Sulfurous acid equilibrium constant

Sulfuric acid is constantly in equilibrium with SO3 and water ... 

Equilibrium moisture content of a hygroscopic material may be determined in a number of ways, the only requirement being a source of constant-temperature and constant-humidity air. Determination may be made under static or dynamic conditions, although the latter case is preferred. A simple static procedure is to place a number of samples in ordinaiy laboratoiy desiccators containing sulfuric acid solutions of known concentrations which produce atmospheres of known relative humidity. The sample in each desiccator is weighed periodically until a constant weight is obtained. Moisture content at this final weight represents the equilibrium moisture content for the particular conditions. 

This ease with which we can control and vary the concentrations of H+(aq) and OH (aq) would be only a curiosity but for one fact. The ions H+(aq) and OH (aq) take part in many important reactions that occur in aqueous solution. Thus, if H+(aq) is a reactant or a product in a reaction, the variation of the concentration of hydrogen ion by a factor of 1012 can have an enormous effect. At equilibrium such a change causes reaction to occur, altering the concentrations of all of the other reactants and products until the equilibrium law relation again equals the equilibrium constant. Furthermore, there are many reactions for which either the hydrogen ion or the hydroxide ion is a catalyst. An example was discussed in Chapter 8, the catalysis of the decomposition of formic acid by sulfuric acid. Formic acid is reasonably stable until the hydrogen ion concentration is raised, then the rate of the decomposition reaction becomes very rapid. 

Sulfur in the +4 oxidation state also forms an oxyacid, sulfurous acid (HjSOj). This compound is not as strong an acid as H2S04. The equilibrium constant for the reaction... 

In the lead-acid battery, sulfuric acid has to be considered as an additional component of the charge-discharge reactions. Its equilibrium constant influences the solubility of Pb2+ and so the potential of the positive and negative electrodes. Furthermore, basic sulfates exist as intermediate products in the pH range where Fig. 1 shows only PbO (cf. corresponding Pour-baix diagrams in Ref. [5], p. 37, or in Ref. [11] the latter is cited in Ref. [8]). Table 2 shows the various compounds. 

That the photoreactive species is the carbonium ion and not the corresponding alcohol is clearly indicated by the relative concentrations of the two species present. The calculated equilibrium constant for 5% aqueous sulfuric acid implies an alcohol content of 2 7 x 10 %, much too low to account for any detectable photoreaction from this covalent species. In addition, when the tropylium salt is irradiated in the absence of acid, neither 3 nor 4 is detected as a product, but rather ditropyl (5) and its photoisomer (6) are observed. 

If ri = rate of reaction of SA and r2 = rate of formation of DES, then rate of formation of MES = r2 - ri. The equilibrium constant for each of the esterification reactions was determined by allowing the reaction to proceed until no further reaction was observed. For the quantities of resin used, the rate of resin-catalyzed esterification is much slower than that for sulfuric acid-catalyzed reaction. All results are summarized in Table 1. 

It will be noted that there is a factor of approximately 105 between successive dissociation constants. This relationship exists between the equilibrium constants for numerous polyprotic acids, and it is sometimes known as Pauling s rule. This rule is also obeyed by sulfurous acid, for which ffj = 1.2 X 10 2 and K2 = 1 X 10 7. 

For this equilibrium, the ion product constant has a value of approximately 2.7 X 1CP4. However, the discussion is complex and other species are present in sulfuric acid as a result of equilibria that can be written as... 

Upon addition of a solution of sulfuric acid in D20 the reaction of A-acetoxy-A-alkoxyamides obeys pseudo-unimolecular kinetics consistent with a rapid reversible protonation of the substrate followed by a slow decomposition to acetic acid and products according to Scheme 5. Here k is the unimolecular or pseudo unimolecular rate constant and K the pre-equilibrium constant for protonation of 25c. Since under these conditions water (D20) was in a relatively small excess compared with dilute aqueous solutions, the rate expression could be represented by the following equation ... 

It is necessary to consider a number of equilibrium reactions in an analysis of a hydrometallurgical process. These include complexing reactions that occur in solution as well as solubility reactions that define equilibria for the dissolution and precipitation of solid phases. As an example, in analyzing the precipitation of iron compounds from sulfuric acid leach solutions, McAndrew, et al. (11) consider up to 32 hydroxyl and sulfate complexing reactions and 13 precipitation reactions. Within a restricted pH range only a few of these equilibria are relevant and need to be considered. Nevertheless, equilibrium constants for the relevant reactions must be known. Furthermore, since most processes operate at elevated temperatures, it is essential that these parameters be known over a range of temperatures. The availability of this information is discussed below. 

Measurements of the solubility of HBr in sulfuric acid at 220 K gave Henry s law constants from 8.5 X 103 M atm-1 for 72 wt% H2S04 to 1.5 X 107 M atm-1 for 54 wt% H2S04 (Williams et al., 1995 Abbatt, 1995). Application of these values to stratospheric aerosol particles typical of midlatitude conditions gives very small equilibrium concentrations of dissolved HBr i.e., most of the HBr will remain in the gas phase. 

The equilibrium constant is 0.013 at 18°C. Sulfur dichloride reacts violently with water, forming hydrogen chloride, sulfur dioxide, hydrogen sulfide, sulfur, and a mixture of thionic acids. 

The "equilibrium boxes" for the solvents (Fig. 10-1) indicate the range over which differentiation occurs outside the range of a particular solvent, all species are leveled. For example, water can differentiate species (i.e., they are weak adds and bases) with pKa s from about 0 to 14 (such as acetic acid). Ammonia, on the other hand, behaves the same toward acetic acid and sulfuric acid because both lie below the differentiating limit of —12. The extent of these ranges is determined by the autoionization constant of the solvent (e.g, —14 units for water). The acid-base behavior of several species discussed previously may be seen to correlate with Fig. 10.1.11... 

PK. A measurement of the complete ness of an incomplete chemical reaction. It is defined as the negative logarithm ito the base 101 of the equilibrium constant K for the reaction in question. The pA is most frequently used to express the extent of dissociation or the strength of weak acids, particularly fatty adds, amino adds, and also complex ions, or similar substances. The weaker an electrolyte, the larger its pA. Thus, at 25°C for sulfuric add (strong acid), pK is about -3,0 acetic acid (weak acid), pK = 4.76 bone acid (very weak acid), pA = 9.24. In a solution of a weak acid, if the concentration of undissociated acid is equal to the concentration of the anion of the acid, the pAr will be equal to the pH. 

Exercise 18-9 Use bond energies and the stabilization energy of ethanoic acid (18 kcal mole-1, Section 18-2A) to calculate AH° for the addition of water to ethanoic acid to give 1,1,1-trihydroxyethane. Compare the value you obtain with a calculated AH° for the hydration of ethanal in the vapor phase. Would you expect the rate, the equilibrium constant, or both, for hydration of ethanoic acid in water solution to be increased in the presence of a strong acid such as sulfuric acid Explain. 

Having found equilibrium constants for the series of bases, we may now use them to characterize the proton-donating ability of any mixture of sulfuric acid and water. Rearranging Equation 3.30, we have... 

Diethyl ether, used as an anesthetic, is synthesized by heating ethanol with concentrated sulfuric acid. Write the equilibrium constant expression for Kc. 

Ethyl Acetate. The esterification of ethanol by acetic acid was studied in detail over a century ago (357), and considerable literature exists on determinations of the equilibrium constant for the reaction. The usual catalyst for the production of ethyl acetate [141-78-6] is sulfuric acid, but other catalysts have been used, including cation-exchange resins (358), a-fluoronitrites (359), titanium chelates (360), and quinones and their pardy reduced products. 

The parameters of the Michaelis-Menten type kinetics were calculated for the reactions and are summarized in Table II. The apparent Michaelis constant values (Km) are rather large, indicating that the concentration of the complex at the equilibrium state is not high, unlike ordinary enzymatic reactions. The ratio of kJKm against the second-order rate constant with sulfuric acid (k2) can be considered to be an indication of the rate enhancement. The ratio increased with increasing mole fraction of the vinyl alcohol repeating unit in the copolymer and with... 

Possible methods of determining the extent of protonation include absorption spectroscopy at a wavelength for which species with n and fir-1 protons have different extinction coefficients, freezing point depression, and electrical conductivity (15). Of these, we have utilized only spectroscopy, which has the disadvantage that only the equilibrium constants for the most highly protonated states are accessible if, as is usual, the species with low protonation are insoluble. In this method, the extinction coefficient c of the compound is determined as a function of the H2SO4 content in the sulfuric acid solvent and correlated with the Hammet acidity function H0 (18) to give the pKB value of the protonated species,... 

The important point is that, at any one particular temperature, the equilibrium constant is just that—constant. This gives us a means of forcing the equilibrium to favour the products (or reactants) since the ratio of the two must remain constant. Therefore, if we increase the concentration of the reactants (or even that of just one of the reactants), more products must be produced to keep the equilibrium constant. One way to make esters in the laboratory is to use a large excess of the alcohol and remove water continually from the system as it is formed, for example by distilling it out. This means that in the equilibrium mixture there is a tiny quantity of water, lots of the ester, lots of the alcohol, and very little of the carboxylic acid in other words, we have converted the carboxylic acid into the ester. We must still use an acid catalyst, but the acid must be anhydrous since we do not want any water present—commonly used acids are toluene sulfonic acid (tosic acid, TsOH), concentrated sulfuric acid (H2SO4), or gaseous HC1. The acid catalyst does not alter the position of the equilibrium it simply speeds up the rate of the reaction, allowing equilibrium to be reached more quickly. 

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