The `%` of `Fe^(+2)` in `Fe_(0.93)O_(1.00)` is .

Options

A. `15%`
B. `85%`
C. `93%`
D. `7%`

Explanation

To find the percentage of \( \text{Fe}^{2+} \) in \( \text{Fe}_{0.93}\text{O}_{1.00} \), we can follow these steps:

### Step 1: Define Variables
Let \( x \) be the amount of \( \text{Fe}^{2+} \) in the compound. The remaining iron will then be in the \( \text{Fe}^{3+} \) state.

### Step 2: Write the Total Amount of Iron
The total amount of iron in the compound is given as \( 0.93 \) moles. Therefore, the amount of \( \text{Fe}^{3+} \) can be expressed as:
\[ \text{Amount of } \text{Fe}^{3+} = 0.93 - x \]

### Step 3: Set Up the Charge Balance Equation
The charge balance must equal zero. The total positive charge from the iron must equal the total negative charge from the oxygen. The equation can be set up as follows:
\[ 2x + 3(0.93 - x) - 2 = 0 \]
Here, \( 2x \) is the contribution from \( \text{Fe}^{2+} \), \( 3(0.93 - x) \) is the contribution from \( \text{Fe}^{3+} \), and \( -2 \) is from the two oxygen atoms (each contributing a charge of -2).

### Step 4: Simplify the Equation
Now, simplify the equation:
\[ 2x + 2.79 - 3x - 2 = 0 \]
Combine like terms:
\[ - x + 0.79 = 0 \]

### Step 5: Solve for \( x \)
Rearranging gives:
\[ x = 0.79 \]

### Step 6: Calculate the Percentage of \( \text{Fe}^{2+} \)
Now, we can find the percentage of \( \text{Fe}^{2+} \) in the compound:
\[ \text{Percentage of } \text{Fe}^{2+} = \left( \frac{x}{0.93} \right) \times 100 \]
Substituting \( x = 0.79 \):
\[ \text{Percentage of } \text{Fe}^{2+} = \left( \frac{0.79}{0.93} \right) \times 100 \approx 84.95\% \]
Rounding gives approximately \( 85\% \).

### Conclusion
Thus, the percentage of \( \text{Fe}^{2+} \) in \( \text{Fe}_{0.93}\text{O}_{1.00} \) is approximately \( 85\% \).

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