Given the following data table for the reaction: 2 H2 + O2 → 2 H2O, if you start with 4.0 mol of H2 and 2.0 mol of O2, how many moles of H2O will be produced?

• A. 2.0 mol
• B. 4.0 mol
• C. 8.0 mol
• D. 6.0 mol

Answer: (B)

Verified step by step guidance

1. Identify the balanced chemical equation: \$2 \mathrm{H}_2 + \mathrm{O}_2 \rightarrow 2 \mathrm{H}_2\mathrm{O}$.
2. Determine the mole ratio between the reactants and the product from the balanced equation. For every 2 moles of \(\mathrm{H}_2\) and 1 mole of \(\mathrm{O}_2\), 2 moles of \(\mathrm{H}_2\mathrm{O}\) are produced.
3. Compare the given moles of reactants to find the limiting reactant. Calculate how many moles of \(\mathrm{H}_2\mathrm{O}\) can be formed from each reactant separately using the mole ratios:
4. \[\text{From } \mathrm{H}_2: \quad 4.0 \text{ mol } \mathrm{H}_2 \times \frac{2 \text{ mol } \mathrm{H}_2\mathrm{O}}{2 \text{ mol } \mathrm{H}_2}\]
5. \[\text{From } \mathrm{O}_2: \quad 2.0 \text{ mol } \mathrm{O}_2 \times \frac{2 \text{ mol } \mathrm{H}_2\mathrm{O}}{1 \text{ mol } \mathrm{O}_2}\]
6. The limiting reactant is the one that produces the smaller amount of \(\mathrm{H}_2\mathrm{O}\). The number of moles of \(\mathrm{H}_2\mathrm{O}\) produced equals the amount calculated from the limiting reactant.