The Dissociation Constant of Boric Acid from 10 to 50°

Owen, Benton Brooks

doi:10.1021/ja01323a014

Cited by 63 publications

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“…The relationship between the activity coefficients of univalent and bivalent ions is given by eq 3. The activity coefficient of a bivalent anion, A =, in the mixture is then given in terms of the mean activity coefficient of hydrochloric acid by (4) The same values of at and (3 are assumed to apply to all ions of a particular mixture at a particular temperature. The values of these parameters are considered to be weighted averages for all the types of ions in the solution.…”

Section: Discussion Of the Ele Ctromotive-force Meth Odmentioning

confidence: 99%

“…The blank correction was determined with Nile Blue A sulfate. 4 The electrometric end point was readily estimated to within 0.2 percent by graphical means, but a further uncertainty of nearly the same magnitude was introduced with the use of a blank correction. The results of the titrations of the phenol group with alkali must, therefore, be considered as only corroborative evidence of the purity of the salt.…”

Section: Methodsmentioning

confidence: 99%

“…Furthermore, such solutions may contain appreciable concentrations of associated polyboric acids [3,4]. Until boric acid solutions have been further investigated, their pH values between pH 8 and pH 9 will be subject to some uncertainty.…”

Section: Alkalineibuffer Standardsmentioning

confidence: 99%

“…The change of pH with temperature of four buffers o~ series A is shown in figure 8.9 Curves A, B, C, and D represent, respectively, solutions A8, A5, A3, and AI. 4. Curves A, B, 0, and D represent respectively solutions AS, A5, A3, andA1.…”

Section: Calculation Of the Ph Values From Electromotive-force Measurmentioning

confidence: 99%

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The second dissociation constant of p-phenolsulfonic acid and pH values of phenolsulfonate-chloride buffers from 0 degrees to 60 degrees C

Bates¹,

Siegel²,

Acree³

1943

J. RES. NATL. BUR. STAN.

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The thermodynamic dissociation constant of the phenol group of p-phenolsulfonic acid was calculated from electromotive-force measurements of hydrogensilver-chloride cells without liquid junction. Thirty-nine buffer mixtures of potassium p-phenolsulfonate, sodium hydroxide, and sodium chloride were studied in five series of experiment s from 0° to 60° C at intervals of 5 degrees. In three series, t he molal ratio of phenolsulfonate ion to phenolate-sulfonate ion was unity, and in two series the buffer ratio was 2 :3. For two series of experiments, one at each buffer ratio, the molality of sodium chloride was m aintained constant near 0.05 for all dilutions of t he buffer. In the other experiments, the molality of each component of t he solut ion varied bet ween 0.0037 and 0.1.The values of pK2' the negative of the common logarithm of the second dissociat ion constant, between 0° and 60° C are given by the equation012139T, where T is in degrees Kelvin.Equations were formulated to give the changes of free energy, heat content, entropy, and heat capacity that accompany t he dissociation, at infinite dilution, of a mole of phenolsulfonate ion at any temperature between 0° and 60° C. For the dissociation of the phenol group at 25° C, ~F o is 12,351 cal, ~Ho is 4,036 cal, ~So is -27.9 cal deg-1, and ~c o p is -33 cal deg-1•The pH value of each buffer-chloride mixture was calculated from the experimental data and the activity coefficients that were found to characterize each series of solutions. The pH values of other phenolsulfonate buffers which have m olal ratios, mdm2, of phenolsulfonate ion to bivalent phenolate-sulfonate ion between 2/3 and 1 can be computed for temperatures between 0° and 60° C with an accuracy of ± 0.002 unit from their compositions and from the dissociation constants and ionic parameters given in this paper. The equation used is pH =pK2 -log (mt/m2) -3A.JJ;f(1 +8B-V JL), where JL is the ionic strength . Buffer solutions of this type are suitable pH standards in the range 8.6 to 9.0. CONTENTS

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Section: Discussion Of the Ele Ctromotive-force Meth Odmentioning

confidence: 99%

Section: Methodsmentioning

confidence: 99%

Section: Alkalineibuffer Standardsmentioning

confidence: 99%

Section: Calculation Of the Ph Values From Electromotive-force Measurmentioning

confidence: 99%

See 2 more Smart Citations

The second dissociation constant of p-phenolsulfonic acid and pH values of phenolsulfonate-chloride buffers from 0 degrees to 60 degrees C

Bates¹,

Siegel²,

Acree³

1943

J. RES. NATL. BUR. STAN.

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“…Provisional values for certain buffer solutions have already been published [1]. 2 Very few materials are suitable for use in the standardization of the pH scale in the alkaline range. Borax, N a2B40 7.lOH20, has been frequently used for this purpose, and the pH of solutions of this material, which is approximately 9.2 at 25° C, has been made to cover the range from 8 to 10.5 by the addition of known quantities of acid or base.…”

Section: Introductionmentioning

confidence: 99%

Ionization constant of boric acid and the pH of certain borax-chloride buffer solutions from 0 degrees to 60 degrees C

Manov¹,

DeLollis²,

Acree³

1944

J. RES. NATL. BUR. STAN.

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The ionization constant of boric acid was determined by the use of the electromotive force of cells without liquid junctions. Hydrogen and silver-silverchloride electrodes were immersed in borax-sodium chloride solutions. To reduce tbe possibility of formation of polyborates, dilute solutions were used. The emf were measured at 5-degree intervals over the temperature range 0° to 60° C.The negative common logarithm of the ionization constant (pK) of boric acid over the temperature range 0° to 60° C may be represented by the equation pK=2237.94/T+0.016883T-3.305,where T is the absolute temperature.The data indicate that the mean s.ctivity coefficients of the ions of sodium chloride and of sodium borate do not differ appreciably. The pH values for the solutions studied and for rounded values of the concentration are tabulated as functions of temperature and ionic strength. These solutions, which range in pH values from 8.934 to 9.465, can be used as standards in the calibration of glass-calomel and other electrometric pH equipment.A discussion is given of the significance of the quantity a i (commonly called the "distance of closest approach" of the ions) and its importance in the calculation of pH value~, especially when the buffer ion is not univalent.The changes in free energy, beat content, entropy, and beat capacity that accompany the ionization of 1 mole of boric acid are listed. CONTENTS Page

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The effects of secular calcium and magnesium concentration changes on the thermodynamics of seawater acid/base chemistry: Implications for Eocene and Cretaceous ocean carbon chemistry and buffering

Hain

Sigman

Higgins

et al. 2015

Global Biogeochemical Cycles

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] enhances "ion pairing," which increases seawater buffering by increasing the concentration ratio of total to "free" (uncomplexed) carbonate ion. An increase in [Ca 2+ ], however, also causes a decline in carbonate ion to maintain a given Ω, thereby overwhelming the ion pairing effect and decreasing seawater buffering. ], a doubling of atmospheric CO 2 would require less carbon addition to the ocean/atmosphere system than under modern seawater composition. Moreover, increasing seawater buffering since the Cretaceous may have been a driver of evolution by raising energetic demands of biologically controlled calcification and CO 2 concentration mechanisms that aid photosynthesis.

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The Dissociation Constant of Boric Acid from 10 to 50°

Cited by 63 publications

References 0 publications

The second dissociation constant of p-phenolsulfonic acid and pH values of phenolsulfonate-chloride buffers from 0 degrees to 60 degrees C

The second dissociation constant of p-phenolsulfonic acid and pH values of phenolsulfonate-chloride buffers from 0 degrees to 60 degrees C

Ionization constant of boric acid and the pH of certain borax-chloride buffer solutions from 0 degrees to 60 degrees C

The effects of secular calcium and magnesium concentration changes on the thermodynamics of seawater acid/base chemistry: Implications for Eocene and Cretaceous ocean carbon chemistry and buffering

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