Chemical Forums
Chemistry Forums for Students => Organic Chemistry Forum => Topic started by: Lucek on April 23, 2020, 02:01:45 AM
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Hi everyone,
when considering naphthalene sulfonate solution (i.e. 1-NSA in water), lets say that everything is happening at room temperature and the water solution has also Na+, Cl- ions and pH = 5.5, we have excess of Na+, Cl- and H+ in comparison to 1-NSA.
What should I look at (bonding energy? pKa?) to establish affinity of the SO3- to Na+ or H+ ? What should I consider to understand what is more likely to happen (SO3-Na+ or SO3-H+)? and what is the proper way to establish which bond will be more stable in the solution if appear?
Cheers,
L
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First of all, the concentrations of all components of the solution are needed for any calculations. You have only provided one value - pH of your solution. In addition, the Ka of naphthalene sulfonic acid is so high that the low acid concentration in the solution and the even lower concentration of sodium naphthalene sulfonate have practically no significant effect on the equilibria in solution.
In more accurate physicochemical calculations, at most, the influence of ionic strength on slight changes in the equilibria would be taken into account.
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The concentrations ratio [ArSO3Na]/[ArSO3H] can be estimated from the Henderson-Hasselbalch equation:
pH = pka + log([ArSO3Na]/[ArSO3H])
and [ArSO3Na] + [ArSO3H] = initial [ArSO3H]
where,
1-naphthalenesulfonic acid pka = 0.17
2-naphthalenesulfonic acid pka = -1.80
Substituted naphthalenesulfonic acids pka, depends on their substitution.
But as AWK has stated above, such calculations are not very accurate for strong acids (say, pka < 1). In more accurate calculations, contribution of water ionization (pkw =14) must also be taken into the account.
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The concentrations ratio [ArSO3Na]/[ArSO3H] can be estimated from the Henderson-Hasselbalch equation:
pH = pka + log([ArSO3Na]/[ArSO3H])
and [ArSO3Na] + [ArSO3H] = initial [ArSO3H]
Not true for pH=5.5
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Whether true or not, the pH adjustment at 5.5 in shampoos and skin care products, is estimated that way.
Dodecylbenzenesulfonic acid pka = 0.7
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Shampoos are buffered by citrate buffer.
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At pH 5.5 and pKa 0.7 ratio base/acid is 105.5-0.7≈63,000 - definitely possible to use it to estimate concentrations from this number. Presence of other buffers doesn't matter much - the only thing that matters is that the pH is established and won't change.
Not that it is going to help answer the original question in any way. Ka can be found in tables, but we are still left with estimating stability constant for Ar_SO3-Na+ (doesn't matter whether we call it a complex, ion pair, or something else).
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Theoretically, such a production is effectuated by continuous monitoring the pH. But in practice, the necessary amount of the base in a 1 ton batch containing e.g. 250 kg of an organosulfonic or organophosphonic acid is pre-estimated that way, followed by pH monitoring after the > 95% of base being added.
Furthermore, addition of a little amount of a buffer to accurately stabilize the already adjusted pH, is useful but not usual in low cost-products mainly due to the defoaming properties of buffers.
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First of all, the concentrations of all components of the solution are needed for any calculations. You have only provided one value - pH of your solution. In addition, the Ka of naphthalene sulfonic acid is so high that the low acid concentration in the solution and the even lower concentration of sodium naphthalene sulfonate have practically no significant effect on the equilibria in solution.
In more accurate physicochemical calculations, at most, the influence of ionic strength on slight changes in the equilibria would be taken into account.
Let's assume that the Naphthalene Sulfonate concentration is extremely low in comparison to Na+ Cl- and of course H+.
Napthalene sulfonate ~ 8·10-8 molal
whereas Na+ Cl- 0.05 molal
will the difference in size of Na+ and H- metter?
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https://www.chembuddy.com/?left=pH-calculation&right=pH-salt-solution
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https://www.chembuddy.com/?left=pH-calculation&right=pH-salt-solution
Have it occurred to you that this is not a pH calculation question?
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@ OP, I don't see how at pH 5.5 that the naphthalene sulfonate can be low in concentration compared to protons.
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@ OP, I don't see how at pH 5.5 that the naphthalene sulfonate can be low in concentration compared to protons.
Please elaborate - pH changes the ratio between protonated and non-protonated forms, but in no way says anything about total concentration.
Plus, as far as I can tell, none of the OP posts suggested pH is a result of the presence of the sulfonate, all we know is that it is 5.5, for whatever reason (it can be buffered separately).
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Borek,
I think I see what you mean, but it would be a very dilute solution of the naphthalene derivative if that were the case, around a micromolar.
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Let’s try to clarify the issue and help the OP who probably started to be confused.
1). NaCl is a salt of a strong acid (HCl) and a strong base (NaOH). Furthermore, naphthalene sulfonic acids are strong acids, too. Consequently, there are (almost) only ions in their solutions; meaning that Na+ cations do not exclusively belong neither to chloride anions, nor to naphthalene sulfonate anions. All ions are in equilibria therein.
2). On the other hand, the concentrations ratio of acid salt and free acid depends on the pka and the pH only, regardless the presence of a buffer or not and where the cation comes from (common ion effect, among others).
3). But as mentioned above, Henderson-Hasselbach equation does not work well for concentrated solutions and strong acids. In these cases, more complex calculations are necessary.
4). However, the error of the Henderson-Hasselbach equation is considered negligible for dilute solutions of strong acids and their salts.
5). Besides and apart a few exceptions, the exact values for concentrated solutions of strong acids with their salts have poor practical value, as measures of comparison between acidic solutions.
6). Indicatively, an electrode pH-meter shows pH = -1 for a solution of NaCl in 6 N HCl; which is wrong (pH < 0 ) but everybody respects it, as being a comparison measure of how acidic this solution is.
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an electrode pH-meter shows pH = -1 for a solution of NaCl in 6 N HCl; which is wrong (pH < 0 )
Wrong why?
And from the very beginning the question was not about pH, but about Na+/sulfonate association. More or less I read the initial post as "is it possible to estimate stability constant for Sulfonate/Na+ knowing stability constant for Sulfonate/H+?". The latter is just Ka-1.
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Because pH = -1 is translated to [H+] = 10 N for a solution of 6 N HCl; or else, the protons concentration is higher than the concentration of the monovalent acid that is not true.
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pH is not -log of concentration but of activity. For 0.01 M HCl measured pH is already not 2, but around 2.05.
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@ OP, I don't see how at pH 5.5 that the naphthalene sulfonate can be low in concentration compared to protons.
pH was controlled by use of HCl
for some reason when only HCl is present than Naphthalene Sulfonate is protonated and we results in naphthalene + SO3 , whereas when I add NaCl to the mixture the Naphthalene Sulfonate remains stable. I am trying to understand the mechanism behind this phenomena. I can image few options... but it might be unrealistic:
1) the Na+ is attracted to Naphthalene Sulfonate (to the sulfonate itself) and this attraction protect the naphthalene backbone from the H+ attack which would appear if no NaCl
2) maybe the ionic strength controls the stability? as H+ activity decreased by increasing salt concentration
3) Solvation ?
I calculate that pH wont change much even if we heat the solution up to 300C
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for some reason when only HCl is present than Naphthalene Sulfonate is protonated and we results in naphthalene + SO3
You mean it decomposes? How do you determine that?
Looks like you asked about secondary things completely derailing the thread from the very beginning :(
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Correct!
Thus:
An electrode pH-meter shows pH = -1 for a solution of NaCl in 6 N HCl; which is wrong (pH < 0 ) because pH = -1 is translated to protons activity = 10 N for a solution of 6 N HCl; or else, the protons activity is higher than the concentration of the monovalent acid that is not true.
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Naphthalenesufonic acids and HCl are strong acids and therefore, their salts are in equilibrium:
ArSO3Na + HCl ← → ArSO3H + NaCl
Thus:
1). Addition of NaCl forces the equilibrium to the left.
2). Heat forces the equilibrium to the right because HCl is a gas.
Conclusion: Better results can be obtained by slow addition of aqueous HCl in presence of NaCl at low temperatures, e.g. cold water bath that may contains a few pieces of ice.
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Correct!
Thus:
An electrode pH-meter shows pH = -1 for a solution of NaCl in 6 N HCl; which is wrong (pH < 0 ) because pH = -1 is translated to protons activity = 10 N for a solution of 6 N HCl; or else, the protons activity is higher than the concentration of the monovalent acid that is not true.
Sorry to say that, but it looks like you have no idea what you are talking about. Activity coefficients for high ionic strength solutions are higher than 1, so it is perfectly OK for activity of H+ to be higher than the concentration in 6N solution. See for example http://www.umich.edu/~chem241/lecture11final.pdf
2). Heat forces the equilibrium to the right because HCl is a gas.
You would need very high concentrations of HCl for that to matter, we are talking pH 5.5 solution. Stop posting nonsense, you are not helping.
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Dear Administrator,
Thank you for your kind compliments!
But please, permit a few remarks:
1). In the provided reference, it is clearly stated that …”So if solubility increases with ionic strength---meaning that concentrations increase---then activity coefficients decrease as you increase ionic strength!!”……
2). In the provided reference, it is also stated that …”Debye-Huckel equation valid from μ =0 --> 0.1 M; beyond, not very accurate at predicting activity coefficient!”
3). Moreover, it is stated therein that …” the smaller the hydrated radius--more effect of μ on activity coeff. (decrease)”…..
4). The cited schema refers to the increase of NaClO4 concentration and change of activity coefficient; but not to the increase of acid (HClO4) concentration, neither to the proton activity coefficient. Besides, ClO4- is a very bulky anion and consequently, it might not be a typical example for general conclusions.
5). By ending, HCl solubilization in water is highly exothermic and thus, fast addition of small amounts of concentrated HCl can significantly increase the HCl partial pressure and significantly decrease its Henry's constant, even at room temperature.
Sincerely yours
PS: Thinking twice, stopping posting does not sound a bad idea.