Chlorine \(\left(\mathrm{Cl}_{2}\right)\) and fluorine \(\left(\mathrm{F}_{2}\right)\) react to form \(\mathrm{ClF}_{3}\). A reaction vessel contains \(2.50\) moles of \(\mathrm{Cl}_{2}\) and \(6.15\) moles of \(\mathrm{F}_{2}\). (a) Write a balanced chemical equation for this reaction. (b) Which reactant is limiting?

One of the fundamental aspects of chemistry is to understand how to balance chemical equations. It involves finding the correct proportions of reactants and products to ensure that the law of conservation of mass is satisfied. This law states that matter is neither created nor destroyed in a chemical reaction.

Luckily, the process of balancing equations can be straightforward with practice. To demonstrate, consider the unbalanced equation from our example:
Cl_2 + F_2 → ClF_3
. Here, we have to make sure the number of chlorine (Cl) and fluorine (F) atoms on each side of the arrow is equal.

To achieve this, look at the number of fluorine atoms in the product, which is three in ClF3. Since the reactant fluorine comes as F2, you need to use 3/2 or 1.5 F2 molecules to have three fluorine atoms available for every ClF3. For chlorine, one Cl2 molecule will provide two Cl atoms, exactly what you need for the two ClF3 molecules produced. Thus, balancing gives you the equation:Cl_2 + \(\frac{3}{2}\)F_2 → 2ClF_3.

This correctly balanced equation ensures that the atoms are conserved across the reaction. Through balancing, not only is the law of conservation of mass upheld, but you also gain essential insights into the reaction ratios which is critical for the next steps involving stoichiometry.

The limiting reactant in a chemical reaction is the substance that is totally consumed first, stopping the reaction from continuing because there is no more of this reactant left to react. Identifying the limiting reactant is crucial, as it determines the amount of product that can be formed.

In the given example where chlorine (Cl2) and fluorine (F2) react to form ClF3, we initially have 2.50 moles of Cl2 and 6.15 moles of F2. Since the balanced chemical equation requires 1.5 moles of F2 for every mole of Cl2, we can quickly run out of Cl2 before using up all the F2 we started with.

To find the limiting reactant, you compare the mole ratio of the reactants with the stoichiometric ratio from the balanced equation. If the actual mole ratio of reactants is less than the stoichiometric ratio, then the first reactant is limiting, and vice versa. This comparison shows that in our scenario, Cl2 is the limiting reactant as its available amount can not fully react with all of the F2 present, leaving some F2 unreacted when all the Cl2 is gone.

Stoichiometry is a section of chemistry that deals with the quantification of reactants and products in a chemical reaction. It allows chemists to predict the amounts of substances consumed and produced in a given reaction.

The key to mastering stoichiometry lies in the balanced chemical equation, which tells us the ratio of moles of reactants needed to produce moles of product. In the reaction between Cl2 and F2 to produce ClF3, stoichiometry helps us understand that one mole of Cl2 will produce two moles of ClF3 as long as there is enough F2, according to the balanced equation Cl2 + \(\frac{3}{2}\)F2 → 2ClF3.

In practical terms, if a question asks you to find the amount of ClF3 produced from 2.50 moles of Cl2, you would use the stoichiometric coefficients from the balanced equation to find that 5.00 moles of ClF3 would be the expected yield—if Cl2 were not the limiting reactant. However, since Cl2 is limiting in this example, the actual yield would be less. Understanding stoichiometry is essential not just for solving textbook problems, but also for practical applications such as calculating the amount of reactants needed for a reaction to take place in the lab or industry settings.