# Na2CO3 solution

1. Aug 10, 2009

### katchum

Does anyone know if a Na2CO3 solution will change in pH when you leave it for a month? Will the pH go up or down? I thought maybe the CO2 would escape which would make the pH go higher?

2. Aug 10, 2009

### chemisttree

If it dries out it could change the concentration of the sodium carbonate. That might change the pH... If the original solution had no dissolved CO2 in it (in the form of H2CO3) equilibration with the atmosphere would cause more CO2 to dissolve and lower pH. In that case, the system would tend toward bicarbonate (NaHCO3) at the expense of Na2CO3. That might change the pH...

3. Aug 10, 2009

### katchum

Thanks for your info, but I still have some questions:
What is wrong with this deduction?

I've seen this equation:

CO3-- + H2O <=> HCO3- + OH-

Then maybe: HCO3- + H2O <=> H2CO3 + OH-
will occur if you put enough CO3-- into solution. (I'm not sure this will occur with high pH...)

Then H2CO3 could become CO2 + H2O and that will evaporate. Which will make pH go higher because of Le Chatelier.

Just an example:
What will happen to the pH if I make 18% Na2CO3 powder dissolved in water? After a month.
I should do the experiment really...

4. Aug 10, 2009

### chemisttree

Yes, that is why sodium carbonate solutions have an alkaline pH. Remember that the OH- is part of the Keq expression for water (Keq = [H+][OH-]/[H2O] = 14) and by increasing the OH- concentration you must decrease the concentration of H+ correspondingly.

Same reason that bicarbonate solutions are slightly alkaline....

So you will have a solution of Na2CO3 (or 2Na+ and CO3-2) that contains the species NaHCO3, Na2CO3, H2CO3 and NaOH. Add CO2 to that mix and what will you produce?

ie. Na+ + OH- + CO2 <-----> NaHCO3 ?

or (and?)

Na2CO3 + H2CO3 <------> 2NaHCO3 ?

5. Aug 11, 2009

### katchum

I think both reactions will occur. Maybe your point is that with alkaline pH CO2 can never escape because of these two reactions that occur?

6. Aug 11, 2009

### chemisttree

I'm saying that there is an equilibrium that you must consider and that CO2 can be absorbed from the air into a sodium carbonate solution.

Correction: I wrote that Keq =[H+][OH-]/[H2O] = 14

That should have been Keq =[H+][OH-] = 10^-14
Sorry if that was confusing. http://www.chembuddy.com/?left=pH-calculation&right=water-ion-product" [Broken]

Last edited by a moderator: May 4, 2017
7. Aug 11, 2009

### katchum

Well, I have this tank that is routed to open air, so all the CO2 that could exit the solution, would go up and mix with the environment air.

I wonder if the concentration of CO2 in air is high enough to absorb back into the Na2CO3 solution?

I am still in the dark about the answer though. Will pH go up or down? I think I'm starting an experiment...

Last edited: Aug 11, 2009
8. Aug 11, 2009

### Staff: Mentor

Think about it this way - what is pH of the starting solution? At this pH, is it likely to lose carbon dioxide, or adsorb it?

As for your question about amount of CO2 in th eair - pH of the so called 'pure' water in equilibrium with air is about 5.6... guess why

Last edited by a moderator: Aug 13, 2013
9. Aug 12, 2009

### katchum

Let's say 25 kg Na2CO3 + 500 l H2O = 0,47 mol/l CO3^2-
pH = 12

Equation for OH-: x^2/[CO3^2-]=1,8x10^-4

I have no idea how to calculate if CO2 would absorb at this pH...

I've read that with pH > 10 this reaction dominates: CO2 + OH- = HCO3-
So I bet pH will go down.

For the second quiz I have no idea what the equilibrium of 5,6 pH means? The solution was at pH 12, and air doesn't have a pH?

Last edited: Aug 12, 2009
10. Aug 12, 2009

### lightarrow

You didn't specify if the solution's container is left open or closed. Anyway, consider that, even in a closed container, a standard made with such a solution cannot be used for analysis more than some days:

CO32- + CO2 + H2O --> 2HCO3-

Of course the pH lowers because CO32- is more basic than HCO3-.

Last edited: Aug 12, 2009
11. Aug 12, 2009

### Staff: Mentor

You doubted if there is enough CO2 in the air - what I wrote is, if you start with a pure water (be it DI, RO or distilled), it should have have a neutral pH of around 7.00. In reality, if it is in contact with the air, pH of this pure water (not your pH 12 solution) goes down to about 5.6. That's because there is enough carbon dioxide present around.

Last edited by a moderator: Aug 13, 2013
12. Aug 12, 2009

### katchum

Ah, then it's easy, it will go down! Experiment cancelled.

Some background why I asked this.

I want to spray some of Na2CO3 solution onto an acidic powder. If I let the container standing there for a month then I'm sure the solution won't be as effective to neutralize the acidic powder. Now I know I'll have to change the container every week or so, otherwise the powder will still be acidic after spraying it with Na2CO3.

Now a more difficult question is, how long does it take for the pH to drop and by how much... but that's too difficult. And it's an open container.

Last edited: Aug 12, 2009