So I have some pH = {4, 7, 10} standard buffer solutions from Fisher Scientific. But they must have went bad or I must have contaminated them because when I try to calibrate my pH meter to 4 using the pH 4 standard buffer solution, it calibrates the pH meter to 2 which obviously isn't right because it should be 4 so I can't trust the standard buffer solutions!!!
I need to do some experiments which require that I accurately calibrate the pH meter which means I can't titrate the standard buffers with HCl or NaOH. I have to know the exact quantities. Might seem a little over-the-top, but I want to be really certain about the math.
I intend to prepare my own calibration standards using exact quantities of chemicals. Total concentration should be 0.1 M, volume needs to be 100 mL. I think 0.1 M is a good choice because that should have low ionic strength and not much crazy debye-huckel variations in pH.
I am pretty certain that I know how to prepare a pH = 4 standard buffer solution using acetic acid and sodium acetate. But I am not quite sure how to prepare a pH 7 standard.
First of all, is this exactly how I should prepare a 0.1 M acetate buffer in 100 mL?
According to pubchem the pKa of acetic acid is 4.76. I have a chem textbook which tells me that the pKa is 4.74. Not much difference so I will assume pubchem is right.
HA + H2O --> A- + H3O+
Acetic Acid + H2O --> Acetate + H3O+
Henderson-Hasselbalch Equation: pH = pKa + log([A-] / [HA])
pH = pKa + log([Acetate] / [Acetic Acid])
= 4.76 + log([Acetate] / [Acetic Acid])
For the total concentration:
[Acetic Acid] + [Acetate] = 0.1 M
The simultaneous solution to these equations is about:
[Acetic Acid] = 0.0852 M
[Acetate] = 0.0148 M
pH = 4.76 + log(0.0148 / 0.0852) = 3.9998 etc ... Pretty much exactly 4!
I have glacial acetic acid and sodium acetate.
Glacial Acetic acid is a pure liquid so its concentration is its density / (Formula Weight)
[Glacial Acetic Acid] = [(1.05 g/mL) / (60.05 g/mole)] * (1000 mL / 1 L) = 17.5 M
Basic Dilution Equation:
Initial Volume = (0.0852 M / 17.5 M) * 100 mL = 0.487 mL glacial acetic acid in 100 mL total volume.
Sodium Acetate --> Na+ (aq) + Acetate- (aq)
grams sodium acetate = (0.0148 moles sodium acetate / L) * (82.03 grams sodium acetate / 1 mole sodium acetate) * (1L / 1000 mL) * 100 mL = 0.121 grams sodium acetate
So I can exactly prepare a pH = 4 is 0.1 M acetate solution by adding 487 microliters of glacial acetic acid & 0.121 grams sodium acetate to 100 mL volumetric flask, and bring the volume up to 100 mL with ultrapure deionized water? Swirl to mix.
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I am confused about how to prepare an exactly pH = 7 phosphate buffer solution.
I know the step wise dissociation of Phosphoric Acid. But the problem is that I am not certain about the second dissociation constant.
H3PO4 --> H2PO4- (aq) + H+ (aq)
H2PO4- (aq) --> HPO42- (aq) + H+ (aq)
If I know the second dissociation constant of phosphoric acid I could figure out exactly how to make a pH = 7 standard 0.1 M phosphate buffer solution using phosphate salts.
My undergrad general chem text book is telling me the second Ka is 6.2 * 10^(-8) or the pKa is 7.208 and pubchem is telling me the second the second pKa is 7.09.
Thus I could use the Henderson-Hasselbalch Equation again with the total concentration to figure out the exact quantities. But I am not certain about what the second pKa of phosphoric acid is! Are undergrad chem textbooks just wrong? That's too much of a margin of error.
pH = pKa + log[(K2HPO4) / (KH2PO4)]
Also I was wondering if a 0.1 M solution of NaCl would have a pH of EXACTLY 7? Would it be more or less accurate than making the phosphate buffer?
NaCl is formed by the titration of a strong acid HCl with a strong base NaOH. Therefore the equivalence point of the titration of a strong acid and strong base should yield a pH 7 of neutrality forming NaCl.
So if I dissolved NaCl in water to yield a 0.1 M concentration would the pH be just as good as making that phosphate buffer solution?
I also need to figure out how to make a pH = 10 standard buffer solution.
Any advice? Can you make my life easier?
Thanks so much!
There must be some problem with the pH meter then.. U should show it to a technician so that it can be fixed.
Update: There was nothing wrong with my pH standard solutions. My pH meter was just having a fun time giving me bogus pH readings of the standard pH solutions, and refusing to calibrate. I had to tinker with the pH meter for 1 and a 1/2 days to fix it.
Turns out I calculated exactly how to make my pH 4 standard solutions. But still my pH meter refused to read it correctly for some time.
There must be some problem with the pH meter then.. U should show it to a technician so that it can be fixed.
For the last part of your question (NaCl vs phosphate buffer). Of course the pH of NaCl 0.1 M is thoretically 7.00, as well as pure water has pH=7.00. The point is that NaCl is not a pH buffer: any external factor (absorption of atmospheric CO2, traces of other chemicals on your glassware, etc...) would not be buffered and the pH of your solution will change in few minutes.
is it your pH meter dose not give slope value after calibration?, normally most of (digital ) pH meter gives slope value after buffer calibration and you can get a rough idea of your instrument and std buffer solution from it.
Try calibrating with glass distilled water, if it did not work, then your pH meter has an issue.
Probably your pH 4 solution is contaminated. You can get a solution with pH 4 manually as follows: Add 20g per liter of citric acid and then add KOH slowly to increase the pH up to 4.00. After preparation leave the buffers to rest for a few hours and measure the pH again to ensure that your solution pH remains stable.
You can buy paper pH it is really usefull when you have some doubts...
And for reapeatability, when performing pH measurement always try to avoid electrode contact with your beaker/falcon
There are all sorts of pitfalls involved in making buffer solutions. Even 0.1 M is sufficiently concentrated to make corrections for activities significant. It is not a good idea to try and develop your own buffer solutions if you want accurate results. Books such as Lange's Handbook of Chemistry contain recipes that you can follow for making up standard buffers - with a standard ionic strength, too. There should be no problem with purchased buffer standards provided the contents of the sachet have been kept dry. The problem is much more likely to be due to your pH electrode. Does your pH meter respond rapidly or slowly to a change in pH? If slowly, then your electrodes are probably not in good condition and you should use fresh ones. If the glass electrode dries out, its resistance goes way up, which gives slow and inaccurate results. In my experience, attempts to revive them are usually unsuccessful. This seems to be a particular problem in a hot climate. Use the "Hendersen-Hasselbalch" equation with caution. It will only give you approximate results, and breaks down entirely in some circumstances. With an acid such as citric acid, which has pKa values which are close together, it won't work. With phosphoric acid, with widely spaced pKa values, it can be used fairly successfully, since only one pKa value need be considered at any given time.
There are a number of factors while determining the pH of unknown solutions. First is the state of the electrode. These days digital pH meters are generally good. If the electrode has dried out then pH can’t be assessed correctly. It is recommended always to keep the standard electrodes in double glass distilled water. Second is the pH of the standard solutions. Generally speaking standard solutions or sachets from well known companies behave ok provided the instructions have been followed to the point. Thirdly always use glass distilled water for making solutions. Companies also sell broad range and narrow range pH papers that can be used to roughly estimate the pH of given solution.
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