Boric acid borate solutions

The composition of aqueous boric acid-borate solutions depends upon pH, the concentration of the total borate, and the nature of the background electrolyte if one occurs (21, 22, 23, 28, 29,30). The simplest equilibrium may be represented by the reaction of boric acid with OH" to form the borate anion. 

SvERRE Stene has measured the pH of a number of phosphate buffer mixtures, biphthalate solutions, and borate buffers with the hydrogen electrode at 150 . He found that the pH of biphthalate-hydrochloric acid solutions at 150° was about 0.2 unit greater than at 20°, the pH of biphthalate-sodium hydroxide mixtures was 0.7 greater than at 20°, while that of boric acid-borate buffers diminished with increasing temperature. Solutions of the latter system with pH s up to 9.0 were 0.5 unit less at 150°, 0.6 unit less for pH 9.2, 0.8 unit less for 9.6, and a whole unit for pH 10.0. Because certain assumptions introduced in his calculations were not entirely justified, these data must be accepted with reserve. Thus the boric acid-borate solutions behave differently from other buffers consisting of a weak acid and one of its salts. Walbum (table, page 250) also has found this diminution of pH with temperature. 

At room temperature, the pH of a sodium borate solution is 9.3. The buffering capacity of boric acid-borate solutions was evaluated at 350°C and 3500 psia by measuring the pH change during titrations with sodium hydroxide and hydrochloric acid[13]. The borate buffer titration experiments were conducted at 350°C and 24.1 MPa (3500 psia) where the water density is 0.622 g/mL and the pKa of boric acid is approximately 9.6. The concentration of sodium borate in the feed solutions was fixed at 6.25 X 10 m, which corresponded to a total boron concentration of 0.025 m. At 350°C, the measurable pH ranges of the various optical indicators are 9.0-11.0 for 2-naphthol, 7.0-9.0 for 2-naphthoic acid, 3.0-5.0 for collidine, and 2.0-4.0 for acridine. The indicator 2-naphthol was used to measure pH for the titration of the borate buffer with NaOH which began at pH of 9.5 and ended at 10.5. For the titration with HCl, 2-naphthol, 2-naphthoic acid, and collidine were used in series to measure pH values between 9.5 and 3.0. 

When a pH meter was standardized with a boric acid-borate buffer with a pH of 9.40, the cell emf was +0.060 V. When the buffer was replaced with a solution of unknown hydronium ion concentration, the cell emf was +0.22 V. What is the pH of the solution ... 

Fig. 2. Population distributions of borate species as a fraction of total boron for a 0.4 mol boric acid equivalent solution at 25 °C...

With these assumptions, the distribution of boron species (among boric acid, borate anion, borate esters, and borate diesters) reduces to a function of boron concentration, polymer concentration, radius of gyration and solvated volume fraction of a polymer unit, solution pH, and the association constants for the borate esters of the functional groups employed. 

Figure 6.3 Schematic of anodic polarization curve of iron,10 showing active-passive behavior of iron in sodium borate-boric acid buffer solution at pH 8.4...

It is often preferable for physiological purposes to use buffer systems other than the boric acid-borate mixtures. L. Michaelis has found that mixtures of veronal (diethylbarbituric acid) with its sodium salt show a satisfactory buffer action in the neighborhood of pH = 8.0. Pure commercial samples of the sodium salt of veronal are readily available, and may be used frequently without previous recrystallization. Buffer mixtures can be prepared by adding hydrochloric acid to the salt. This compound is water-free, and should suffer no loss in weight when dried at 100°. A 0.1 N solution in water requires exactly an equivalent quantity of 0.1 N hydrochloric acid when neutralized against methyl red. A stock solution should contain 10.30 g. of the sodium salt per 500 c.c. Only carbon dioxide free water should be used. 

For type 4a the three-phase line LGSn runs through a pressure maximum between the triple point of the pure component II and the quadrupole point Q where a liquid and a gaseous phase coexist with the pure solid components I and II. This curve represents the vapour pressure curve above a solution saturated with solid component II it can be determined quite easily by experiment in a closed autoclave filled with component I and an excess of solid component II and has been found for aqueous solutions of NaCl, boric acid, borates, borax, etc. ... 

Boron trioxide is not particularly soluble in water but it slowly dissolves to form both dioxo(HB02)(meta) and trioxo(H3B03) (ortho) boric acids. It is a dimorphous oxide and exists as either a glassy or a crystalline solid. Boron trioxide is an acidic oxide and combines with metal oxides and hydroxides to form borates, some of which have characteristic colours—a fact utilised in analysis as the "borax bead test , cf alumina p. 150. Boric acid. H3BO3. properly called trioxoboric acid, may be prepared by adding excess hydrochloric or sulphuric acid to a hot saturated solution of borax, sodium heptaoxotetraborate, Na2B407, when the only moderately soluble boric acid separates as white flaky crystals on cooling. Boric acid is a very weak monobasic acid it is, in fact, a Lewis acid since its acidity is due to an initial acceptance of a lone pair of electrons from water rather than direct proton donation as in the case of Lowry-Bronsted acids, i.e. 

Lead borate moaohydrate [14720-53-7] (lead metaborate), Pb(B02)2 H20, mol wt 310.82, d = 5.6g/cm (anhydrous) is a white crystalline powder. The metaborate loses water of crystallization at 160°C and melts at 500°C. It is iasoluble ia water and alkaHes, but readily soluble ia nitric and hot acetic acid. Lead metaborate may be produced by a fusion of boric acid with lead carbonate or litharge. It also may be formed as a precipitate when a concentrated solution of lead nitrate is mixed with an excess of borax. The oxides of lead and boron are miscible and form clear lead-borate glasses in the range of 21 to 73 mol % PbO. 

In general, hydrated borates of heavy metals ate prepared by mixing aqueous solutions or suspensions of the metal oxides, sulfates, or halides and boric acid or alkali metal borates such as borax. The precipitates formed from basic solutions are often sparingly-soluble amorphous soHds having variable compositions. Crystalline products are generally obtained from slightly acidic solutions. 

Determination of borate as nitron tetrafluoroborate Discussion. Boric acid (100-250 mg) in aqueous solution may be determined by conversion into tetrafluoroboric acid and precipitation of the latter with a large excess of nitron [see Section 11.11(E)] as nitron tetrafluoroborate, which is weighed after drying at 110°C. The accuracy is about 1 per cent. 

From this material, samples are cut and swelled to constant weight in a buffered saline solution prepared from 8.43 g sodium chloride (NaCl), 9.26 g boric acid (H3BO3), 1.0 g sodium borate (Na3B03), and 0.1 g of the disodium salt of the dihydrate of ethylenediaminetetraacetic acid [Na2 EDTA -(/ 0)21 ini L of distilled water. 

The general procedure for the preparation of vanadium borates consists in heating a concentrated H2O solution of boric acid and vanadium oxide in an autoclave at 170 °C for several days [143]. Two different vanadium borate clusters 105 and 106 are obtained, one with two polyborate chains coordinated to a contorted vanadium oxide ring (105) and another one with a macrocyclic Bi8036(0H)6 ring (106). The latter ring is composed of six B306(0H) units and has a chair-like conformation (Fig. 27) [143]. 

Delayed Crosslinking Additives. Glyoxal [458,460,461] is effective as a delay additive within a certain pH range. It bonds chemically with both boric acid and the borate ions to limit the number of borate ions initially available in solution for subsequent crosslinking of a hydratable polysaccharide (e.g., galactomannan). The subsequent rate of crosslinking of the polysaccharide can be controlled by adjusting the pH of the solution. 

Anions of weak acids can be problematic for detection in suppressed IEC because weak ionization results in low conductivity and poor sensitivity. Converting such acids back to the sodium salt form may overcome this limitation. Caliamanis et al. have described the use of a second micromembrane suppressor to do this, and have applied the approach to the boric acid/sodium borate system, using sodium salt solutions of EDTA.88 Varying the pH and EDTA concentration allowed optimal detection. Another approach for analysis of weak acids is indirect suppressed conductivity IEC, which chemically separates high- and low-conductance analytes. This technique has potential for detection of weak mono- and dianions as well as amino acids.89 As an alternative to conductivity detection, ultraviolet and fluorescence derivatization reagents have been explored 90 this approach offers a means of enhancing sensitivity (typically into the low femtomoles range) as well as selectivity. 

Buffers contain mixtures of weak acids and their salts (i.e., the conjugate bases of acids), or mixtures of weak bases and their conjugate acids. Typical buffer systems used in pharmaceutical dosage forms include mixtures of boric acid and sodium borate, acetic acid and sodium acetate, and sodium acid phosphate and disodium phosphate. The reason for the buffering action of a weak acid, HA (e.g., acetic acid) and its ionized salt, A" (e.g., sodium acetate) is that A" ions from the salt combine with the added hydrogen ions, removing them from solution as undissociated weak acid. 

What molar ratio of acid to salt is required to adjust the pH of a solution to 8.8 using boric acid-sodium borate buffer ... 

What molar ratio of salt/acid (sodium borate to boric acid) is required to prepare a buffer solution having a pH of 9.6 The pKa of boric acid is 9.24 at 25°C. 

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