Use the chemical equation to answer the question. 2Ag(s) + H2S(g) → Ag2S(s) + H2(g) The molar mass of silver (Ag) is 108 g/mol. The molar mass of sulfur (S) is 32 g/mol. The reaction uses 0.04 mol of silver. Which steps show how to determine the mass of silver sulfide (Ag2S) produced in the reaction? (1 point) Responses 2(108 g/mol)+32 g/mol=248 g/mol ; (248 g/mol)(0.02 mol)=4.96 g 2(108 g/mol)+32 g/mol=248 g/mol ; (248 g/mol)(0.02 mol)=4.96 g 108 g/mol+2(32 g/mol)=172 g/mol ; (172 g/mol)(0.04 mol)=6.88 g 108 g/mol+2(32 g/mol)=172 g/mol ; (172 g/mol)(0.04 mol)=6.88 g 2(108 g/mol)+32 g/mol=248 g/mol ; (248 g/mol)(0.04 mol)=9.92 g 2(108 g/mol)+32 g/mol=248 g/mol ; (248 g/mol)(0.04 mol)=9.92 g 108 g/mol+2(32 g/mol)=172 g/mol ; (172 g/mol)(0.02 mol)=3.44 g

Answer 1

The correct steps to determine the mass of silver sulfide (Ag2S) produced in the reaction are:

1. Calculate the molar mass of silver sulfide (Ag2S):
2(108 g/mol) + 32 g/mol = 248 g/mol

2. Calculate the mass of silver sulfide produced using the number of moles of silver used in the reaction (0.04 mol):
248 g/mol x 0.04 mol = 9.92 g

Therefore, the correct answer is:
2(108 g/mol) + 32 g/mol = 248 g/mol ; (248 g/mol)(0.04 mol) = 9.92 g.