If a 250.0 mL saturated solution of ionic compound MX₂ (Molar mass = 89.86 g/mol) contains 0.131 mg MX₂, which is the Ksp of MX₂? O 1.2 x 10-8 O 6.8 x 10-11 O 3.4 x 10-11 O 7.9 x 10-16

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### Solubility Product Constant (Ksp) Calculation

**Problem Statement:**

If a 250.0 mL saturated solution of ionic compound MX₂ (Molar mass = 89.86 g/mol) contains 0.131 mg MX₂, what is the Ksp of MX₂?

**Options:**

- \( 1.2 \times 10^{-8} \)
- \( 6.8 \times 10^{-11} \)
- \( 3.4 \times 10^{-11} \)
- \( 7.9 \times 10^{-16} \)

---

**Explanation:**

To solve this problem, we first need to calculate the molarity of the MX₂ solution and then use it to find the solubility product constant (Ksp).

1. **Convert mg to g:**
   \[
   0.131\ \text{mg} = 0.131 \times 10^{-3}\ \text{g} = 0.000131\ \text{g}
   \]

2. **Calculate the moles of MX₂:**
   \[
   \text{Moles of MX₂} = \frac{\text{mass}}{\text{molar mass}} = \frac{0.000131\ \text{g}}{89.86\ \text{g/mol}} = 1.46 \times 10^{-6}\ \text{mol}
   \]

3. **Calculate the molarity of the solution:**
   \[
   \text{Molarity (M)} = \frac{\text{moles of solute}}{\text{volume of solution in liters}} = \frac{1.46 \times 10^{-6}\ \text{mol}}{0.250\ \text{L}} = 5.84 \times 10^{-6}\ \text{M}
   \]

4. **Dissociation of MX₂:**
   \[
   \text{MX₂} \rightleftharpoons \text{M}^{2+} + 2\text{X}^{-}
   \]
   - Let \( s \) be the molarity of MX₂ (5.84 \times 10^{-6} M).
   - Then, [M²⁺] = \( s \) and [X⁻] = 2\(
Transcribed Image Text:### Solubility Product Constant (Ksp) Calculation **Problem Statement:** If a 250.0 mL saturated solution of ionic compound MX₂ (Molar mass = 89.86 g/mol) contains 0.131 mg MX₂, what is the Ksp of MX₂? **Options:** - \( 1.2 \times 10^{-8} \) - \( 6.8 \times 10^{-11} \) - \( 3.4 \times 10^{-11} \) - \( 7.9 \times 10^{-16} \) --- **Explanation:** To solve this problem, we first need to calculate the molarity of the MX₂ solution and then use it to find the solubility product constant (Ksp). 1. **Convert mg to g:** \[ 0.131\ \text{mg} = 0.131 \times 10^{-3}\ \text{g} = 0.000131\ \text{g} \] 2. **Calculate the moles of MX₂:** \[ \text{Moles of MX₂} = \frac{\text{mass}}{\text{molar mass}} = \frac{0.000131\ \text{g}}{89.86\ \text{g/mol}} = 1.46 \times 10^{-6}\ \text{mol} \] 3. **Calculate the molarity of the solution:** \[ \text{Molarity (M)} = \frac{\text{moles of solute}}{\text{volume of solution in liters}} = \frac{1.46 \times 10^{-6}\ \text{mol}}{0.250\ \text{L}} = 5.84 \times 10^{-6}\ \text{M} \] 4. **Dissociation of MX₂:** \[ \text{MX₂} \rightleftharpoons \text{M}^{2+} + 2\text{X}^{-} \] - Let \( s \) be the molarity of MX₂ (5.84 \times 10^{-6} M). - Then, [M²⁺] = \( s \) and [X⁻] = 2\(
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