Silver ions can be precipitated from aqueous solutions by the addition of aqueous chloride: $\mathrm { Ag } ^ { + }$ (aq)+ $\mathrm { Cl } ^ { - }$ (aq)? AgCl(s) Silver chloride is virtually insoluble in water so that the reaction appears to go to completion.How many grams of solid NaCl must be added to 25.0 mL of 0.376 mol L^-1 $\mathrm { AgNO } _ { 3 }$ solution to completely precipitate the silver?

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The balanced chemical equation is:$$\mathrm { Ag } ^ { + }\) (aq)+ $\mathrm { Cl } ^ { - }$ (aq)? AgCl(s)From the stoichiometry of the balanced chemical equation, we can see that 1 mol of $\mathrm { Ag } ^ { + }$ reacts with 1 mol of $\mathrm { Cl } ^ { - }$. Therefore, the number of moles of $\mathrm { Cl } ^ { - }$ required to precipitate all of the $\mathrm { Ag } ^ { + }$ ions in 25.0 mL of 0.376 mol L^-1 $\mathrm { AgNO } _ { 3 }$ solution is:$$25.0 \text{ mL} \times \frac{1 \text{ L}}{1000 \text{ mL}} \times 0.376 \text{ mol/L} = 9.40 \times 10^{-3} \text{ mol}$$The molar mass of NaCl is 58.44 g/mol. Therefore, the mass of NaCl required to provide this number of moles of $\mathrm { Cl } ^ { - }$ ions is:$$9.40 \times 10^{-3} \text{ mol} \times 58.44 \text{ g/mol} = 0.549 \text{ g}$$Therefore, the correct answer is B) 0.549 g.

Silver Ions Can Be Precipitated from Aqueous Solutions by the Addition