pH and Molarity ```Name: Jose Status: student Grade: 9-12 Location: N/A Country: N/A Date: April 2007 ``` Question: Dear Argonne scientists, I am a 11th grade student in Peru and i will carry out an experiment showing how P.H of water used for watering certain bean plans would affect their growth (height, biomass etc). The problem is, we can't use pH papers, litmus papers, so we will have to measure its acidity or how alkaline it is by Molarity. If i want to chose 6, 6.5, 7, 7.5, 8 as my range of pH's, then how can i calculate the Molarity of each? Replies: Jose Luis, You may find this process a little bit difficult in terms of precision if you are using a strong acid. The problem is that you are shooting for pH's that are very nearly neutral (pH = 7) and so you would have to obtain very dilute strong acids in order to get a pH in the range of 6 - 6.5. The equation is from the definition of pH which is: pH = -log([H+]) where [H+] = M of the strong acid and converting this to solve for the target molarity you would use: 1 / 10^(pH) = M and as you can see, to get a pH of 6.5 you would need a M that is 3.16 x 10^-7. And as you can see, this is very dilute and very hard to obtain. On the other hand, if you use a weak acid then it's still pH = -log([H+]), but the [H+] this time is not just the molarity of the weak acid. Rather: [H+] = Square root of (ka x M of the weak acid) Thus, if you wanted to get a pH = 6.5, you would need an [H+] = 3.16x10^-7 which is equivalent to: ([H+])^2/ka = M and depending on the ka of your acid, you would still be working with very dilute solutions. For example if you use acetic acid which has a ka = 1.8 x 10^-5 then you would need a M that is 5.6 x 10^-9. For the range you are looking for, you might be better off using solutions of weak acid and weak base SALTS or buffer solutions. You can find tables of typical salt solutions or buffer solutions to prepare particular pHs, rather than to dilute acids. Greg (Roberto Gregorius) Click here to return to the Chemistry Archives

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