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The following balanced equation shows the formation of water.

[tex]\[ H_2 + O_2 \rightarrow 2 H_2O \][/tex]

How many moles of oxygen [tex]\(\left( O_2 \right)\)[/tex] are required to react completely with [tex]\(1.67 \, \text{mol} \, H_2\)[/tex]?

A. [tex]\(0.835 \, \text{mol} \, O_2\)[/tex]
B. [tex]\(1.67 \, \text{mol} \, O_2\)[/tex]
C. [tex]\(3.34 \, \text{mol} \, O_2\)[/tex]
D. [tex]\(6.68 \, \text{mol} \, O_2\)[/tex]

Sagot :

GinnyAnswer
To determine how many moles of oxygen ([tex]\(O_2\)[/tex]) are required to react completely with [tex]\(1.67\)[/tex] moles of hydrogen ([tex]\(H_2\)[/tex]), let's first examine the given balanced chemical equation:

[tex]\[ H_2 + O_2 \rightarrow 2 H_2O \][/tex]

The balanced equation indicates the stoichiometric relationship between the reactants and products. According to this equation:
- 1 mole of [tex]\(O_2\)[/tex] reacts with 2 moles of [tex]\(H_2\)[/tex].

Given this stoichiometric relationship:

[tex]\[ 1 \text{ mole of } O_2 \text{ reacts with } 2 \text{ moles of } H_2 \][/tex]

we can set up the following proportion to find out how many moles of [tex]\(O_2\)[/tex] are needed for [tex]\(1.67\)[/tex] moles of [tex]\(H_2\)[/tex]:

[tex]\[ \frac{\text{moles of } O_2}{1} = \frac{\text{moles of } H_2}{2} \][/tex]

We know the moles of [tex]\(H_2\)[/tex] is [tex]\(1.67\)[/tex]:

[tex]\[ \frac{\text{moles of } O_2}{1} = \frac{1.67}{2} \][/tex]

Solving for the moles of [tex]\(O_2\)[/tex]:

[tex]\[ \text{moles of } O_2 = \frac{1.67}{2} = 0.835 \][/tex]

Therefore, the number of moles of [tex]\(O_2\)[/tex] required to react completely with [tex]\(1.67\)[/tex] moles of [tex]\(H_2\)[/tex] is:

[tex]\[ 0.835 \text{ moles of } O_2 \][/tex]

So, the correct choice is:

[tex]\[ \boxed{0.835 \text{ mol } O_2} \][/tex]
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