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Harlamova29_29 [7]
3 years ago
13

How many grams of carbon monoxide are needed to react with an excess of iron (III) oxide to produce 198.5 grams of iron? Fe2O3(s

) + 3CO(g) -----à 3CO2(g) + 2Fe(s)
Chemistry
1 answer:
erastova [34]3 years ago
7 0

Answer : The grams of carbon monoxide needed are 148.89 g

Solution : Given,

Mass of iron, Fe = 198.5 g

Molar mass of iron, Fe = 56 g/mole

Molar mass of carbon monoxide, CO = 28 g/mole

First we have to calculate the moles of iron, Fe.

\text{ Moles of Fe}=\frac{\text{ Mass of Fe}}{\text{ Molar mass of Fe}}=\frac{198.5g}{56g/mole}=3.545moles

The balanced chemical reaction is,

Fe_2O_3(s)+3CO(g)\rightarrow 3CO_2(g)+2Fe(s)

From the balanced reaction, we conclude that

2 moles of iron produces from the 3 moles of carbon monoxide

3.545 moles of iron produces from the \frac{3}{2}\times 3.545=5.3175 moles of carbon monoxide

Now we have to calculate the mass of carbon monoxide, CO.

\text{ Mass of CO}=\text{ Moles of CO}\times \text{ Molar mass of CO}

\text{ Mass of CO}=(5.3175moles)\times (28g/mole)=148.89g

Therefore, the grams of carbon monoxide needed are 148.89 g

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How many grams of Sg is required to produce 83.10 g SF6? S: +24F-->8SF
ozzi

Answer : The mass of S_8 required is 18.238 grams.

Explanation : Given,

Mass of SF_6 = 83.10 g

Molar mass of SF_6 = 146 g/mole

Molar mass of S_8 = 256.52 g/mole

The balanced chemical reaction is,

S_8+24F_2\rightarrow 8SF_6

First we have to determine the moles of SF_6.

\text{Moles of }SF_6=\frac{\text{Mass of }SF_6}{\text{Molar mass of }SF_6}=\frac{83.10g}{146g/mole}=0.569moles

Now we have to determine the moles of S_8.

From the balanced chemical reaction we conclude that,

As, 8 moles of SF_6 produced from 1 mole of S_8

So, 0.569 moles of SF_6 produced from \frac{0.569}{8}=0.0711 mole of S_8

Now we have to determine the mass of S_8.

\text{Mass of }S_8=\text{Moles of }S_8\times \text{Molar mass of }S_8

\text{Mass of }S_8=(0.0711mole)\times (256.52g/mole)=18.238g

Therefore, the mass of S_8 required is 18.238 grams.

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3 years ago
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