`2.1g` of `Fe ` combines with `S` evoling `3.77KJ` . The heat oof formation of `FeS` in `KJ//mol` is

Options

A. `-3.77`
B. `-1.79`
C. `-100.5`
D. none of these

Explanation

To find the heat of formation of FeS in kJ/mol, we can follow these steps:

### Step 1: Calculate the moles of Iron (Fe)
The molar mass of Iron (Fe) is approximately 56 g/mol. Given that we have 2.1 g of Fe, we can calculate the number of moles of Fe using the formula:

\[ \text{Moles of Fe} = \frac{\text{mass of Fe}}{\text{molar mass of Fe}} = \frac{2.1 \, \text{g}}{56 \, \text{g/mol}} \]

Calculating this gives:

\[ \text{Moles of Fe} = \frac{2.1}{56} \approx 0.0375 \, \text{mol} \]

### Step 2: Determine the heat evolved per mole of FeS
The problem states that 3.77 kJ of energy is evolved when 2.1 g of Fe reacts with sulfur. Since the reaction is:

\[ \text{Fe} + \text{S} \rightarrow \text{FeS} \]

This indicates that 1 mole of Fe reacts to form 1 mole of FeS. Therefore, the moles of FeS formed will also be 0.0375 mol.

### Step 3: Calculate the heat of formation per mole of FeS
To find the heat of formation per mole of FeS, we can use the energy evolved and the moles of FeS formed:

\[ \text{Heat of formation of FeS} = \frac{\text{Total heat evolved}}{\text{Moles of FeS}} = \frac{3.77 \, \text{kJ}}{0.0375 \, \text{mol}} \]

Calculating this gives:

\[ \text{Heat of formation of FeS} = \frac{3.77}{0.0375} \approx 100.53 \, \text{kJ/mol} \]

### Step 4: Account for the sign of the heat of formation
Since the reaction releases heat, the heat of formation is considered negative:

\[ \text{Heat of formation of FeS} = -100.53 \, \text{kJ/mol} \]

### Final Answer
The heat of formation of FeS is approximately **-100.5 kJ/mol**.

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