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The problem in the image asks to calculate the mole fraction, partial pressure of each gas, and total pressure in a tank filled with sulfur tetrafluoride (SF₄) and chlorine pentafluoride (ClF₅) gases. The gases are assumed to behave as ideal gases. 

We are given the following dataThe problem in the image asks to calculate the mole fraction, partial pressure of each gas, and total pressure in a tank filled with sulfur tetrafluoride (SF₄) and chlorine pentafluoride (ClF₅) gases. The gases are assumed to behave as ideal gases. 

We are given the following 
- Volume of the tank: \( V = 5.00 \, \text{L} \)
- Temperature: \( T = 4.84^\circ \text{C} \) \( = 4.84 + 273.15 = 278.0 \, \text{K} \)
- Mass of sulfur tetrafluoride (SF₄): \( 4.59 \, \text{g} \)
- Mass of chlorine pentafluoride (ClF₅): \( 16.4 \, \text{g} \)
  
The steps to solve this are:
1. Convert the given masses to moles using the molar masses:
   - Molar mass of SF₄: \( 144.06 \, \text{g/mol} \)
   - Molar mass of ClF₅: \( 130.44 \, \text{g/mol} \)

2. Calculate the moles of each gas:
   \[
   n_{\text{SF₄}} = \frac{\text{mass of SF₄}}{\text{molar mass of SF₄}}
   \]
   \[
   n_{\text{ClF₅}} = \frac{\text{mass of ClF₅}}{\text{molar mass of ClF₅}}
   \]

3. Calculate the mole fractions of each gas:
   \[
   \text{mole fraction of SF₄} = \frac{n_{\text{SF₄}}}{n_{\text{SF₄}} + n_{\text{ClF₅}}}
   \]
   \[
   \text{mole fraction of ClF₅} = \frac{n_{\text{ClF₅}}}{n_{\text{SF₄}} + n_{\text{ClF₅}}}
   \]

4. Use the ideal gas law \( PV = nRT \) to calculate the total pressure in the tank, where:
   - \( R = 0.0821 \, \text{L·atm/(mol·K)} \)
   - \( V = 5.00 \, \text{L} \)
   - \( T = 278.0 \, \text{K} \)
   - \( n_{\text{total}} = n_{\text{SF₄}} + n_{\text{ClF₅}} \)
   
   The total pressure is:
   \[
   P_{\text{total}} = \frac{n_{\text{total}} \cdot R \cdot T}{V}
   \]

5. Finally, calculate the partial pressures using:
   \[
   P_{\text{SF₄}} = \text{mole fraction of SF₄} \times P_{\text{total}}
   \]
   \[
   P_{\text{ClF₅}} = \text{mole fraction of ClF₅} \times P_{\text{total}}
   \]

Let me calculate these values for you.Here are the results:

1. **Mole fraction of sulfur tetrafluoride (SF₄)**: \( 0.202 \)
2. **Mole fraction of chlorine pentafluoride (ClF₅)**: \( 0.798 \)
3. **Partial pressure of SF₄**: \( 0.145 \, \text{atm} \)
4. **Partial pressure of ClF₅**: \( 0.574 \, \text{atm} \)
5. **Total pressure in the tank**: \( 0.719 \, \text{atm} \)

Let me know if you'd like further clarification or details!

---

**Relative Questions:**
1. How does temperature affect the total pressure in the tank?
2. What would happen to the mole fraction if more sulfur tetrafluoride were added?
3. How would you calculate the pressure if one of the gases were removed?
4. Can the ideal gas law be applied to non-ideal gases in real-world conditions?
5. How would the total pressure change if the tank volume were reduced?

**Tip:** The sum of the partial pressures of all gases in a mixture equals the total pressure (Dalton's Law of Partial Pressures).

A 5.00 L tank at 4.84°C is filled with 4.59 g of sulfur tetrafluoride gas and 16.4 g of chlorine pentafluoride gas. You are asked to calculate the mole fraction and partial pressure of each gas, and the total pressure in the tank, assuming ideal gas behavior.

This problem involves calculating the mole fractions, partial pressures, and total pressure in a 5.00 L tank containing sulfur tetrafluoride and chlorine pentafluoride gases at 4.84°C. It uses the Ideal Gas Law and Dalton's Law of Partial Pressures.

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