[Lucek]

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

[AWK]

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.

[pgk]

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.

[AWK]

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

[pgk]

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

[AWK]

Shampoos are buffered by citrate buffer.

[Borek]

At pH 5.5 and pKa 0.7 ratio base/acid is 10^5.5-0.7≈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_SO_3^-Na⁺ (doesn't matter whether we call it a complex, ion pair, or something else).

[pgk]

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.

[Lucek]

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?

[AWK]

https://www.chembuddy.com/?left=pH-calculation&right=pH-salt-solution

[Borek]

https://www.chembuddy.com/?left=pH-calculation&right=pH-salt-solution

Have it occurred to you that this is not a pH calculation question?

[Babcock_Hall]

@ OP, I don't see how at pH 5.5 that the naphthalene sulfonate can be low in concentration compared to protons.

[Borek]

@ 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).

[Babcock_Hall]

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.

[pgk]

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.