Question

Consider the following chemical reaction:

2H₂O(l) → 2H₂(g) + O₂(g)
What mass of H₂O is required to form 1.6 L of O₂ at a temperature of 310 K and a pressure of 0.916 atm?
A) 1.43 g
B) 2.87 g
C) 4.30 g
D) 5.74 g

Answer 1

To determine the mass of H₂O needed to form 1.6 L of O₂ at a temperature of 310 K and a pressure of 0.916 atm, we need to use the ideal gas law and stoichiometry. First, let's convert the given volume of O₂ into moles using the ideal gas law. According to the balanced chemical equation, 1 mole of O₂ reacts with 2 moles of H₂O, so we can calculate the moles of H₂O and then convert it to grams using the molar mass.

Step-by-step explanation:

To determine the mass of H₂O needed to form 1.6 L of O₂ at a temperature of 310 K and a pressure of 0.916 atm, we need to use the ideal gas law and stoichiometry.

First, let's convert the given volume of O₂ into moles using the ideal gas law:

n = PV / RT = (0.916 atm) * (1.6 L) / (0.0821 L*atm/mol*K * 310 K) ≈ 0.0803 moles of O₂

According to the balanced chemical equation, 1 mole of O₂ reacts with 2 moles of H₂O. Therefore, 0.0803 moles of O₂ would react with 0.0803 * 2 = 0.1606 moles of H₂O.

Finally, we can convert the moles of H₂O into grams using the molar mass:

Mass = moles * molar mass = 0.1606 moles * 18.02 g/mol ≈ 2.89 g