Answer:
a.3.2204 g
b.Ag+ = 0 M
NO3? = 0.105 M
Ca2+ = 0.07M
Cl? =0.035 M
Explanation:
The balanced equation for this reaction is:
[tex]2AgNO_{3}+CaCl_{2} - - -> 2AgCl + Ca(NO_{3})_{2}[/tex]
mol of [tex]AgNO_{3}[/tex]
[tex]mol= Molarity*volume(L)=0.21M*0.107L=0.02247 mol[/tex]
mol of [tex]CaCl_{2}[/tex]
[tex]mol= Molarity*volume(L)=0.14M*0.107L=0.01498 mol[/tex]
from the reaction stoichiometry we know that we need 2 moles of [tex]AgNO_{3}[/tex] per every mol of [tex]CaCl_{2}[/tex].
so for 0.01498 mol of [tex]CaCl_{2}[/tex] we need 0.01498*2 =0.02996 mol of [tex]AgNO_{3}[/tex]. Since we have just 0.02247 mol, this will be the reaction limit reagent for this reaction.
we will use all the [tex]AgNO_{3}[/tex]
mol of [tex]CaCl_{2}[/tex] used =[tex]\frac{0.02247}{2}=0.011235mol[/tex]
For the calculation we always use the limit reagent.
a. How many grams of silver chloride.
From the reaction stoichiometry we know that 2 mol of [tex]AgNO_{3}[/tex] produce 2 mol of AgCl so, 0.02247 mol of [tex]AgNO_{3}[/tex] will the same number of mol of AgCl.
AgCl produced = 0.02247 mol
molecular weight of AgCl = 143.32 g/mol
mass of AgCl = [tex]mol*Mw = 0.02247 mol*143.32 \frac{g}{mol}=3.2204 g[/tex]
b.concentrations of each ion
volume of the remained solution 107 mL + 107mL= 214 mL or 0.214 L
[tex]Ag^{+}[/tex] since we used all the [tex]AgNO_{3}[/tex] to produced AgCl and this precipitated we don't have [tex]Ag^{+}[/tex] in the solution after the precipitation.
[tex]NO_{3}[/tex] this ion was used to produce [tex]Ca(NO_{3})_{2}[/tex] and this is soluble in water. The moles of[tex]NO_{3}[/tex] will be 2 times the number of moles of [tex]Ca(NO_{3})_{2}[/tex] produced o simply the same number of moles of [tex]2AgNO_{3}[/tex] before the reaction because this ion did not precipitate but it's concentration change because the volume changed.
concentration = [tex]\frac{0.02247 mol}{0.214L}=0.105 M[/tex]
following the same analysis for made for [tex]NO_{3}[/tex] we know that [tex]Ca^{2+}[/tex] did not precipitated so we have the same number of moles of [tex]CaCl_{2}[/tex] in the original sample but in a different volume
concentration = [tex]\frac{0.01498 mol}{0.214L}=0.07 M[/tex]
the Cl in the solution after the precipitation will be the Cl that did not reacted since all the Cl that reacted to produce AgCl precipitated.
Cl that reacted is the same number of AgCl produced this is=0.02247 mol
Cl before the reaction is 2 times the number of moles of [tex]CaCl_{2}[/tex] in the sample this is: 0.01498*2 =0.02996 mol
Cl remained = 0.02996 mol -0.02247 mol=0.00749 mol
concentration = [tex]\frac{0.00749 mol}{0.214L}=0.035 M[/tex]
