Activity Coefficient of the Magnesium Ion in Sea Water

1965 ◽  
Vol 22 (1) ◽  
pp. 113-116 ◽  
Author(s):  
R. F. Platford
Keyword(s):  
Activity Coefficient ◽  
Sodium Hydroxide ◽  
Sea Water ◽  
Magnesium Ion ◽  
Activity Product ◽  
Hydroxyl Ion ◽  
Ion Activity ◽  
Total Magnesium

Sodium hydroxide was added to carbonate-free sea water and the point at which Mg(OH)2 precipitated was detected optically. At this point the hydroxyl ion activity was calculated from the measured pH of the solution and the magnesium ion activity was calculated from the activity product for Mg(OH)2.From a knowledge of the magnesium ion activity and other data, it was estimated that about 3% of the total magnesium in the sea water was in the form of [Formula: see text].

scholarly journals Some Chemical Considerations on the Distribution of Iron in the Sea

1948 ◽  
Vol 27 (2) ◽  
pp. 314-321 ◽  
Author(s):  
L. H. N. Cooper
Keyword(s):  
Ferric Iron ◽  
Theoretical Study ◽  
Sea Water ◽  
Precipitated Phase ◽  
Ferric Phosphate ◽  
Biological Interest ◽  
Hydroxyl Ion ◽  
Ion Activity ◽  
High Dilutions ◽  
Existing Data

A theoretical study of the behaviour of ferric phosphate in sea water has been made with existing data. The conditions least favourable for the continued existence of ferric phosphate should occur in the upper layers of the sea during active plant growth. The formation and stability of ferric phosphate is likely to be most favoured under the more acid conditions found in the gut of many animals. Ferric phosphate is likely to be introduced into sea water in animal faeces. Experimental confirmation of this deduction is essential.In all experimental work involving ferric phosphate in sea water, the need is indicated for very accurate control of hydroxyl-ion activity and probably pressure, and for attention to the time factor and to the physical condition of the precipitated phase. The topic is of geochemical as well as of biological interest.Dipyridyl is a reagent for ferrous iron. Reasons are set out why in sea and lake water at high dilutions it may also determine ferric iron without presence of an added reducing agent.Evidence is presented that ferrifluoride plays no part in the iron cycle in he sea.

Speciation and Geochemical Implications of Carcass Burial Leachate

2018 ◽  
Vol 61 (2) ◽  
pp. 559-570
Author(s):  
Dyan L. Pratt ◽  
Terrance A. Fonstad
Keyword(s):  
Chemical Composition ◽  
Ionic Strength ◽  
Activity Coefficient ◽  
Swine Manure ◽  
Burial Site ◽  
Mass Mortality Event ◽  
Ammonium N ◽  
Ion Activity ◽  
Common Activity ◽  
The Impact

Abstract. In the event of a mass livestock mortality situation, disposal routes such as burial are commonly chosen. The impact of burial on the environment could be substantial, but the composition of the leachate arising from a burial site has not been well documented. This study was performed to determine the chemical composition of leachate arising from animal mortalities in a burial setting. Three species of livestock were used: bovine, swine, and poultry. Leachate collected from lined burial pits over two years of decomposition was analyzed for major and minor ions. Analysis indicated that livestock mortality leachate contains, on average, concentrations of 46,000 mg L-1 of alkalinity (as bicarbonate), 12,600 mg L-1 of ammonium-N, 2600 mg L-1 of chloride, 3600 mg L-1 of sulfate, 2300 mg L-1 of potassium, 1800 mg L-1 of sodium, and 1500 mg L-1 of phosphorus, along with lesser amounts of iron, calcium, and magnesium. Select samples had maximum concentrations of ammonium-N and bicarbonate up to 50% higher than these average values. In comparison to earthen swine manure storages and landfills, the ionic strength of the leachate was 2 to 4 times higher, and therefore its impact on water resources could be greater. Following the study of the chemical composition of livestock mortality leachate, the potential impacts of this leachate on the soil/water systems below a burial site were investigated. The ionic strength of the leachate presents its own set of challenges. Basic modeling of ion activity using the five most common activity coefficient equations (Debye-Hückel, extended Debye-Hückel, Truesdell-Jones, Davies, and Pitzer) were considered to assess the sensitivity of these methods for calculated ion activity as impacted by the ionic strength of the leachate. This was completed to further enhance the modeling and speciation efforts. Based on the results and the applicability of the Truesdell-Jones equation, PHREEQC was used to assess the chemical speciation of the leachate. The speciation of this leachate provides evidence of phosphate and sulfate compounds available for potential unattenuated transport. Understanding the geochemical implications of livestock mortality burial will give scientists and regulators more information for performing future risk analyses when considering mortality burial as a management option, either routinely or during a mass mortality event. Keywords: Ion activity coefficient, Ionic strength, Leachate chemical composition, Livestock burial leachate, Speciation.

Which value for the first dissociation constant of carbonic acid should be used in biological work?

1991 ◽  
Vol 260 (5) ◽  
pp. C1113-C1116 ◽  
Author(s):  
R. W. Putnam ◽  
A. Roos
Keyword(s):  
Ionic Strength ◽  
Activity Coefficient ◽  
Dissociation Constant ◽  
Carbonic Acid ◽  
Mass Action ◽  
Bicarbonate Concentration ◽  
Apparent Dissociation Constant ◽  
Logarithmic Form ◽  
Ion Activity ◽  
Hasselbalch Equation

The apparent first dissociation constant of carbonic acid has been defined in different ways in the literature. Harned and co-workers (8-10) have defined it in terms of molalities of the participating species, including H ions: Ks = mHmHCO3/mCO2. In contrast, Hastings and Sendroy have defined an apparent constant in which acidity is expressed as H ion activity: K'1 = aHmHCO3/mCO2. These constants differ by a factor gamma H, the activity coefficient of H ions at the prevailing ionic strength. Therefore, pK'1 is greater than pKs by an amount equal to -log gamma H, which, at mu = 0.16 M, is approximately 0.1. It is important that the correct value for the apparent dissociation constant or its logarithmic form be entered in the mass action expression or in the Henderson-Hasselbalch equation in order to prevent significant errors in the computation by means of these equations of quantities that cannot be directly measured. Specifically, for the derivation of bicarbonate concentration from PCO2 and pH (-log aH), pK'1 is to be used and not an uncorrected pKs.

scholarly journals Dissolution, Stability and Solubility of Tooeleite [Fe6(AsO3)4(SO4)(OH)4·4H2O] at 25–45 °C and pH 2–12

Minerals ◽  
2020 ◽  
Vol 10 (10) ◽  
pp. 921 ◽  
Author(s):  
Zongqiang Zhu ◽  
Jun Zhang ◽  
Yinian Zhu ◽  
Jie Liu ◽  
Shen Tang ◽  
...  
Keyword(s):  
Aqueous Medium ◽  
Solubility Product ◽  
Arsenic Concentration ◽  
Dissolved Iron ◽  
Activity Product ◽  
Free Energy Of Formation ◽  
Ion Activity ◽  
Initial Ph ◽  
Energy Of Formation

Tooeleite [Fe6(AsO3)4(SO4)(OH)4·4H2O] was synthesized and characterized to investigate its possible immobilization for arsenic in acidic and alkali environments by a long-term dissolution of 330 d. The synthetic tooeleite was platy crystallites of ~1μm across, giving the lattice parameters of a = 6.4758 Å, b = 19.3737 Å and c = 8.9170 Å. For the tooeleite dissolution, the dissolved arsenic concentration showed the lowest value of 427.3~435.8 mg/L As at initial pH 12 (final pH 5.54). The constituents were dissolved preferentially in the sequence of SO42− > AsO33− > Fe3+ in the aqueous medium at initial pH 2–12. The dissolved iron, arsenite and sulfate existed mainly as FeSO4+/Fe3+, H3AsO30 and SO42− at initial pH 2, and in the form of Fe(OH)30/Fe(OH)2+, H3AsO30 and SO42− at initial pH 12, respectively. The tooeleite dissolution was characterized by the preferential releases of SO42− anions from solid surface into aqueous medium, which was fundamentally controlled by the Fe-O/OH bond breakages and the outer OH− group layers. From the data of the dissolution at 25 °C and initial pH 2 for 270–330 d, the ion-activity product [logˍIAP], which equaled the solubility product [Ksp] at the dissolution equilibrium, and the Gibbs free energy of formation [ΔGfo] were estimated as −200.28 ± 0.01 and −5180.54 ± 0.07 kJ/mol for the synthetic tooeleite, respectively.

INCREASE IN NEUTRAL SALT EXTRACTABLE CATION EXCHANGE CAPACITY OF SOME ACID SOILS AS AFFECTED BY CaSO4 APPLICATIONS

1984 ◽  
Vol 64 (2) ◽  
pp. 153-161 ◽  
Author(s):  
S. SHAH SINGH
Keyword(s):  
British Columbia ◽  
Cation Exchange ◽  
Cation Exchange Capacity ◽  
Acid Soils ◽  
Exchange Capacity ◽  
Neutral Salt ◽  
Activity Product ◽  
Three Samples ◽  
Ion Activity ◽  
Extractable Cations

The equilibration of acid soils, a Sombric Ferro-Humic Podzol (CSSC-2) from British Columbia, an Orthic Ferro-Humic Podzol (CSSC-19) from Quebec and two horizons of a Dystric Brunisol (SSD-330, SSD-331) from British Columbia, with CaSO4 solution demonstrated that SO4 ions reacted with components of these soils. These reactions increased soil pH, ion activity product (Al)(OH)3 and neutral salt extractable exchangeable cations. The increase in pH and ionic activity product (Al)(OH)3 were noticeable on a single equilibration; however, increase in neutral salt extractable cations was only observed after subsequent equilibrations. After three equilibrations, the sums of NaCl extractable cations were 6.56, 11.99, 5.62 and 4.31 meq/100 g for soil samples CSSC-2, CSSC-19, SSD-330 and SSD-331, respectively. The corresponding values for the unequilibrated soils were 5.20, 7.49, 4.30 and 2.50 meq/100 g. On further equilibration there was no increase in total extractable cations for sample CSSC-2; however, for the other three samples there were increases which became progressively smaller. The reaction of SO4 ions with aluminum hydroxy clay complexes seems to be the mechanism for the increase of negative sites. Key words: Cation exchange capacity, CaSO4, acid soils

scholarly journals SALT ERROR OF INDICATORS DUE TO STANDARD ALKALINE BUFFERS THEMSELVES. II

1929 ◽  
Vol 12 (5) ◽  
pp. 695-710 ◽  
Author(s):  
J. W. McBain ◽  
M. E. Laing ◽  
O. E. Clark
Keyword(s):  
Sodium Hydroxide ◽  
Aqueous Sodium Hydroxide ◽  
Thymol Blue ◽  
Phenol Red ◽  
Dilute Solutions ◽  
Aqueous Sodium ◽  
Previous Communication ◽  
Alizarin Yellow ◽  
Hydroxyl Ion

Previous results of the comparison of colors given by indicators in alkaline buffers and pure aqueous sodium hydroxide have been repeated and confirmed. The electrometric determinations show that the sodium hydroxide was pure and gave theoretical values for the concentration of hydroxyl ion. The slight but distinct neutralising effect of dilute solutions of alkali has been measured electrometrically and the allowances to be made are recorded graphically. It is found that whereas alizarin yellow G, tropæolin O and thymol violet may be used without appreciable error (in accordance with our previous communication) the grave discrepancies remain for phenolphthalein, o-cresol phthalein and thymol blue and phenol red which must be ascribed to salt error in the alkaline buffer itself.

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