Carbonic acid equilibrium

Cement mortar will be attacked by waters that have an excess of free carbon dioxide compared with that of waters that are in a lime-carbonic acid equilibrium. There is a two-step mechanism with a carbonization process according to... 

At equilibrium, the concentration of H+ will remain constant. When a strong acid (represented by H+ or HA) is introduced into solution, the concentration of H+ is increased. The buffer compensates by reacting with the excess H ions, moving the direction of the above reaction to the left. By combining with bicarbonate and carbonate ions to form the nonionic carbonic acid, equilibrium is reestablished at a pH nearly the same as that existing before. The buffer capacity in this case is determined by the total concentration of carbonate and bicarbonate ions. When no more carbonate or bicarbonate ions are available to combine with excess H+ ions, the buffer capacity has been exceeded and pH will change dramatically upon addition of further acid. 

A specific example from this patent was discussed in Chapter 8 where it was offered that extraction of carboxylic acids from solutions of pH greater than pK was not a very attractive process because all the anion is not converted to the acid form by action of the carbonic acid equilibrium. The absffact of the patent states that CO2 will react with the salt to form the (free) carboxylic acid which can dissolve in CO2. That is suictly correct, but we always suggest that the absolute values be examined to determine such factors as yield and distribution coefficient. 

A weak acid that can be found in rain water in relatively large quantities is a carbonic acid the minimum pH value of a solution of carbonic acid in rain water is about 5.5. In a rain water sample having a pH 5.5 the carbonic acid equilibrium is reversed and CO2 will disappear from the sample. In such samples the pH value is entirely determined by the free H ions of strong mineral acids such as sulfuric acid and nitric acid. 

If the water to be analyzed lies in the range of the so-called lime/carbonic acid equilibrium, i.e. approximately in the range of calcium carbonate saturation, the difference in pH before and after adding calcium carbonate is smaller than 0.04. 

Respiratory acidosis is sometimes the result of slow or shallow breathing (hypoventilation) caused by an overdose of narcotics or barbiturates. Anesthetists need to be particularly aware of respiratory acidosis, because most inhaled anesthetics depress respiration rates. Lung diseases, such as anphysema and pneumonia, or an object lodged in the windpipe can also result in hypoventilation. In respiratory acidosis, too little carbon dioxide is exhaled, and its concentration within the blood increases. The carbon dioxide/carbonic acid equilibrium shifts to the left to restore the equilibrium ... 

Urea (the diamide of carbonic acid) can be prepared by the historic method of Wohler. When an aqueous solution of ammonium cyanate is allowed to stand, the cyanate undergoes molecular rearrangement to urea, and an equilibrium mixture containing about 93% of urea is thus formed. Urea is... 

The equilibrium constant for the overall reaction is related to an apparent equilibrium constant Ki for carbonic acid ionization by the expression... 

Section 19 9 Carbon dioxide and carbonic acid are m equilibrium m water Carbon dioxide IS the major component... 

Phenol IS a weaker acid than carbonic acid The equilibrium shown lies to the left... 

In the discussion of the relative acidity of carboxylic acids in Chapter 1, the thermodynamic acidity, expressed as the acid dissociation constant, was taken as the measure of acidity. It is straightforward to determine dissociation constants of such adds in aqueous solution by measurement of the titration curve with a pH-sensitive electrode (pH meter). Determination of the acidity of carbon acids is more difficult. Because most are very weak acids, very strong bases are required to cause deprotonation. Water and alcohols are far more acidic than most hydrocarbons and are unsuitable solvents for generation of hydrocarbon anions. Any strong base will deprotonate the solvent rather than the hydrocarbon. For synthetic purposes, aprotic solvents such as ether, tetrahydrofuran (THF), and dimethoxyethane (DME) are used, but for equilibrium measurements solvents that promote dissociation of ion pairs and ion clusters are preferred. Weakly acidic solvents such as DMSO and cyclohexylamine are used in the preparation of strongly basic carbanions. The high polarity and cation-solvating ability of DMSO facilitate dissociation... 

Table 4-1 lists some rate constants for acid-base reactions. A very simple yet powerful generalization can be made For normal acids, proton transfer in the thermodynamically favored direction is diffusion controlled. Normal acids are predominantly oxygen and nitrogen acids carbon acids do not fit this pattern. The thermodynamicEilly favored direction is that in which the conventionally written equilibrium constant is greater than unity this is readily established from the pK of the conjugate acid. Approximate values of rate constants in both directions can thus be estimated by assuming a typical diffusion-limited value in the favored direction (most reasonably by inspection of experimental results for closely related... 

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... 

In this chapter it is clearly impossible to do more than sample the extensive literature on the carbon acidity of sulfinyl and sulfonyl compounds, as it illuminates the electronic effects of these groups, particularly in connection with linear free-energy relationships. There are three main areas to cover first, as already indicated, equilibrium acidities (pKa values) second, the kinetics of ionization, usually studied through hydrogen isotopic exchange and finally, the kinetics of other reactions proceeding via carbanionic intermediates. 

Water often is a reagent in an aqueous equilibrium. For example, when carbon dioxide dissolves in water, it reacts with a water molecule to form carbonic acid ... 

The human body generates a steady flow of acidic by-products during its normal metabolic processes. Foremost among these is carbon dioxide, which is a major product of the reactions the body uses to produce energy (see Section 14-). An average person produces from 10 to 20 mol (440 to 880 g) of CO2 every day. Blood carries CO2 from the cells to the lungs to be exhaled. In aqueous solution, dissolved CO2 is in equilibrium with carbonic acid H2 O + CO2 H2 CO3... 

This proton transfer reaction involves the second acidic hydrogen atom of carbonic acid, so the appropriate equilibrium constant is. a 2 > whose p is found in Appendix E p. a 2 — 10.33. Because this is a buffer solution, we apply the buffer equation ... 

By comparing the approximate pK values of the bases with those of the carbon acid of interest, it is possible to estimate the position of the acid-base equilibrium for a given reactant-base combination. For a carbon acid C-H and a base B-H,... 

A chemical reaction in which the products react to re-form the original reactants is called a reversible reaction. For example, club soda is a mixture of carbon dioxide gas and water. The water and carbon dioxide react forming carbonic acid (H2C03). Carbonic acid decomposes to again form water and carbon dioxide. A state of equilibrium is reached in which the amounts of carbonic acid, water, and carbon dioxide remain constant. The overall reaction can be written as follows. 

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 ... 

In considering relative acidity, classically it is only the thermodynamics of the situation that are of interest in that the pKa value for the acid (cf. p. 54) can be derived from the equilibrium above. The kinetics of the situation are normally of little significance, as proton transfer from atoms such as O, N, etc., is extremely rapid in solution. With carbon acids such as (1), however, the rate at which proton is transferred to the base may well be sufficiently slow as to constitute the limiting factor the acidity of (1) is then controlled kinetically rather than thermodynamically (cf. p. 280). 

The partial ionization of carbonic acid produces hydronium ion, H+, driving the indicator equilibrium to the weak acid form. A colorless solution results. As the water in the ink evaporates, the white residue of sodium carbonate remains. 

In water the bicarbonate ion is in equilibrium with carbonic acid ... 

The small value of the equilibrium constant indicates that the formation of carbonic acid is not very extensive in neutral water. However, the formation of carbonic acid is quite favored in acidic solution (arising from the citric acid also contained in the product) ... 

In real systems (hydrocarbon-02-catalyst), various oxidation products, such as alcohols, aldehydes, ketones, bifunctional compounds, are formed in the course of oxidation. Many of them readily react with ion-oxidants in oxidative reactions. Therefore, radicals are generated via several routes in the developed oxidative process, and the ratio of rates of these processes changes with the development of the process [5], The products of hydrocarbon oxidation interact with the catalyst and change the ligand sphere around the transition metal ion. This phenomenon was studied for the decomposition of sec-decyl hydroperoxide to free radicals catalyzed by cupric stearate in the presence of alcohol, ketone, and carbon acid [70-74], The addition of all these compounds was found to lower the effective rate constant of catalytic hydroperoxide decomposition. The experimental data are in agreement with the following scheme of the parallel equilibrium reactions with the formation of Cu-hydroperoxide complexes with a lower activity. 

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