# Is a .1M HCl solution more acidic than a .2M HCL solution ?

1. Mar 8, 2013

### morrobay

With .1M HCl the pH = -Lg 10-1 = 1
With .2M HCL the pH = 2 Since both solutions are low concentrations of a strong acid
they are fully ionized and the Ka HCL = [H+] [A-]/[HA]
is not relevant.
But the .2M HCl solution has higher concentration of H3O+
But by definition the .1 M HCL solution is more acidic ?

2. Mar 8, 2013

### aroc91

Your pH calculation for 0.2M is wrong. -log(0.2) is 0.70, not 2.

3. Mar 8, 2013

### morrobay

Right - You got here first, should be -lg 2 -lg 10-1 = -.3+1=.7 pH
Stupid mistake at my expense

4. Mar 8, 2013

### morrobay

Follow up question : The Ka for HCl = 1 x 103
At what concentration , with increasing concentration , would disassociation of HCl stop.
And the pH no longer decreasing.
Note this is why it is advisable to take Chem 1A and 1B inorganic at the same school.
Otherwise there can be overlap and missing material

5. Mar 9, 2013

### Staff: Mentor

Disassociation or dissociation?

Unless I am missing something, this is crazy difficult. First, we don't have a good theory for calculations at high ionic strengths. Second, in high concentrations ions tend to create pairs (associate). This can change the apparent dissociation constant.

6. Mar 13, 2013

### ajkoer

I disagree with the comment "we don't have a good theory for calculations at high ionic strengths" unless you qualify it as your personal opinion. The theoretical equation and numerical models based thereon (which are non-linear and require interative solutions) to predict the ionic strength of mixed saturated salt solutions is a mini-industry, obviously producing satisfactory results.

I do agree with the statement "This can change the apparent dissociation constant".

Now, to those who know some of this theory, I would ask the question can one add a sufficient quantity of MgCl2 to the .1 M HCl to make it many times stronger than the .2 M HCl, which is absence any other salts?

Last edited: Mar 13, 2013
7. Mar 14, 2013

### Staff: Mentor

As far as I know the best approach to the activity calculations in high ionic strengths at the moment are the Pitzer equation and Specific Interaction Theory. They are expected to be good for ionic strengths up to 5 molal. Problem is, they require experimentally determined interaction coefficients - and these are known only for a limited number of ions pairs, which makes calculations impossible in most cases. So even if these results are satisfactory, they are severely limited.

Thus, while we have a theory that is capable of yielding correct results, we don't have enough data to use it. For most practical purposes it means we can't calculate what we need, and that's what I meant by "we don't have a good theory for calculations".

8. Mar 14, 2013

### ajkoer

Yes, I agree with most of your comments.

For those interested in more details on this topic, please see the following reference (full paper) at http://www.jim.or.jp/journal/e/pdf3/45/04/1317.pdf . To quote from the abstract:

"We developed a chemical model to analyze ionic equilibria in a cobalt chloride solution at 298K. The chemical model consisted of chemical equilibria, mass and charge balance equations. The activity coefficients of solutes and water activity were calculated with Bromley equation. Values of the equilibrium constants for the formation of cobalt chloride complexes at zero ionic strength and of the interaction parameters were estimated by applying Bromley equation to the reported equilibrium constants at different ionic strength".