6B.9 Hw problem

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Michael Du 1E
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6B.9 Hw problem

Postby Michael Du 1E » Tue Jan 14, 2020 4:12 pm

Given that the H3O+ concentration is 1.50 mol/l, why is it that the OH- concentration is 1.50x10^(-14) mol/l? When they are multiplied together, it does not equal the water constant (1.0 x10^-14). Thank you!

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Re: 6B.9 Hw problem

Postby Diana_Diep2I » Tue Jan 14, 2020 4:36 pm

I think we get that when the pH scale exceeds 14. This usually occurs when the molarity or concentration of H3O+ or OH- is greater than 1.

Bryce Ramirez 1J
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Re: 6B.9 Hw problem

Postby Bryce Ramirez 1J » Tue Jan 14, 2020 10:44 pm

I also did the math for this problem and the book said my answers were wrong, though I think the book is wrong. When calculating the pH for part a using 1.5 moles of H30, my calculator gives the pH as -.176. The answer shows it being positive .176, which is wrong. If 1 mole of Hydrochloric Acid (HCI) has a pH of 0, then how could an acid with a higher amount of moles have a more basic pH (.176 < 0).

Ariel Davydov 1C
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Re: 6B.9 Hw problem

Postby Ariel Davydov 1C » Wed Jan 15, 2020 12:19 am

When calculating pH, if you end up with a negative value you must convert it to its positive form, which will be the answer. Since the pH scale ranges from 0 to 14, there cannot be a “true” negative pH value. This can be explained by the fact that strong acids at concentrations higher than 1.00 mol*L^-1 typically do not entirely dissociate in water due to being in extremely concentrated form. Thus, if you end up with -.176 for your pH value, your “true pH value” will be the positive form of the number, which is .176. Hope this helped!

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