If 25.0 mL of 0.100 M #HCl# is titrated with 0.150 M #Ba(OH)_2#. what volume of barium hydroxide is required to neutralize the acid?

1 Answer
anor277
Jan 19, 2017

A little less than half the volume of the starting hydrochloric acid.

Explanation:

We need (i) a stoichiometric equation:

#Ba(OH)_2(aq) + 2HCl(aq) rarr BaCl_2(aq) + 2H_2O#,

Which shows us that 2 equiv acid react with the 1 equiv of barium hydroxide.

And (ii) we need the equivalent quantities of acid and base, observing the stoichiometry.

#"Moles of hydrochloric acid"# #=# #25.00xx10^-3*Lxx0.100*mol*L^-1=2.50xx10^-3*mol#

Given the stoichiometry, this molar quantity represents HALF of the molar quantity of barium hydroxide, which of course had an initial concentration of #0.150*mol*L^-1#. So we take the quotient,

#1/2xx(2.50xx10^-3*mol)/(0.150*mol*L^-1)xx10^3*mL*L^-1~=8*mL#

I have not checked the solubility of the barium salt. Barium hydroxide has limited aqueous solubility. Presumably this question is consistent with experiment.

Related questions
See all questions in Neutralization
Impact of this question
31249 views around the world
You can reuse this answer
Creative Commons License