Question #2070f

You can't actually convert a #"0.1 M"# solution into a #"0.1 N"# solution, because normality depends on what solution you're dealing with.

Normality is usually used in acid-base reactions and is a means of expressing the number of reactive units that are present in a solution.

In acid-base reactions, a reactive unit is simply a proton, #H^(+)#, or a hydroxide ion, #OH^(-)#.

When you have something like hydrochloric acid, #HCl#, which is a strong acid that dissociates completely in aqueous solution to give #H^(+)# and #Cl^(-)# ions, a 0.1 M solution will also be a 0.1 N solution.

This happens because you get 1 mole of reactive units, in this case #H^(+)#, for every 1 mole of #HCl#.

For a diprotic acid like #H_2SO_4#, you get

#H_2SO_((4aq)) -> 2H_((aq))^(+) + SO_(4(aq))^(2-)#

SInce 1 mole of sulfuric acid produces 2 moles of #H^(+)#, a 0.1 M solution will be a 0.2 N solution - you get 2 equivalents of #H^(+)# per liter of solution.

Likewise, when you have a 0.1 M #NaOH# solution, you'll also have a 0.1 N solution, since sodium hydroxide dissociates completely into #Na^(+)# and #OH^(-)# ions.

In this case, the reactive unit is #OH^(-)#.

So, as a conclusion, the normality of a solution depends entirely on what that solution contains; when you have 1 mole of reactive units produced by every 1 mole of a compound, then a 0.1 M solution is also a 0.1 N solution.

Normality was also used for redox reactions and precipitation reactions, the reactive units being the electrons lost or gained for the former, and the number of ions that will precipitate, for the latter.