When 3.00 moles of hydrogen molecules and 1.50 moles of oxygen molecules react they form 3.00 moles of water according to the reaction below. Does the mass of oxygen required equal the mass of the water formed? 2H_(2)+O_(2)arrow 2H_(2)O no yes

To determine whether the mass of oxygen required equals the mass of the water formed, we need to compare the masses involved in the reaction.<br /><br />First, let's write down the balanced chemical equation:<br />\[ 2H_{2} + O_{2} \rightarrow 2H_{2}O \]<br /><br />Given:<br />- 3.00 moles of \( H_2 \)<br />- 1.50 moles of \( O_2 \)<br />- 3.00 moles of \( H_2O \)<br /><br />From the balanced equation, we see that 2 moles of \( H_2 \) react with 1 mole of \( O_2 \) to produce 2 moles of \( H_2O \).<br /><br />Now, let's calculate the molar masses:<br />- Molar mass of \( H_2 \): \( 2 \times 1.01 \text{ g/mol} = 2.02 \text{ g/mol} \)<br />- Molar mass of \( O_2 \): \( 2 \times 16.00 \text{ g/mol} = 32.00 \text{ g/mol} \)<br />- Molar mass of \( H_2O \): \( 2 \times 1.01 + 16.00 \text{ g/mol} = 18.02 \text{ g/mol} \)<br /><br />Next, calculate the mass of oxygen required:<br />\[ \text{Mass of } O_2 = 1.50 \text{ moles} \times 32.00 \text{ g/mol} = 48.00 \text{ g} \]<br /><br />Calculate the mass of water formed:<br />\[ \text{Mass of } H_2O = 3.00 \text{ moles} \times 18.02 \text{ g/mol} = 54.06 \text{ g} \]<br /><br />Comparing the masses:<br />- Mass of \( O_2 \) required: 48.00 g<br />- Mass of \( H_2O \) formed: 54.06 g<br /><br />Since the mass of oxygen required (48.00 g) does not equal the mass of the water formed (54.06 g), the answer is:<br /><br />**No**