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My textbook says:

Boric acid is a weak monobasic acid but on the addition of certain organic polyhydroxy compounds, such as mannitol, glycerol, dextrose or invert sugar, it is transformed into a relatively stronger acid. Ethylene glycol cannot give this test (the reason is uncertain).

I do know that cis-diols form a boron-ether compound with the $\ce{[B(OH)4]-}$, leading to an increase in acidity. But is it true that ethylene glycol is exceptional? And is there any possible reason for it?

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  • $\begingroup$ Maybe it is about the molecular geometries. I can't run the necessary calculations until at least the weekend, though. Would be interesting to know how buta-2,3-diol would behave (but then again, maybe it does not dissolve in water). $\endgroup$
    – TAR86
    Mar 20 '19 at 15:17
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    $\begingroup$ But a question in iit jee advanced(2014) says that on addition of ethylene glycol too the acidity of orthoboric acid is increased. ques- The correct statement(s) for orthoboric acid is/are (A) It behaves as a weak acid in water due to self ionization. (B) Acidity of its aqueous solution increases upon addition of ethylene glycol. (C) It has a three dimensional structure due to hydrogen bonding. (D) It is a weak electrolyte in water. ans- B,D (official) i am also confused now beacuse its opposite is written in j d lee as per your question $\endgroup$ Apr 20 '19 at 14:36
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    $\begingroup$ Yes it has been written in jd Lee that ethylene glycol don't give this test $\endgroup$
    – Kuvcharu
    May 9 at 1:42
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    $\begingroup$ Will somebody please make a solution of boric acid, check its pH, then add some ethylene glycol to see if the pH decreases, and by how much? It would save time for so many people checking their textbooks. $\endgroup$ May 9 at 13:54
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    $\begingroup$ I think all are reading the Indian version of JD Lee the Sudarshan Guha one the original does have the ethylene glycol debate $\endgroup$
    – lakshman
    Jun 8 at 16:23
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Boric acid is a very weak acid with a $\mathrm{p}K_\mathrm{a}$ value of 9.2 and does not dissociate in aqueous solution as a Brønsted acid. Yet, in aqueous solutions, it acts as a Lewis acid by accepting a hydroxyl ion to form the tetrahydroxyborate ion $(\ce{B(OH)4-})$, as confirmed by Raman spectroscopy (Ref.1):

$$\ce{B(OH)3 + 2H2O <=> B(OH)4- + H3O+} \tag1$$

Thus, the dominant forms of inorganic boron in natural aqueous systems are mononuclear species such as boric acid $(\ce{B(OH)3})$ and borate ion $(\ce{B(OH)4-})$. The distribution of these two components depends on the first dissociation constant $K_\mathrm{a}$ of boric acid, which is equal to $\pu{5.80 \times 10^{−10} mol L-1}$ in fresh water at a temperature of $\pu{25 ^\circ C}$ (Ref.2).

Note that in solutions that are more concentrated than $\pu{0.1 M}$, boric acid acts as a much stronger acid than in diluted solutions, and becomes comparable to acetic acid. As shown in the equation $(1)$, the apparent ratio of the concentration of borate ions to that of boric acid molecules in the solution progressively increases from $1:1000$ at $\pu{0.2 M}$ to $1:5$ at $\pu{3.5 M}$ (Ref.1).

In an aqueous environment, boric acid and borates react with alcohols and multiple hydroxyl-containing compounds (polyols) forming boron esters. For example:

$$\ce{B(OH)3 + 3 R-OH <=> B(OR)3 + 3 H2O} \tag2$$

Borate ester complexes

As shown in the above figure, complexation with polyols increases the acidity of boric acid due to the formation of cyclic borate esters. For example, the dissociation constant of boric acid becomes about $7 \times 10^{−6}$ in the presence of mannitol, i.e. a $10$- to $1000$-fold increase compared to that without the polyol (Ref.1). Such behavior is not shown on the addition of monofunctional alcohols, nor of glycols, nor of the trans-form of cyclopentane diols, but it is shown upon the addition of the cis-form of the latter compounds. AS described in the Ref.1:

The phenomenon depends on the formation of monocyclic or dicyclic compounds with the polyol groups, which are more highly dissociated than the boric acid itself, and it follows that the only compounds with two hydroxyl groups suitably placed on the same side of the C–C link can react in this way. Thus, the stability of the borate complex formed is strongly dependent on the type of diol used. If the diol involves the OH groups oriented in such a way that they accurately match the structural parameters required by a tetrahedrally coordinated boron, a strong complex is formed. For instance, erythritan (3,4-cis-dihydroxytetrahydrofuran) has produced a much lower $\mathrm{p}K_\mathrm{a} \ (\approx 4.8)$ with boric acid (more acidity, stronger complex; Ref.3).

Ester formation

The equilibrium constants of borate complexes have been investigated in several studies, and some of the data are listed in the following Table (Ref.3):

$$\begin{array}{l|c|r} \hline \text{Polyol} & K_1 & K_2\\ \hline \text{1,2-Ethanediol} & 2.15 & 1.15\\ \text{1,3-Propanediol} & 1.27 & 0.11\\ \text{Glycerol} & 16.0 & 41.2\\ \text{Catechol} & 7.8 \times 10^3 & 1.42 \times 10^4\\ \text{D-Mannitol} & 1.1 \times 10^2 & 1.37 \times 10^5\\ \text{D-Glucose} & 1.50 \times 10^3 & 7.60 \times 10^2\\ \hline \end{array} $$

It has been shown that the amount of acidification produced upon the addition of polyol is proportional to the extent of ester formation (Ref.3). Basically, 1,2-Ethanediol is ethylene glycol and it also has a $K_1$ and $K_2$ values as well although they are not pronounce as those of D-Mannitol or even Catechol. Therefore, the increase in acidity may not significant, but it sure does the trick. The lower values of $K_1$ and $K_2$ are probably due to restrictions in structural parameters required by a tetrahedrally coordinated boron.

References:

  1. Victor Kochkodan, Nawaf Bin Darwish, and Nidal Hilal, "Chapter 2: The Chemistry of Boron in Water," In Boron Separation Processes; Nalan Kabay, Marek Bryjak, and Nidal Hilal, Editors; Elsevier B.V.: Amsterdam, The Netherlands, 2015, Pages 35-63 (DOI: https://doi.org/10.1016/B978-0-444-63454-2.00002-2)(ISBN of the Book: 978-0-444-63454-2).
  2. Benton Brooks Owen, "The Dissociation Constant of Boric Acid from 10 to 50°," J. Am. Chem. Soc. 1934, 56(8), 1695–1697 (DOI: https://doi.org/10.1021/ja01323a014).
  3. Philip P. Power and William G. Woods, "The chemistry of boron and its speciation in plants," Plant and Soil 1997, 193, 1–13 (DOI: https://doi.org/10.1023/A:1004231922434).
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  • $\begingroup$ In the figure, a particular polyol, reacts with both: boric acid and borate ions to increase the acidity. Can we say that the value of $K_1$ and $K_2$ (boric acid) should be less than $K_3$ and $K_4$ (borate ion) to increase the acidity more? Thus, the polyol preferably reacts to $\ce{[B(OH)4]-}$ to remove it and the equilibrium of reaction 1 also shifts in forward direction. $\endgroup$
    – Apurvium
    Jul 21 at 2:10
  • $\begingroup$ Does the vale of $K_1$ and $K_2$ in the table below that figure represents the same $K_1$ and $K_2$? And I think, in the balanced reaction for $K_2$, there will be $\ce{1H2O}$ and that for $K_4$, there won't be any $\ce{H+}$. $\endgroup$
    – Apurvium
    Jul 21 at 2:15
  • $\begingroup$ @ Apurvium 20: The reference did not report $K_3$ and $K_4$ so I won't able to discuss the difference. $\endgroup$ Jul 21 at 23:06
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I do not see any reason why the following reaction cannot take place, which will consume $\ce{[B(OH)4]-}$ and the concerned reaction will move in forward direction to increase the acidity of orthoboric acid$^1$.

enter image description here

According to wikipedia, "With polyols containing cis-vicinal diols, such as glycerol and mannitol, the acidity of the boric acid solution is increased."

If someone has aceess to this or this journal, then please add.


Reference

  1. Page 375-376, Concise Inorganic Chemistry, 5th Edition, J. D. Lee, Wiley (ISBN 978-81-265-1554-7)
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    $\begingroup$ I think JD Lee writes "In effect, it acts as a strong acid in the presence of the cis-diol except ethylene glycol" which might be the source of confusion for others as it was for me as well. It probably means that all cis-diols enhance the acidity to an extent so as to make it a strong acid except ethylene glycol which does enhance it but not to the extent of others. $\endgroup$
    – Ashish
    Jul 17 at 4:21
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    $\begingroup$ I do not see any reason either. The question is whether ethylene glycol is exceptional. Or, rather HOW exceptional. I predict that free-rotating ethylene glycol, like the structurally tight cis-diols, will increase the acidity of boric acid, but probably not as much as the sugars and other tight cis-diols. So if you do the acidity test with paper, you will see no definite pH drop; however, if you do the test with a two-place pH meter, you will see a small drop in pH. Can I offer a bounty for someone to do this test and report the actual numbers so we can all discuss it? $\endgroup$ Jul 17 at 14:23
  • $\begingroup$ @Apurvium The reaction looks similar on paper to the formation of acetals in organic chemistry. These are equilibrium reactions, synthetically useful (i.e., at good yield within reasonable time) if run under acid catalysis, elevated temperature, and active removal of water (e.g., Dean-Stark apparatus). Does your source details out such conditions in addition to an balanced reaction equation? $\endgroup$
    – Buttonwood
    Jul 18 at 18:35
  • $\begingroup$ @Buttonwood No, my reference book does not provide more details but I added few more sources to my answer. $\endgroup$
    – Apurvium
    Jul 19 at 3:08

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