# Why does my Patton-Reeder indicator immediately turn blue? (Ca2 conc.)

## Main Question or Discussion Point

I was trying to use the ethylenediaminetetraacetic acid (EDTA) method using this simple guide to measure the calcium ions in my CaCl2 (2.4 mM) + NaHCO3 (6.4 mM) solution. But the solution immediately turns dark blue (like really dark) when I add the indicator to the calcium solution.

From the guide; excess calcium ions form a complex with the PR indicator (pink), but then as the EDTA is added they start to form a complex with EDTA and it turns blue.

$[Ca-PR] + EDTA^{4-}\rightarrow PR + [Ca-EDTA]^{2-}$​

The background to this comes from some papers that say the EDTA method is via magnesium precipitatation as magnesium hydroxide by addition of potassium hydroxide. Then, calcium forms a stronger complex with EDTA than magnesium. Where does the magnesium fit into the above reaction? Where's the potassium?

Now, there is no magnesium or potassium in my calcium solution, so I believe this may be the problem.

My PR indicator immediately turns dark blue when added to my calcium solution. Do I need to add magnesium to my solution and use potassium hydroxide? However, nowhere in the "Solutions Needed" part do they mention magnesium. The guide is quite confusing.

Also, in order to make the calcium solution, NaOH (0.1 N) is added to the $NaHCO_{3}$, to get the pH to 9.1. Then it is mixed with the $CaCl_{2}$ resulting in a final pH of 8.8. In my experiment the pH was measured at 12.78 and it stayed at that value when the $CaCl_{2}$ was added (no idea why, will try using less NaOH). Could the pH have been to high?

Please could someone give me a point in the right direction? What do I need to do to my calcium solution in order to use the EDTA method?

## Answers and Replies

Borek
Mentor
Potassium and sodium don't matter here, they don't react with neither EDTA nor the PR indicator (or, to be more precise, stability constants of these complexes make them negligible). NaOH or KOH is added to keep pH as high as possible. That serves two purposes - it makes the EDTA/Ca reaction easier (EDTA is a weak acid, ionizing it in a high pH is beneficial), but more importantly it masks the Mg precipitating it as Mg(OH)2. If there is no Mg present, added base just keeps the pH high.

Calcium solution made by using carbonate buffer sounds off, as it will precipitate calcium, removing it from the solution. Try if preparing calcium solution without a buffer won't make the titration work.

Potassium and sodium don't matter here, they don't react with neither EDTA nor the PR indicator (or, to be more precise, stability constants of these complexes make them negligible). NaOH or KOH is added to keep pH as high as possible. That serves two purposes - it makes the EDTA/Ca reaction easier (EDTA is a weak acid, ionizing it in a high pH is beneficial), but more importantly it masks the Mg precipitating it as Mg(OH)2. If there is no Mg present, added base just keeps the pH high.
Thank you Borek. Yes, there should be no (or negligible) Mg as I used DI to make the calcium solution. The pH was certainly high!

Calcium solution made by using carbonate buffer sounds off, as it will precipitate calcium, removing it from the solution. Try if preparing calcium solution without a buffer won't make the titration work.
Trying the $CaCl_{2}$ solution is a good idea, for me to learn the process, I will go do that now thanks.

The reason for the bicarbonate is to make a calcium carbonate saturated solution, for precipitation purposes. NaOH is added to the bicarbonate as here:

Paper here for your interest.​

It actually works pretty well! For example, I compared $CaCl_{2} + Na_{2}SO_{4}$ solution to $CaCl + NaHCO_{3} (+NaOH)$ solution, the precipitation is significantly reduced. I am now just trying to measure the calcium ion concentration to confirm enough excess $Ca^{2+}$ ions in solution for my process.

Ok the PR with the $CaCl_{2}$ (pH 7) turned blue, the PR with $CaCl_{2}$ and NaOH (pH 12) turned the pink I need. pH 12 was about the pH of my calcium solution that I made. So I made a new solution, adjusted to pH 12, and this time it turned pink. I added the EDTA, then it turned blue. No idea what went wrong the first time, maybe too much indicator so the pink looked purple / blue. Anyway, now back on track!

Thanks Borek, otherwise I would have ended up with some concoction of magnesium, potassium, sodium etc etc.

Borek
Mentor
Most of the indicators used in complexometric titrations are also reacting to pH changes. I am not saying that's the case here, but it is something to always consider.