What volume of a concentrated HCl solution, which is 36.0% HCl by mass and has a density of 1.179 g/mL, should be used to make 5.10 L of an HCl solution with a pH of 1.7?

1 Answer
anor277
Mar 30, 2018

This is an illustrative question...and excuse me while I belabour the points and definitions...and there is also a very important ISSUE of safety to consider.

Explanation:

By definition, #"concentration"="amount of SOLUTE"/"amount of solution"#...and normally we use units of #mol*L^-1# or #g*L^-1#... Here we probably should first find the concentration of the mother solution in #mol*L^-1#...and so...working from a #1*mL# volume...

#[HCl(aq)]=((1*mLxx1.179*g*mL^-1xx36%)/(36.46*g*mol^-1))/(1*mLxx10^-3*L*mL^-1)=11.64*mol*L^-1#...this is concentrated stuff...the conc, #HCl# we use in the lab is #10.6*mol*L^-1#.

And we want #5.10*L# of a solution whose #pH=1.7#... Now, by definition...#pH=-log_10[H_3O^+]#...i.e. #[H_3O^+]=10^(-1.7)*mol*L^-1=0.0200*mol*L^-1#...and so...

#"moles of HCl required"-=5.10*Lxx0.0200*mol*L^-1=0.102*mol#..

#"Volume of conc. acid"=(0.102*mol)/(11.64*mol*L^-1)=0.00874*L#, i.e. #8.70*mL#.

And #"YOU MUST ADD ACID TO WATER"# and never the #"REVERSE"#. Why not? Because if you spit in concentrated acid, it spits back. I kid you not....

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