Talk:RICE chart

Needs some qualifiers

There's no mention of the restricted applicability of this method for calculating pH. There's an underlying assumption that [H+] >> [OH-]. Another way of saying this is that this only applies to solutions whose pH is much less than 7. Klortho (talk) 05:34, 13 March 2011 (UTC)[reply]

SOMEONE DO SOMETHING ABOUT THIS PAGE

Someone who knows HTML needs to create an ICE table to demonstrate to others what exactly is an ICE table... --XH 20:55, 20 January 2007 (UTC)User:Xinyu

I've now done this, basically by paraphrasing the material below. I hope LestatdeLioncourt does not mind. Petergans 19:05, 17 May 2007 (UTC)[reply]

pH value from dissociation constant

The following was from the March 12 Science Reference Desk.

This is based on a homework question with specific values, but I am asking only for general formula.

If I know the acid dissociation constant (pKa) and the concentration of a weak acid, how would I find the pH? How do I find the pH after adding a certain volume of a strong base? − Twas Now ( talk • contribs • e-mail ) 00:58, 12 March 2007 (UTC)[reply]

Take a look at Henderson-Hasselbalch equation

{\textrm {pH}}={\textrm {pK}}_{a}+\log _{10}{\frac {[{\textrm {A}}^{-}]}{[{\textrm {HA}}]}}

From the pKa page the ionised acid and H+ concentration will be the same, so you end up with pH+pH=-log10(concentration of a weak acid)+pKa.

{\textrm {pH}}=({\textrm {pK}}_{a}-\log _{10}{[{\textrm {HA}}]})/2

GB 02:18, 12 March 2007 (UTC)[reply]

Thanks. I guess what confuses me is how do I determine [HA] and [A^-] for a given weak acid (and its conjugate base)? − Twas Now ( talk • contribs • e-mail ) 03:09, 12 March 2007 (UTC)[reply]

It's really very simple. All you need to do is set up an ICE table for the ionization equilibrium of the acid, i.e. HA ⇌ A^- + H⁺ :

	[HA]	[A^-]	[H⁺]
I	c₀	0	0
C	-x	+x	+x
E	c₀ - x	x	x

Because $pH=-\log {[H^{+}]}$ , and you can tell from the table that when the ionization is over [H⁺] = x, then $pH=-\log {x}$ . So, your job is to find x.

The equilibrium constant expression for the ionization is:

$K_{a}={\frac {[H^{+}][A^{-}]}{[HA]}}$

Substitute the concentrations with the values found in the last row of the ICE table.

$K_{a}={\frac {x.x}{c_{0}-x}}$

Now plug in the specfic values for c₀ and K_a ( $K_{a}=10^{-pK_{a}}$ ) provided in your question, then solve for x, and you're done!

If you do some math, you can quickly figure out that the relation GB provided is only a different (and handier) way of expressing K_a, especially when you're working with buffer solutions.

Now for the second part of your question: when you added a certain quantity of a strong base, the OH^- ions you just added will react with some of the H⁺ in solution. From the OH^- - H⁺ reaction stoichiometry, you can calculate the number of moles of H⁺ that have disappeared and how many moles are left. Then you can find the new concentration of H⁺ ions (keep in mind the change in volume) and the new pH. —LestatdeLioncour t 14:45, 12 March 2007 (UTC)[reply]

Typesetting

Terrible (and inconsistent) typesetting in this article. Needs to be fixed. —DIV (1.129.104.52 (talk) 12:02, 18 August 2019 (UTC))[reply]