To solve this, we'll need to calculate the amount of silver sulfide (Ag₂S) produced based on the given quantities.

Step 1: Calculate the moles of silver nitrate (AgNO₃)

We are given:

• Volume of AgNO₃ solution = 0.0150 L
• Molarity of AgNO₃ solution = 2.50 M (mol/L)

The number of moles of AgNO₃ is given by the formula: $\text{moles of AgNO₃} = \text{molarity} \times \text{volume (in L)}$ $\text{moles of AgNO₃} = 2.50 \, \text{mol/L} \times 0.0150 \, \text{L} = 0.0375 \, \text{mol}$

Step 2: Use the stoichiometry of the reaction to find moles of Ag₂S

From the balanced chemical equation: $2 \, \text{AgNO₃} (aq) + \text{Na₂S} (aq) \rightarrow \text{Ag₂S} (s) + 2 \, \text{NaNO₃} (aq)$ From the equation, 2 moles of AgNO₃ react to produce 1 mole of Ag₂S. Thus, the moles of Ag₂S produced are: $\text{moles of Ag₂S} = \frac{\text{moles of AgNO₃}}{2} = \frac{0.0375 \, \text{mol}}{2} = 0.01875 \, \text{mol}$

Step 3: Calculate the mass of silver sulfide (Ag₂S)

Next, we need to find the molar mass of Ag₂S. Using the atomic masses from the periodic table:

• Silver (Ag) = 107.87 g/mol
• Sulfur (S) = 32.07 g/mol

The molar mass of Ag₂S is: $\text{molar mass of Ag₂S} = 2 \times 107.87 \, \text{g/mol} + 32.07 \, \text{g/mol} = 247.81 \, \text{g/mol}$

Now, we can calculate the mass of Ag₂S produced: $\text{mass of Ag₂S} = \text{moles of Ag₂S} \times \text{molar mass of Ag₂S}$ $\text{mass of Ag₂S} = 0.01875 \, \text{mol} \times 247.81 \, \text{g/mol} = 4.65 \, \text{g}$

Final Answer:

The amount of silver sulfide produced is 4.65 g, which corresponds to option B.

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Would you like further details or have any questions?

Related Questions:

1. How would the amount of silver sulfide change if a different volume of AgNO₃ solution were used?
2. What would happen if the sodium sulfide (Na₂S) were in excess?
3. How could you determine the limiting reagent if you were given excess AgNO₃?
4. How do you calculate the molarity of a solution from the mass and volume of solute?
5. What are some real-world applications of silver sulfide in chemistry?

Tip: Always check if a reactant is in excess, as this will help simplify stoichiometric calculations.