reacts with carbon dioxide to produce calcium carbonate. This process is represented by a balanced chemical equation showing that one mole of CaO reacts with one mole of CO2 to yield one mole of CaCO3. Understanding the stoichiometry of a chemical reaction is crucial for predicting the amounts of products formed and for calculating the limiting reactant, theoretical yield, and percent yield. It also enables us to grasp the relationship between reactants and products, which is fundamental for all chemical processes, whether in a laboratory setting or industrial production.