Calculate the number of moles of [tex]O_2[/tex] and [tex]H_2O[/tex] required to react completely with 1.55 mol [tex]Fe[/tex] in the following reaction:

\[ 2Fe(s) + O_2(g) + 2H_2O(l) \rightarrow 2Fe(OH)_2(s) \]

A) 0.775 mol [tex]O_2[/tex] and 1.55 mol [tex]H_2O[/tex]
B) 1.55 mol [tex]O_2[/tex] and 0.775 mol [tex]H_2O[/tex]
C) 3.10 mol [tex]O_2[/tex] and 1.55 mol [tex]H_2O[/tex]
D) 1.55 mol [tex]O_2[/tex] and 3.10 mol [tex]H_2O[/tex]

To react completely with 1.55 moles of Fe, 0.775 moles of O₂ and 1.55 moles of H2O are needed, therefore, option A is the correct answer. This is determined by the stoichiometry of the balanced chemical equation.

Explanation:

Based on the balanced reaction equation, 2Fe(s) + O₂(g) + 2H₂O(l) → 2Fe(OH)2(s), it can be seen that for every two moles of Fe reacted, one mole of O₂ and two moles of H₂O are required. Therefore, just by looking at the stoichiometry (the mole ratios) you can determine the required amount of O₂ and H₂O. So, for 1.55 moles of Fe, you will need 0.775 moles of O₂ and 1.55 moles of H₂O, which corresponds to option A. These calculations are a fundamental aspect of stoichiometry, which involves relationship between reactants and/or products in a chemical reaction.

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