Bicarbonate equilibrium

The pH-buffering of extracellular fluid depends in part on the carbon dioxide/ bicarbonate equilibrium so that the intake of sodium bicarbonate is followed by a brief alkalosis and an increased excretion of sodium carbonate in the urine. Depending on its carbonate concentration, the pH of the urine may rise to 8.07. Large doses (80—100 g/day) of sodium bicarbonate were needed if the pH of stomach contents was to be maintained at 4 or over in patients with duodenal ulcers8. Oxidation of organic anions in the body to carbon dioxide and water permits the use of sodium citrate, lactate or tartrate instead of sodium bicarbonate. In an analogous manner the ingestion of ammonium chloride induces a brief acidosis as a result of the metabolic conversion of ammonia to urea and lowers the pH of the urine. 

At this stage of development the model is restricted to a constant pH reactor and considers only two (pH and volatile acids concentration) of the five variables considered important for monitoring digester operation. This restriction of constant pH can be removed and the model extended to incorporate the interaction with bicarbonate alkalinity by considering the carbon dioxide-bicarbonate equilibrium as shown in Equations 16 and 17. 

Precipitation may be due to an increase in environmental pH and/or a shift in the carbonate bicarbonate equilibrium system, as the result of photosynthetic CO2 incorporation. However, deposition of the carbonates in the immediate vicinity of the cyanobacteria has edso been demonstrated by Friedman (1955) for Geitleria and by Jaag (1945) for Rivularia. Golubic (1973) concluded that incrustation of cyanobacteria with carbonates is not species-specific and not exclusively dependent on photosynthetic activity. Potts and Krumbein (1978) demonstrated that the cyanobacteria Pleuro-capsa minor and Plectonema gloeophilum, caused the precipitation of different kinds of carbonate particles at the same site in desert stromatolites. 

The exchange of HCO3" /OH- across cell membrane results in the accumulation of the hydroxyl ion in the exterior microenvironment, which causes a local pH increase and shifting of the bicarbonate equilibrium toward carbonate see Figure 2.2. The bacterial cell wall may also act as a nucleation site for the precipitation of magnesite by binding the Mg2+ ion. 

The equilibrium equations that normally have to be considered in the EKR modeling of a soil contaminated by heavy metals can be classified into one of the following categories complex formation reactions, precipitation of the metal hydroxides or of other species, ion exchange reactions, surface complexation reactions, etc. Anyway, the autoionization of water always has to be considered and the precipitation of carbonates, together with the carbonate-bicarbonate equilibrium, should normally also be considered. However, the above equations have only considered the species in aqueous phase, so if a species precipitates, a new master species has to be included in this equilibrium system, whose concentration would be the amount of the precipitated species per unit volume of water. This additional degree of freedom is constrained by the solubility product constant of the precipitate (KO, because the new solid phase is in equilibrium with the aqueous phase. If there exists Np precipitated species, the pure-phase equilibria can be represented with the following equation ... 

Various body processes—collectively called metabolism— produce acidic substances that liberate H in solntion. The diffusion of these substances into the bloodstream causes a shift in the carbonic add-bicarbonate equilibrium ... 

The addition of water to carbon dioxide, CO2 + H2O OC(OH)2 H +HCOJ, is formally very similar to its addition to aldehydes and ketones, although here only 0.2% of the carbon dioxide is hydrated at equilibrium, and observations make use of the further equilibrium with and HCO3. A summary of work up to 1958 has been given by Edsall and Wyman,and several later kinetic studies have been made " the hydration process shows general catalysis by basic anions. It is particularly interesting that the enzyme carbonic anhydrase, which is active in maintaining the carbon dioxide-bicarbonate equilibrium in the body, is also an effective catalyst for the hydration of acetaldehyde and other carbonyl compounds. ... 

Because of the great solubility of sulphonic acids in water and the consequent difficulty in crystallisation, the free sulphonic adds are not usually isolated but are converted directly into the sodium salts. The simplest procedure is partly to neutralise the reaction mixture (say, with solid sodium bicarbonate) and then to pour it into water and add excess of sodium chloride. An equilibrium is set up, for example ... 

A species that can serve as both a proton donor and a proton acceptor is called amphiprotic. Whether an amphiprotic species behaves as an acid or as a base depends on the equilibrium constants for the two competing reactions. For bicarbonate, the acid dissociation constant for reaction 6.8... 

The component reactions in eqn. (2) are very fast, and the system exists in equilibrium. Additional carbon dioxide entering the sea is thus quickly converted into anions, distributing carbon atoms between the dissolved gas phase, carbonate and bicarbonate ions. This storage capacity is clear when the apparent equilibrium constants for the two reactions in eqn. (2) are examined, namely... 

X 10 M), and an equivalent amount of OH (its usual concentration in plasma) would swamp the buffer system, causing a dangerous rise in the plasma pH. How, then, can this bicarbonate system function effectively The bicarbonate buffer system works well because the critical concentration of H2CO3 is maintained relatively constant through equilibrium with dissolved CO2 produced in the tissues and available as a gaseous CO2 reservoir in the lungs. ... 

Under the conditions of temperature and ionic strength prevailing in mammalian body fluids, the equilibrium for this reaction lies far to the left, such that about 500 CO2 molecules are present in solution for every molecule of H2CO3. Because dissolved CO2 and H2CO3 are in equilibrium, the proper expression for H2CO3 availability is [C02(d)] + [H2CO3], the so-called total carbonic acid pool, consisting primarily of C02(d). The overall equilibrium for the bicarbonate buffer system then is... 

The most important property of the dissolved solids in fresh waters is whether or not they are such as to lead to the deposition of a protective film on the steel that will impede rusting. This is determined mainly by the amount of carbon dioxide dissolved in the water, so that the equilibrium between calcium carbonate, calcium bicarbonate and carbon dioxide, which has been studied by Tillmans and Heublein and others, is of fundamental significance. Since hard waters are more likely to deposit a protective calcareous scale than soft waters, they tend as a class to be less aggressive than these indeed, soft waters can often be rendered less corrosive by the simple expedient of treating them with lime (Section 2.3). 

When gaseous CO2 is equilibrated with aqueous buffer solution in a closed vessel, a large portion of the CO2 is dissolved in the aqueous phase, mostly in the form of bicarbonate, maintaining the equilibrium of the following three phases ... 

It should be noted that calcium bicarbonate does not exist in the solid state rather, it exists as an unstable salt in water, provided that an excess of free carbon dioxide is available to maintain equilibrium. The reaction is shown below. 

The complexation of Pu(IV) with carbonate ions is investigated by solubility measurements of 238Pu02 in neutral to alkaline solutions containing sodium carbonate and bicarbonate. The total concentration of carbonate ions and pH are varied at the constant ionic strength (I = 1.0), in which the initial pH values are adjusted by altering the ratio of carbonate to bicarbonate ions. The oxidation state of dissolved species in equilibrium solutions are determined by absorption spectrophotometry and differential pulse polarography. The most stable oxidation state of Pu in carbonate solutions is found to be Pu(IV), which is present as hydroxocarbonate or carbonate species. The formation constants of these complexes are calculated on the basis of solubility data which are determined to be a function of two variable parameters the carbonate concentration and pH. The hydrolysis reactions of Pu(IV) in the present experimental system assessed by using the literature data are taken into account for calculation of the carbonate complexation. 

The concentrations of free carbonate and bicarbonate ions determined at solubility equilibrium as a function of pH. Decrements of the concentration near pH = 10 suggest the formation of the Pu(0H)2C03 precipitate and hence lowering solubilities of Pu02 (cf. Figure 2). 

Rainwater and snowmelt water are primary factors determining the very nature of the terrestrial carbon cycle, with photosynthesis acting as the primary exchange mechanism from the atmosphere. Bicarbonate is the most prevalent ion in natural surface waters (rivers and lakes), which are extremely important in the carbon cycle, accoxmting for 90% of the carbon flux between the land surface and oceans (Holmen, Chapter 11). In addition, bicarbonate is a major component of soil water and a contributor to its natural acid-base balance. The carbonate equilibrium controls the pH of most natural waters, and high concentrations of bicarbonate provide a pH buffer in many systems. Other acid-base reactions (discussed in Chapter 16), particularly in the atmosphere, also influence pH (in both natural and polluted systems) but are generally less important than the carbonate system on a global basis. 

Any body tissue in direct equilibrium with the major excretum should reflect the isotopic composition of the diet as a whole. This seems to be the case for bioapatite carbonate, which is thought to be in equilibrium with plasma bicarbonate, which itself is in equilibrium with respired CO2. In fact the Ambrose and Norr (1993) and Tieszen and Fagre (1993) data sets (among others) show clearly that the bioapatite carbonate differs from total diet (or respired CO ) by an amount approximating to the equilibrium isotopic fractionation in the system (Mook 1989) ... 

Acid-base reactions of buffers act either to add or to remove hydrogen ions to or from the solution so as to maintain a nearly constant equilibrium concentration of H+. For example, carbon dioxide acts as a buffer when it dissolves in water to form carbonic acid, which dissociates to carbonate and bicarbonate ions ... 

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