Are You Overreacting? Understanding Limiting and Excess Reactants in Chemical Reactions

In the captivating world of chemistry, understanding how different substances interact is vital. One of the key concepts to grasp is the idea of limiting and excess reactants. These terms describe how reactants in a chemical reaction are consumed, and determine the amount of products produced (aka the yield of the reaction). Figuring out the limiting and excess reactants in a reaction is key to solving a stoichiometry problem, so buckle up to do some learning today!

The Concept of Limiting Reactants

In any chemical reaction, you may have multiple reactants, but they do not always get used up at the same rate. The limiting reactant is the substance that gets fully consumed first, which limits the amount of product that can be formed. This also limits and determines the yield of the products, and the amount (moles) of limiting reactant is directly related to that of the products. Once the limiting reactant is finished/used up, the reaction stops, even if other reactants are still available.

For example, consider the reaction of hydrogen and oxygen to form water:

2 H2 + 1 O2 -> 2 H2O

Let's say we have 2 g of H2, and 50 g of O2. According to the balanced reaction, you need 2 H2:1 O2 to get 2 moles of H2O. Let's calculate the moles of H2 and O2 (mass/molar mass):

H2 moles= 2 g/2g/mol= 1 mole

O2= 50 g/32 g/mol= 1.56 mole

According to the ratio of the balanced reaction, you would need 2 moles of H2 for every one mole of O2. Clearly, in this case, you have more O2 moles than H2, which doesn't fit the ratio! In this case, H2 is the limiting reactant.

To identify the limiting reactant, follow these steps:

1. Write the balanced equation for the chemical reaction.

2. Determine the number of moles of each reactant available.

3. Use stoichiometry from the balanced equation to find out how much product can be formed from each reactant.

4. The reactant that produces the fewest moles of product is the limiting reactant.

   
   

By mastering these calculations, anyone can predict reaction outcomes more accurately, making those high school Chemistry exams a breeze!

This is you, the next time you need to take a stoichiometry exam :)

Finding the Theoretical Yield of the Product using the Limiting Reactant

Now, let's say you wanted to measure how much water you produced using the previous reaction. How would you find it? Let's take a look at it again:

2 H2 + 1 O2 -> 2 H2O

2 g of H2= 1 mole of H2

50 g of O2= 1.56 mole of O2

Since H2 is the limiting reactant, when it gets used up, the reaction stops. Therefore, it determines the amount of product, and we can use the moles of the limiting reactant to find the moles of product, and then convert that into the mass of product, finally finding our yield!

Still confused? Okay, let's use this example. Suppose your baking a bunch of sandwiches for a bake sale, and you have 4 dozen loaves of bread, but only a dozen slices of cheese. Can you make a decent sandwich without cheese? Not really! So the number of sandwiches you make is *LIMITED* by the slices of cheese you have at your disposal.

Now, let us say that you now want to calculate your yield of water using the example above. First, you will want to determine the moles of water produced: since the coefficients of both H2 (our limiting reactant) and H2O (our product) are both 2, the ratios are both 1:1, and the moles of our limiting reactant is EQUAL to our moles of product:

1 mole of H2:1 mole of H2O because 2H2:2H2O!

Next, you will need to convert the moles of H2O to the mass of H2O. Mass= moles*

molar mass, so:

1 mol H2O*18 g/mol= 18 g of H2O

There you have it! We have figured out our theoretical yield!

A high angle view showing bubbling chemical reaction in a glass container.

Stay tuned for comparing theoretical yield and actual yield, as well as more blog posts on both beginner Chemistry like these, and more advanced Chemistry later! If you like articles like these, consider booking a session at Meenakshi's Tutoring. See you soon!

BYEEE