which reactant is limiting when 5 moles of N2 gas react with 10 moles of H2 gas to make ammonia?
The correct answer and explanation is:
To determine the limiting reactant in the reaction between nitrogen gas (N₂) and hydrogen gas (H₂) to form ammonia (NH₃), we first look at the balanced chemical equation: N2+3H2→2NH3\text{N}_2 + 3\text{H}_2 \rightarrow 2\text{NH}_3
Step 1: Understand the mole ratio
From the balanced equation:
- 1 mole of N₂ reacts with 3 moles of H₂.
Step 2: Compare available moles
We are given:
- 5 moles of N₂
- 10 moles of H₂
Now, let’s see how many moles of H₂ are needed to fully react with the 5 moles of N₂: 5 mol N2×3 mol H21 mol N2=15 mol H25 \text{ mol N}_2 \times \frac{3 \text{ mol H}_2}{1 \text{ mol N}_2} = 15 \text{ mol H}_2
So, we need 15 moles of H₂ to react with 5 moles of N₂, but we only have 10 moles of H₂.
Step 3: Identify the limiting reactant
Since we don’t have enough H₂, it will run out first. Therefore, H₂ is the limiting reactant.
Explanation (Approx. 300 words):
In chemical reactions, the limiting reactant is the substance that is completely used up first, limiting the amount of product formed. To determine the limiting reactant between nitrogen gas (N₂) and hydrogen gas (H₂) in the synthesis of ammonia (NH₃), we must use the balanced chemical equation: N2+3H2→2NH3\text{N}_2 + 3\text{H}_2 \rightarrow 2\text{NH}_3
This tells us that for every 1 mole of nitrogen, 3 moles of hydrogen are required. Given 5 moles of N₂ and 10 moles of H₂, we calculate how much hydrogen would be needed to react with all the nitrogen: 5 mol N2×3=15 mol H25 \text{ mol N}_2 \times 3 = 15 \text{ mol H}_2
However, we only have 10 moles of hydrogen, which is insufficient. Because we don’t have the full 15 moles needed, the reaction cannot consume all 5 moles of nitrogen. Hydrogen will be used up first, stopping the reaction from proceeding further. That makes hydrogen the limiting reactant.
Identifying the limiting reactant is important in stoichiometry because it determines the maximum amount of product that can be formed. Even if there’s excess of the other reactant, the reaction stops once the limiting one is gone. In this case, since H₂ is the limiting reactant, the amount of ammonia produced depends on the 10 moles of hydrogen available.