[Tobias]

Here in the Netherlands every High School student has to do a PWS (profielwerkstuk), which is basically a bit of research you have to spend 100 hours on.

Every year our school has fieldwork (basically doing a bit of research in the outdoors) and part of that is an acid-base titration of soilwater or water from a creek. Now the problem with that titration is that it usually goes wrong. This is because the water you have to determine the pH of is too close to 7 (my teacher said something about the "waterbalance" starting to play a role).

My partner and I decided to try to find out how to get better results from the titration for our PWS. But since we can't find out on our own how to fix it I decided to take a look online and here I am.

So basically my question is: how can you get better results from an acid-base titration when the pH is close to 7?

-Tobias

[Borek]

Explain what you are trying to determine with the titration, as it is not clear from your post; at least to me.

[Tobias]

Well, basically just the pH of the (ground)water. But there aren't a lot of H⁺ (or H₃O⁺ if you wish) particles to begin with so it's quite difficult to get clear results.

I hope this clarifies it a bit.

-Tobias

[Borek]

You can't determine pH with titration. Use pH meter.

[kevins]

If you want to perform the titration with indicator,
1. To find out a indicator (or indicator mixture) which change the colour close to 7.
2. Use the very dilute titratant.

Please try.

[Tobias]

I know you can use a pH-meter to determine the pH, but the point of our "fieldwork" is using things you learned in the classroom. And they want us to use a titration to determine the pH of the water.
But I don't see what you mean with "You can't determine pH with titration". With a titration you can determine the equivalence point and on that point [H⁺]=[OH^-]. And if you know the concentration of H⁺ you can calculate the pH (= -log[H⁺])

Kevins:
We prefer to use a pH-meter to determine the equivalence point instead of an indicator. And we already tried using a very dilute titrant, but it still gave foggy results (at best).

To give an example of what I mean:
Lets say we have some groundwater. We determine whether it's accid or basic with an pH-meter (we need to do this to determine which titrant to use). We find that the pH is 6.2. That means that the concentration is 10^-6.2=6.31E-7 mole/liter.

We take 100ml of that sample to use in the titration. That means there are 0.1*6.31E-7=6.31E-8 H⁺ mole present. Lets use a very dilute titrant, 0.001M NaOH. If we put 1mL of that titrant in our sample it means that we put 0.001M*0.001L= 1E-6 mole in our sample of water. Which is a lot more than the 6.31E-8 we had in our sample.

You see the problem? We would have to dilute the titrant so much it's probably not going to be accurate anyway.

If anyone has any solutions it would be really appreciated.

[Borek]

I know you can use a pH-meter to determine the pH, but the point of our "fieldwork" is using things you learned in the classroom. And they want us to use a titration to determine the pH of the water.

Sorry to say that, but it means they know nothing about chemistry. General chemistry to be precise, as that's the moment where everyone learns about buffers. You have already adressed one of the problems in your post - need for using insanely diluted reactants. But that's the last thing I will be afraid of.

But I don't see what you mean with "You can't determine pH with titration".

OK, pH = -log[H⁺] - that's what you have posted, and (while that's not the whole truth) that will be our starting point.

Your water contains dissolved carbonic acid - you can be sure it does, each solution that is in contact with air contains some dissolved carbon dixoide, which means it contains carbonic acid. Carbonic acid is a weak acid, that is dissociated only partially:

H₂CO₃ <-> H⁺ + HCO₃^-

There is also second step of dissociation, but at pH around 7 it doesn't play any role so we can ignore it for now. Equilibrium of this dissociation is described by acid dissociation constant:

K_a = [H⁺][HCO₃^-]/[H₂CO₃] = 4.27x10^-7

This equation can be rearranged to form

pH = pKa + log([HCO₃^-]/[H₂CO₃])

(pKa is just -log(Ka), by analogy to pH) or (using given Ka value)

pH = 6.37 + log([HCO₃^-]/[H₂CO₃])

and it is called then Henderson-Hasselbalch equation.

As you see pH of the solution depends on the ratio of concentrations of carbonic acid and hydrogen carbonate anion. When you add titrant - strong base - to your solution, you are not neutralizing H⁺ (in which case you will be able to determine [H⁺] and pH). You are neutralizing carbonic acid - thus changing ratio of [HCO₃^-]/[H₂CO₃]. However, the higher concentration of carbonic acid, the slower changes in pH. That's so called buffering effect - pH of the solution that contains weak acid and its conjugated base (H₂CO₃ is a weak acid, HCO₃^- is its conjugated base) changes very slowly on the addition of bases and/or acids.

Why does it matter? Well - result of your titration won't have ANYTHING to do with the solution pH. You will titrate all buffer systems present in the water - and you can be sure natural water contains not only carbonic buffer, but also some humic acids that'll do their best to keep pH unchanged on addition of titrant. With titration you can - at best - determine buffer capacity of the system.

Let's assume you start with 25 mL of solution containing 0.001M carbon dioxide. That means that sum of concentrations of carbonic acid and hydrogen carbonate is always 0.001M. Let's assume we start with pH 6.00. At this pH concentrations of carbonic acid and hydrogen carbonate are 7x10^-4M and 3x10^-4M (put these numbers into Henderson-Hasselbalch equation to check, note that 7x10^-4M + 3x10^-4M = 0.001M). Now, we want to end our titration at pH 7 - at this pH concentration of carbonic acid will be 1.9x10^-4 and concentration of hydrogen carbonate 8.1x10^-4M (check the numbers again with Henderson-Hasselbalch equation, check if the sum is still OK). To change pH we had to neutralize carbonic acid to change its concentration from 7x10^-4M to 1.9x10^-4M. As the volume was 25 mL we had to add 0.025x(7x10^-4-1.9x10^-4)=1.275x10^-5 mole of strong base.

Now you do the same calculation assuming there is no carbonic acid in the solution - check what will be the final pH of the pH 6.00 solution that contains "only" H⁺ once you add 1.275x10^-5 mole of strong base to 25 mL. You see why it won't work?

Read also these pages:

http://www.chembuddy.com/?left=pH-calculation&right=introduction-acid-base-equilibrium

http://www.chembuddy.com/?left=pH-calculation&right=bronsted-lowry-theory

[Tobias]

I see what you mean.  And I don't mind you doubting the chemical abilities of my teacher and the Technical Education Assistant (he does all the practicals), they're not my favorite people to say the least.

But we have titrated weak acids/bases and it worked. The equivalence point wasn't at pH=7 when you titrate weak acids/bases, but titration curves did tell us where the equivalence point was. But I must say I'm a bit rusty when it comes to buffers. I take it not all weak acids and their conjugated  bases are buffers?
This is not really part of the problem, but I want to know; if not all acids and their conjugated bases are buffers, then what does it need to "be"  in order to act as a buffer?

[Borek]

For most practical purposes all weak acids and bases can be used for buffer preparation. If they are too weak or too strong it becomes tricky, but still possible if you know how.

Titration of weak acid doesn't tell you its pH - it tells you amount of that acid. That's completely different thing. In fact it will be even a better example than what I wrote yesterday. If as a result of titration you have found that your solution contains 0.01M of some acid it doesn't mean pH was 2 - it could be as well 3.4 for acetic acid, or 3.1 for benzoic, or 2.9 for formic - and so on.

If you are looking for some titration that really makes sense you may consider determination of water hardness with EDTA. Not especially easy (but not that difficult) and at least you can be sure your result have some practical meaning.

[Tobias]

We already did an EDTA titration (just part of a regular practical, not fieldwork). But the point of the fieldwork is that you examine an area and learn as much as you can about it. And they insist that we calculate

But just to summarize what you said (I like bringing order to stuff in my head):
1. If you titrate a weak acid (which is almost always a buffer) and you determine the equivalence point, you determined how much of the acid+conjugated base were present in the solution
2. If that solution had been "pure"  you could have calculated the pH with Ka (but don't you need the ratio of HA:H⁺+A_- for that? and how do you get that ratio?)
3. But since groundwater isn't a "pure" solution you can't determine the pH of the water with a titration

If there's something wrong in my summary please correct it. I get what you're saying (most of it anyway), but I want to understand it completely.

[Borek]

We already did an EDTA titration (just part of a regular practical, not fieldwork). But the point of the fieldwork is that you examine an area and learn as much as you can about it.

Water hardness is a very important parameter - so IMHO it fits perfectly. Just like TOD (total oxygen demand) and pH it describes water quality. For example many species of fish can survieve in hard water (well, not very hard) but they are not able to breed. See any FAQ on breeding fish in aquarium.

And they insist that we calculate

1. If you titrate a weak acid (which is almost always a buffer) and you determine the equivalence point, you determined how much of the acid+conjugated base were present in the solution

No - after titration you don't know amount of conjugated base, just that of the acid not neutralized earlier. Check discussion of the carbonic acid at pH 6.00 - already 30% of analytical concentration of carbonate is in the form of conjugated base.

2. If that solution had been "pure"  you could have calculated the pH with Ka (but don't you need the ratio of HA:H⁺+A_- for that? and how do you get that ratio?)

If the solution has been that of pure, not neutralized acid you can calculate initial pH (assuming you know its pKa). That's just pH of a weak acid/url]. But in field you will always have a solution containing mixture of acid and conjugated base, moreover, water will contain mixture of several different acids.

3. But since groundwater isn't a "pure" solution you can't determine the pH of the water with a titration

Exactly. pH is pH, acid amount is acid amount.

[Tobias]

Sorry about that sentence that ends so abruptly. What I wanted to say was that they want us to determine the pH of groundwater (and they insist on a titration, but maybe we can change that one to the EDTA one).

But there is still a small thing I don't get about titrating buffers.

Lets say you have the weak acid HA and it forms the following equilibrium =

HA <--> H⁺ + A^-

Now we titrate this buffer with a NaOH solution.

So wouldn't the OH^- react with both the HA and H⁺?

OH^- + HA --> H₂O + A^-
OH^- + H⁺ --> H₂O

Or is my reasoning flawed somewhere? Because if it isn't you determine the amount of HA+A^- if you titrate...

I'm very grateful for your help and patience.

EDIT:
The acidity of ground and surface water is mainly determined by H₂CO₃ right? How big a difference would it make if we would ignore all the other buffers and just focus on the H₂CO₃ and calculate the pH that way?

[Borek]

they want us to determine the pH of groundwater (and they insist on a titration

As I told you - it doesn't make any sense.

So wouldn't the OH^- react with both the HA and H⁺?

OH^- + HA --> H₂O + A^-
OH^- + H⁺ --> H₂O

Or is my reasoning flawed somewhere? Because if it isn't you determine the amount of HA+A^- if you titrate...

If you have only acid in the solution - you are right. But if you have mixture of HA and - say - NaA, you will be not able to determine A^- from NaA dissociation, only that from HA dissociation.

The acidity of ground and surface water is mainly determined by H₂CO₃ right?

Not necesarilly. Usually it contains also so called humic acids.