Reef Chemistry Question of the Day #53 Boron

[Randy Holmes-Farley]

Reef Chemistry Question of the Day #53

Boron is present in seawater at pH 8.2 in which of the following forms?


A. 70% Boric acid (B(OH)₃), 30% boride (B^-).
B. 70% Boric acid (B(OH)₃), 30% borate (B(OH)₄^-)
C. 30% Boric acid (B(OH)₃), 70% borate (B(OH)₄^-)
D. 30% Boric acid (B(OH)₃), 70% boride (B^-)

Good luck!













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[hatfielj]

I'm going to guess C

 

[zemuss]

I think its C

 

[DFW]

C!! The relative concentrations of borate and boric acid are dependent on pH, temperature and salinity. The borate ion becomes more prevalent at higher pH, higher salinity, and higher temperature. If it is not C, then it definitely is B!

 

[Steven@]

C...

 

[ingtar_shinowa]

interesting

 

[s.h.moradi]

I think B is correct

 

[gdwats]

I say C

 

[Randy Holmes-Farley]

Anyone else what to try before I give the answer?

 

[Randy Holmes-Farley]

And the answer is...B. 70% Boric acid (B(OH)₃), 30% borate (B(OH)₄^-)

The pKa of boric acid in seawater is about pH 8.55. That means that at pH 8.55, there are equal concentrations of boric acid and borate. At any pH value below the pKa there is more boric acid than borate. Since this question asked about pH 8.2, the boric acid predominates

This has more:

Boron in Seawater
In natural seawater, boron is present at about 0.41 mM (4.4 ppm total boron) and takes two different chemical forms. The predominate form is boric acid, comprising about 70% of the total boron present. Boric acid, B(OH)₃, consists of a central boron atom and three hydroxyl groups arranged in a triangle. The second form is borate, B(OH)₄^-. It consists of a central boron atom and four hydroxyl groups arranged in a tetrahedron. It carries a net negative charge, while boric acid is neutral. These two forms can interconvert in less than a second, so the two forms are in chemical equilibrium with each other.
Boric acid and borate have a variety of chemical and biological effects in normal seawater, though biological effects are rarely discussed in the marine aquarium hobby. The biological effects are, however, important in certain situations, and will be discussed later in this article. The most commonly known property of this system is its ability to buffer seawater against pH changes, and this aspect is discussed in detail in the next section.


Buffering of Normal Seawater
The exact percentage of boric acid and borate in any aqueous system is dependent on pH. The pKa of boric acid in seawater is about 8.55, depending on the temperature.1 That is, the pH where both forms are equally represented is about 8.55 in a normal tropical reef tank. At lower pH values, such as those in most reef tanks, there is more boric acid than borate.


The fact that the two forms are related by equation 1 explains why boric acid and borate together form a buffer system:


1) B(OH)₃ + H₂O → B(OH)₄^- + H⁺ (pKa ~8.55)


If the pH in a system containing both forms were to begin to rise for any reason, some of the B(OH)3 would be converted into B(OH)4-, releasing a proton, H+. The pH would then not rise as much as it otherwise would. Likewise, if the pH in that system were to begin to drop for any reason, some of the B(OH)4- would be converted into B(OH)3, taking up a proton. The pH would then not drop as much as it otherwise would.

his effect is exactly how a standard buffer works.


The other system that significantly buffers the pH in normal seawater is the bicarbonate/carbonate system:


2) HCO₃^- → CO₃^2- + H⁺ (pKa ~ 8.92)


The relative buffering of these two systems (equations 1 and 2) depends on the amounts of each system present, and also on the pH. For any given buffer system, it turns out that the pH at which it gives optimal buffering corresponds to the pKa, where there are equal amounts of each of the two forms of the buffer present. Why exactly this is true is beyond this article, but it relates to the fact that for a given incremental change in pH, more of the buffer will change form at that pH than would change form at any higher or lower pH, and moreover, the farther that the pH is from the pKa, the smaller the effective capacity of that buffer system to resist pH changes.


Buffering capacity can be quantified using the buffer intensity, b, defined mathematically in a way that is easy to calculate, but that isn’t worth detailing here.2 The units of the buffering intensity can be expressed as meq/L or meq/L/pH unit (these are equivalent since pH is really a dimensionless parameter). Thinking about it as meq/L/pH unit makes it easier to understand that it is a measure of the amount of alkalinity (or acidity; either one measured as meq/L) that needs to be added to impact the pH up or down by one unit (though that is a substantial simplification).


In the case of normal seawater at pH 8.2, b = 0.19 meq/L/pH unit for the boric acid/borate system, and 0.63 meq/L/pH unit for the bicarbonate/carbonate system. These values are additive, and result in a total buffering of b = 0.82 meq/L/pH unit. Under these conditions, the boric acid/borate system provides about 23% of the total buffering, while the bicarbonate/carbonate system provides about 77%.


If the pH of normal seawater is raised to 8.5, the total buffering is b = 1.2 meq/L/pH unit, or about 40% greater than at pH 8.2 (because both systems are closer to the pKa). At this pH, the relative contribution of the two systems to the total capacity is only slightly different than at pH 8.2, with 20% from borate and 80% from carbonate.

If the pH of normal seawater is lowered to 7.8, the total buffering is b = 0.42 meq/L/pH unit, or about half that at pH 8.2 (because both systems are farther from the pKa). At this pH, the relative contribution of the two systems to the total capacity is also only slightly different than at pH 8.2, with 29% from borate and 71% from carbonate.