Salifert KH/Alk testing help

[JimWelsh]

I agree that there are differences of opinion. Dickson came up with pH 4.5, and if I plug the raw data from that paper into the calculator I cite above, the inflection point method finds the endpoint at pH 4.45, FWIW (regardless which conductivity value I provide).

 

[JimWelsh]

Of course, I couldn't just let this rest, so I satisfied my curiosity, and sorted the math on this out.

The exact definition of total alkalinity (Dickson, 1981) is:

TA = [HCO3-]+2[CO3--]+[B(OH)4]+[OH-]+[HPO4--]+2[PO4---]+[SiO(OH)3-]+[HS-]+[NH3]-[H+]-[HSO4-]-[HF]-[H3PO4]

At any point in the titration, the alkalinity can be expressed as:

(V0*TA - V*N)/(V0+V) = [HCO3-]+2[CO3--]+[B(OH)4]+[OH-]+[HPO4--]+2[PO4---]+[SiO(OH)3-]+[HS-]+[NH3]-[H+]-[HSO4-]-[HF]-[H3PO4]

where V0 is the original sample volume, N is the normality of the titrant acid, and V is the volume of acid added. At the equivalence point, the term (V0*TA - V*N) becomes zero, so we can rearrange the formula to show that the net charge at the endpoint is exactly equal to [H+]:

[H+] = [HCO3-]+2[CO3--]+[B(OH)4]+[OH-]+[HPO4--]+2[PO4---]+[SiO(OH)3-]+[HS-]+[NH3]-[HSO4-]-[HF]-[H3PO4]

Now, my goal is to calculate the pH of this equivalence point to, say, 0.01 pH unit. At a pH of 4.60, a change of 0.01 pH units means a change in [H+] of around 1/10^4.60 - 1/10^4.61 = 5.7E-07. At a pH of 4.20, this 0.01 pH unit change means a change in [H+] of around 1.4E-06. So, within the range from pH 4.20 to 4.60 where we suspect the equivalence point will occur, we can safely disregard any of the factors in the equation above that will contribute significantly less than 5.7E-07 to the total. It can be demonstrated that within this pH range, for seawater, all of the terms except for [H+] and [HCO3-] are at least an order of magnitude smaller than 5.7E-07. This fact allows us to substantially simplify the formula to this, while remaining within 0.01 pH units of the correct answer:

[H+] = [HCO3-]

[HCO3-] can be calculated from the total carbonate alkalinity Ct, if the carbonic acid dissociation constants pKa1 and pKa2 are known. Dickson and Millero (1987) have published formulas for calculating these pKa in seawater based on temperature and salinity (http://www-naweb.iaea.org/napc/ih/documents/global_cycle/vol I/cht_i_09.pdf). Assuming S=35 and T=20C, we get pKa1 = 5.882 and pKa2 = 9.035. For a given total carbonate alkalinity Ct, [HCO3-] can be shown to be equal to Ct*[H+]*Ka1/([H+]^2+[H+]*Ka1+Ka1*Ka2), so that means that at the equivalence point, we have:

[H+] = Ct*[H+]*Ka1/([H+]^2+[H+]*Ka1+Ka1*Ka2)

Rearranging this, we get:

[H+]^2+[H+]*Ka1+Ka1*(Ka2-Ct) = 0

This is a simple quadratic equation where a = 1, b = Ka1, and c = Ka1*(Ka2-Ct). Assuming a Ct value of, say, 7 dKH, and using the pKa1 and pKa2 values for S=35 and T=20C, solving the quadratic equation for [H+] gives an answer of 5.665E-05, for a pH value of 4.25.

So the theoretical equivalence point for S=35, T=20C, and Ct=7dKH would happen at a pH of 4.25.

Now, this assumes zero dilution of the sample by the titrant, which we know is not possible, so we need to adjust the pKa1 and pKa2 values appropriately for this dilution factor. Taking the Salifert alkalinity test as an example, we would start with 4 mL of sample, but then it would be diluted by the 4 drops of indicator (figure 0.05 mL per drop), plus the 0.46 mL of acid that it would take to titrate 7.0 dKh, for a total volume at the equivalence point of 4 + 4*0.05 + 0.47 = 4.67 mL. So, our S=35 sample would become S=35*4/4.67=29.98. Using S=29.98 to calculate pKa1 and pKa2, we get pKa1 = 5.903 and pKa2 = 9.089. Solving for [H+] using these constants still gives a pH of 4.26.

Regarding my earlier titration curves where I was reporting a value closer to pH 4.6 for the endpoint, I was seriously omitting the dilution factor. Per the Hach alkalinity burette method, I was taking 10 mL of sample and diluting it to appx. 50 mL, and was also using a rather dilute titrant, so that the total volume at the endpoint was around 95 mL. The resulting salinity would have been S = 35 * 10 / 95 = 3.68, which in turn would give pKa1 = 6.111 and pKa2 = 9.469. Even still, the calculated equivalence point pH would be 4.31 in this case. Not sure how I managed to arrive at a pH of 4.6 -- perhaps it was a pH probe calibration issue?

So, assuming that we can trust Dickson and Millero's calculations for pKa1 and pKa2, the theoretical endpoint does appear to fall between pH 4.2 and 4.3.

 

[chipmunkofdoom2]

Good discussion here, thanks Randy and Jim.

As far as "when is the titration done" with regard to the Salifert test kits, I drip until the sample is fully pink with no hints of blue/green. I was told by several friends in academia that as soon as the sample begins to change colors, that is when the titration is complete. However, as @Scott.h mentioned, there is only a 0.1 - 0.3 difference between the beginning of the color change to the change to completely pink. Even if my carbonate alkalinity is a bit lower than my test says as a result of using more titrant, I don't think it matters much if I am testing consistently each time.

 

[gerald91]

All the test a wrong.
In my point of view the most important is to follow always the same protocole. And 1 time per quater do a laboratory test .. after you know the gap beetwen your protocole test vs laboratory result

 

[Tennsquire]

My Salifert alk test came with a bottle of reference solution. If you have that, test it so you know what the color looks like at the designated endpoint.

 

[Mindi]

Interesting discussion. I sometimes do Calcium with both Salifert and Hanna Checker. The Hanna figure consistently matches the Salifert figure if you stop the Salifert titration at the first hint of blue. If you make your end point clear blue you consistently get 15-20ppm higher reading than Hanna (checked against their 400ppm standard). Not suggesting this directly relevant to the Alk discussion.

 

[Hans-Werner]

In Germany another expression is Säurekapazität pH 4,3 (acid capacity pH 4.3) expressed in mmol/l hydrochloric acid. So per definition it should be pH 4.3.
At pH 4.3 all carbonates and bicarbonates which form the alkalinity are destroyed and converted to CO2 which should be given the chance to degas from the sample by shaking carefully for a while. This may cause the color to turn back (to blue in this case). If so another drop of acid has to be added.

 

[Big E]

I appreciate all the comments but my question doesn't seem to get answered which is when do you stop adding reagent? When the blue is completely gone and only the pink is visible or when the color change begins and you can still see the blue and pink together?

I take note of both colors and just use a number that falls directly between the two.numbers. The important point is to just consistently do it the same way. I seriously doubt using one or the other is going to matter, just be consistent what you do.

I started to do it this way because I would test against my La Motte test kit and the Salifert consistently matched by me taking that median number.

 

[Clittrell]

This is slightly over my head but the interesting non the less. I use the Hanna alk checker and it seems to be the same as above. You get the same number as when you get the first color change on the salifert. Does anyone know how they came up with the color chart that they program in to the unit. The thing about the Hanna is I always get the same number if I retest. Something I seem to have more issues with when I use the salifert.

 

[Rick Mathew]

It is most likely they did not use a color chart at all for the Hanna Checker. They probably made know solutions and created a calibration curve and that information is what is used by the checker to do the measurement...That wouldbe my guess

Rick

 

[Randy Holmes-Farley Verified Community Expert]

I do not believe that Hanna has stated exactly what the alk checker analyzes.

 

[Rick Mathew]

The Hanna Literature says they use the colorimetric method with a 610nm LED as the source and silicone photocell as the detector...sounds like light absorption method of some sort...

 

[Randy Holmes-Farley Verified Community Expert]

Yes, thanks. It would be nice to know more, such as how much acid they add, etc.

 

[JimWelsh]

The calculator may (or may not) be accurate in seawater. It askes for conductivity to:

"The specific conductance is used to estimate the ionic strength of the sample. Ionic strength and sample temperature are used to calculate the equilibrium dissociation constants of water and carbonic acid, which are then used in the calculation of hydroxide, carbonate, and bicarbonate concentrations."

But the pKa of carbonate and bicarbonate in NaCl of similar conductivty is not the same as it is in actual seawater.

Sorry to resurrect this old thread, but the subject of the exact color of the Salifert alk test endpoint has once again come up in another thread.

Randy, wouldn't the influence of the seawater matrix also have a similar effect on the pKa of the indicator dye(s) as it does on carbonate and bicarbonate? And wouldn't this tend to cause the dye color at the true equivalence point to be similar for various matrices?

Saying it another way: Just as the seawater matrix lowers the pKa of carbonate and bicarbonate, wouldn't it also similarly lower the pKa of bromocresol green and methyl red, and just as the species distributions of the carbonate species are shifted towards a lower pH by the lowered pKa, then, too, the species distribution of the two different colored forms of the dye(s) would also be shifted to a similarly lower pH, causing the same color to appear at a lower pH in seawater as the color in freshwater at a higher pH?

 

[Randy Holmes-Farley Verified Community Expert]

Sorry to resurrect this old thread, but the subject of the exact color of the Salifert alk test endpoint has once again come up in another thread.

Randy, wouldn't the influence of the seawater matrix also have a similar effect on the pKa of the indicator dye(s) as it does on carbonate and bicarbonate? And wouldn't this tend to cause the dye color at the true equivalence point to be similar for various matrices?

Saying it another way: Just as the seawater matrix lowers the pKa of carbonate and bicarbonate, wouldn't it also similarly lower the pKa of bromocresol green and methyl red, and just as the species distributions of the carbonate species are shifted towards a lower pH by the lowered pKa, then, too, the species distribution of the two different colored forms of the dye(s) would also be shifted to a similarly lower pH, causing the same color to appear at a lower pH in seawater as the color in freshwater at a higher pH?

Yes, but not necessarily to the same degree.

Salinity impacts the pKa of different molecules differently. Salinity tends to reduce the effects of charge-charge interactions, and also makes it easier to more highly charge individual ions.

For example, in Table 8.2 of Millero, he shows that the third pKa of phosphate is 3.65 pH units lower in seawater (8.69) than in fresh water (12.34), while the first pKa is only 0.58 units lower (1.57 vs 2.15). Salinity similarly shifts the pKa of carbonic acid less than bicarbonate in seawater vs fresh.

Indicator dyes generally do not have multiple charges in a single atom, but they can have multiple charges in the molecule, and so the salinity may impact how easily those charges are made differently in different dyes.

 

[JimWelsh]

The paper "Determination of seawater pH from 1.5 to 8.5 using colorimetric indicators" says that the pKa of bromocresol green in 35 PPT seawater at 25C is 4.410. Various references for its pKa in freshwater range from 4.7 to 4.9, making the shift downward for the dye in seawater to be between 0.3 to 0.5 pH units, which is just about right for the difference between the pH values of the endpoint ranging from 4.6 in freshwater to 4.2 in seawater being thrown around by the various sources and calculations.

So, while the downward shift in the dye's pKa due to the seawater matrix might not correspond *exactly* to the downward shift in the endpoint pH, it looks to be very close.

 

[Randy Holmes-Farley Verified Community Expert]

The paper "Determination of seawater pH from 1.5 to 8.5 using colorimetric indicators" says that the pKa of bromocresol green in 35 PPT seawater at 25C is 4.410. Various references for its pKa in freshwater range from 4.7 to 4.9, making the shift downward for the dye in seawater to be between 0.3 to 0.5 pH units, which is just about right for the difference between the pH values of the endpoint ranging from 4.6 in freshwater to 4.2 in seawater being thrown around by the various sources and calculations.

So, while the downward shift in the dye's pKa due to the seawater matrix might not correspond *exactly* to the downward shift in the endpoint pH, it looks to be very close.

Yes, it might be.

 

[JimWelsh]

I have started another thread that is at least partially relevant to this thread: Alkalinity titration endpoint pH & Salifert endpoint color