Chemical Forums
Chemistry Forums for Students => Inorganic Chemistry Forum => Topic started by: ahmed_saleh on June 08, 2008, 07:15:55 AM
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can any one advise please:
What is the expected range of pH of the Sodium Bromide (44.7 % wt. aqueous solution with a density of 12.4 lb/gal = 1486 kg/m3)? Is it expected that the pH changes during storing of this product?
My question is based on the fact that we produce a NaBr batch with pH ~ 7, but after some time of storage it gets down to about 6.5, even when raising the pH again by using an NaOH solution to 7 or a litte higher it gets down again to 6.5 after storing! Thanks for assistance...
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Is it stored in a sealed container or open to the atmosphere?
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How do you measure pH of this syrup?
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DrCMS,
It is stored in an atmospheric tank (opened to atmosphere through the goose neck vent only). Do you think about absorbing CO2 from the atmosphere, which may affect/lower the pH value? If so, what is the best approach to kill this issue? Thanks for your assistance.
Borek,
pH is measured in the lab using an instrumental pH-meters. The pH of the solution can be measured by two approaches: directly or after 1:10 dilution. If the issue is related to absorption of atmospheric CO2, what would be the magnitude of such effect for the both approaches of measurement.Thanks for your assistance.
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Borek,
I have found a topic & was about: " Is NaCl a base?". You & AWK have responed to this topic, where I found the answers interesting. based on this information, I have the following questions:
1-Would NaBr behave in the same way as NaCl :[ pH of NaCl solutions will be slightly below 7, but this is caused by "ionic strength effect" (influence of ionic strength on activity of H+ ion) ] ?
2-What is the effect of the solution's concentration on this phenomena?
3-What is the expected relation/trend between the Halide (as F,Cl,Br & I) of Na & the magnitude of this effect?
Please advise & thanks for your assistance.
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Would NaBr behave in the same way as NaCl :[ pH of NaCl solutions will be slightly below 7, but this is caused by "ionic strength effect" (influence of ionic strength on activity of H+ ion) ] ?
Yes, I would expect exactly the same problems here. Note - they will be not time dependent, ie pH will not change in time. If the pH changes it is most likely due to some additional effect, like CO2 contamination.
What is the effect of the solution's concentration on this phenomena?
Good question. Most often used theory (Debye-Huckel theory (http://www.chembuddy.com/?left=pH-calculation&right=ionic-strength-activity-coefficients)) works OK for solutions of ionic strength up to 0.1, with some extensions they can be used up to ionic strength 0.5 (in the case of Me+X- salt ionic strength is identical with its molar concentration). Your solution is far more concentrated. Most likely activity coefficients are much higher than 1 then.
What is the expected relation/trend between the Halide (as F,Cl,Br & I) of Na & the magnitude of this effect?
As a first approximation, ionic strength is what counts, type of halide is of secondary importance. F- is a weak base, so it will hydrolize.
I would definitely stick to pH measurements in the diluted solution.
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Yes I think you are absorbing CO2 from the atmosphere during storage.
You could try to prevent this but before you try that the question i have is does it matter for the end use?
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Borek,
Thanks for your answers. I would highly appreciate any additional scientific clarification for the above described issue.
DrCMS,
Thanks for your answer. Regarding the importance of the pH value: We don't use this product in our plant, but we are trying to keep stick to the required value of pH as stated in the specification sheet, which is =7, to avoid any customer complaints in future.
Borek / DrCMS,
I have additional question:
If the reduction in pH value can be related to the absorption of CO2: What is the expected minimum value of pH that this solution can reach? Why?
I'm asking this question because we noticed that the pH didn't go lower than 6.4~6.5 !!
I would highly appreciate any additional scientific clarification for the above described issue. Thanks a lot for your efforts.
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You probably adjusted pH to 7 with a small amount of NaOH. When you absorb some CO2 then a byffer of NaHCO3/H2CO3 can be formed
with pH close to 6.4
printing error correction
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Regarding the importance of the pH value: We don't use this product in our plant, but we are trying to keep stick to the required value of pH as stated in the specification sheet, which is =7, to avoid any customer complaints in future.
I'm sure the specification must have a range for pH ie 7+/-1 or a minimum or maximum. No pH specification should ever be set with a single figure as experimental errors/varriations will always lead to a range of reported values for the same material.
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DrCMS,
You are right. I mentioned you the lower end of the range, which is 7. The range of pH in the specification sheet is 7 to 8.5 .
AWK,
As I mentioned above, each time we found the pH lower than 7 we adjust it to 7 by adding NaOH. It is very interesting to hear about the possibility of forming a buffer solution, & I have the following comments/questions regarding that:
1-Can the buffer form in the storage tank due to CO2 absorption even without adding the NaOH solution? Or there must be some excess NaOH in the NaBr solution to initiate the buffer forming?
[I'm afraid this excess NaOH could be received during the NaBr batch endpointing. As I stated above,we produce the NaBr solution with pH ~ 7, but after some time of storage it gets down to about 6.5 (without adding the NaOH solution) & as a result we adjust it to 7 by adding NaOH].
2-Regarding the quantities: What quantities of formed buffer (for example, in % of total volume/or/ mass NaBr solution) that would create such effect of keeping the pH around ~ 6.4 ? [ I beleive that the storage size of NaBr solution may play role in estimating the needed time to have a decline in the pH value to about 6.5 after each adjustment to 7].
3- If installing CO2 filters at the tank openings to atmosphere would that help avoiding this phenomena by stopping the forming of this buffer?
Thanks for your assistance.
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By the very definition of buffer solution answers to your questions are:
1-Can the buffer form in the storage tank due to CO2 absorption even without adding the NaOH solution?
No.
Or there must be some excess NaOH in the NaBr solution to initiate the buffer forming?
Yes.
Regarding the quantities: What quantities of formed buffer (for example, in % of total volume/or/ mass NaBr solution) that would create such effect of keeping the pH around ~ 6.4 ?
Concentration of buffer present is irrelevant. See http://www.chembuddy.com/?left=pH-calculation&right=pH-buffers-henderson-hasselbalch
If installing CO2 filters at the tank openings to atmosphere would that help avoiding this phenomena by stopping the forming of this buffer?
That's what we all expect.
What are your storage tanks made off?
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Borek,
Thanks for your answers.
The Material Of Construction of our tanks is either Stainless Steel or epoxy painted Stainless Steel.
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I can think of a few different answers to the problem.
1) Can you change the spec to 6-8 for example.
2) Pack the product off to sealed containers. If you ship in bulk to the oil industry this one is out unless you make to order.
3) Add a different buffer, bicarbonate/carbonic acid buffers blood to pH~7.4 so maybe add sodium bicarbonate to the solution.
4) Prevent the CO2 from entering the storage tank.
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DrCMS ,
Thanks for your answers.
I'm also thinking about implementing Option (4).
Regarding the other ones: I must discuss these options with our team, but in general we try to avoid any actions that may seriously affect the product specifications or working plans.
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If you are adding NaOH you are most likely already creating carbonate buffer. If you add the buffer in the minute quantities at the very beginning you may have things under control from the start.
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Borek,
I'm not sure that I have got your point... Would you clarify please? Thanks
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in general we try to avoid any actions that may seriously affect the product specifications or working plans.
If you are adding NaOH to keep pH at 7, and your solution becomes contaminated with carbon dioxide, it contains carbonate buffer (CO2/HCO3-). In effect it doesn't matter (much) if you add bicarbonate to prepare buffer by yourself, or NaOH to keep pH at requested level - in both cases final effect (final composition of the solution) is the same.
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See lease the Conclusions from the above answers & my questions based on them:
Conclusions:
1- Due to CO2 abosrbtion & available some amount of NaOH: buffer NaHCO3/H2CO3 can be formed with pH close to 6.4 (not desired for our case).
2- If adding bicarbonate, at the beginning when sending the batch to storage : buffer of Na2CO3/H2CO3 can be formed with pH ~ 7.4 (desired for our case).
Questions:
1-When adding Na2CO3: Would the formed H2CO3 convert some of this bicarbonate to become NaHCO3?
2- What would be the pH value after adding Na2CO3? This question is because, as I know, at pH<9-10 you do not have carbonate ion, but rather bicarbonate ion (HCO3-) !!
3-If the facts in questions no.1 & 2 take place: What buffer would have the prevailing effect: NaHCO3/H2CO3 with pH ~6.4 or Na2CO3/H2CO3 with pH ~7.4 ?
Thanks for your response & assistance
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1-When adding Na2CO3: Would the formed H2CO3 convert some of this bicarbonate to become NaHCO3?
You mean H2CO3 formed by dissolving CO2? Yes.
2- What would be the pH value after adding Na2CO3? This question is because, as I know, at pH<9-10 you do not have carbonate ion, but rather bicarbonate ion (HCO3-) !!
It will depend on the concentration, you may use my BATE (pH calcualtor, see my signature) to calculate by yourself.
3-If the facts in questions no.1 & 2 take place: What buffer would have the prevailing effect: NaHCO3/H2CO3 with pH ~6.4 or Na2CO3/H2CO3 with pH ~7.4 ?
No such thing as CO32-/H2CO3 buffer. It is either CO32-/HCO3- or HCO3-/H2CO3. Which pair will be at work depends on the initial concentrations of carbonates and amount of CO2 dissolved.