Carbonic acid first dissociation constant

Table 9 includes data on the first dissociation constants of seven weak acids it will be recalled that we expect these to fall into class III. The table includes the second dissociation constants of five acids, phosphoric, sulfuric, oxalic, malonic, and carbonic, which fall into class IV, while the amino acids glycine and alanine provide four examples that should fall into class II. 

Thus, k = k2K2/(KsK]) if aHC0(s) - aHC0- at equilibrium, if, is the first dissociation constant of carbonic acid. Equation (136) becomes... 

It is interesting that the presence of a charged atom in a molecule induces a fractional charge not only on the other atoms of the chain but also on those atoms which are near by in space. Thus if the electrostatic attraction were conducted only through the carbon skeleton of a molecule, we should expect to find the dissociation constants of fumaric and maleic acids to be identical. In actual fact the first dissociation constants are somewhat similar, but there is a considerable difference in the second dissociation constants. A possible explanation of this phenomenon is that in the case of the mono acid salts of maleic acid, the near-by negative charge on the carboxylate group hinders the removal of the second proton. 

Table III. Data for the Computation of the First Dissociation Constant of Carbonic Acid. Solution Equal Concentrations of KHCO,...

In clean natural water the pH can be calculated from the content of free CO2 and hydrogen carbonates using the expression for the first dissociation constant of carbonic acid. Dissociation to the 2nd degree can be neglected as its effect becomes significant only as pH > 8.3. Due to the inaccurate determination of free CO2 the calculation provides only rough results. On the contrary, from a known value of pH and the content of HCO3 the content of free CO2 can be calculated. 

Where k is the first dissociation constant of carbonic acid. An internal pH-electrode is used to register the pH-value of this (internal) buffer. The pH-value then directly correlates with the C02-partial pressure in the culture medium. 

First dissociation constant of carbonic acid (/iTi) in seawater as a function of temperature and salinity (after Roy et al, 1993,1996 DOE, 1994). 

Ascorbic acid is a white, crystalline compound, with a melting ppetroleum ether. In aqueous solution the hydrc en atom of the enol group on carbon atom 3 dissociates, yielding a solution with a pH of about 3. Thus, in aqueous solution, ascorbic add behaves as a monobasic acid, forming salts containing one monovalent metal atom or equivalent. In alkalme solution the hydrogen of the enol group on carbon atom 2 dissociates and is i laced by metal. The of the first dissociation constant is 4.17 and that of the second dissociation constant is 11.57. 

C is the concentration of heme iron in the solution, j8 is the buffering power of hemoglobin per heme, K4 is the first dissociation constant of carbonic acid, q is the solubility of carbon dioxide, and A is a constant. The first term on the left represents the negative charge carried by the protein, which is very nearly linear in pH in the range 6.5 < pH < 7.8 and the second term on the left represents the concentration of bicarbonate ion. /3 may be taken as 2.9 equivalents per mole iron per pH unit, the value obtained for horse hemoglobin from the results of Ger.man and Wyman (56). The value of the constant A depends on the amount of base present in the solution. 

TTie first dissociation constant is typically greater than the second i.e., K i > K,2- For example, the weak unstable carbonic acid [H2CO3) can lose one proton to form bicarbonate anion (HC03 ) and lose a second to form carbonate anion [C03 "). Both Ka values are small, but K,] > K. ... 

Care must be taken when quoting and using the first dissociation constant of carbonic acid. In aqueous solution carbonic acid only exists in equilibrium with carbon dioxide, and the concentration of H2CO3 is much lower than the dissolved CO2 concentration. Since it is not possible to distinguish between H2CO3 and dissolved CO2 (referred to as C02(aq)) by conventional methods, H2CO3 is used to represent the two species when writing the aqueous chemical equilibrium equation. The equation may be rewritten as follows cf. sulfurous acid) ... 

There are four reactions that deal explicitly with the H+ ion (Table 3.3) one is the dissociation constant for water (Kw), two are the first and second dissociation constants of carbonic acid (K and K2), and the fourth deals with the dissociation of the bisulfate (HSOJ) ion (iFbisuifate)-... 

When this is the case, the heat of reaction must be quite independent of the nature of the anion and of the cation, aa these are not affected by the reaction. This is clearly true for nitric and hydrochloric acids with all the bases given in the table. For sulphuric and carbonic acids, however, the conditions for the validity of the theory are apparently not fulfilled. In the first case, the heat of dilution of sulphuric acid amounts to 2000 cal., and this amount must be subtracted from the figure given in the table, as it is evolved when the alkali and acid are mixed. In the second case, carbonic acid is so weak an acid that it is practically undissociated. The heat necessary for the dissociation into ions therefore uses up part of the heat of neutralisation. From the table it follows that the electrolytic dissociation of J mol. HgCOg requires 13700 — 10200 = 3500 calories. The constant heat of neutrahsation 13700 cal. is the heat of ionisation of water, i.e, the quantity of heat required for the dissociation of water, and liberated on the combination of its ions. 

Note that in this case, only intermolecular hydrogen bonding is possible, because of the rigid nature of the central carbon/carbon double bond. The difference in the stability of the intermolecular versus the intramolecular hydrogen bond is reflected in the difference in the acidities of the first hydrogen to be removed in each case. The trans isomer has a dissociation constant of 3.02, while for the cis isomer it is 1.92, i.e. the latter is a stronger acid. 

Equation 3 is not important in the atmosphere, so the pH of a droplet of water in equilibrium with atmospheric C02 can be determined by combining the first two equilibrium constant equations that govern the dissolution (i.e. Henry s law, as discussed in Box 3.4) and dissociation. If carbonic acid (H2C03) is the only source of protons, then aH+ must necessarily equal aHCOj. Thus the equilibrium equation for equation 2 can be written ... 

Where is the total concentration of acid species and A, and K2 are the first and second stepwise dissociation constants of the acids. This equation can be used to compute the buffer index of a polyprotic acid as long as successive dissociation constants differ by at least 20 times (this assures a calculation error of 5% or less). In other words, for a diprotic acid K2fK should be less than 0.05 (cf. Butler 1964). Thus, for example, Eq. (5.114) may be used to compute the buffer index due to species of carbonic acid, for which A = 10 and K2 = 10" °, or/ for species of silicic acid, for which a , = lO- and K, = 10... 

The 5,8-dihydroxy-4(3//)-quinazolinones possess three protons capable of dissociation in aqueous buffer, viz. N3—H and the protons of the 5- and 8-hydroxyl groups (Scheme 107). Initial acid dissociation of the N3 proton does not take place at low pH because electron-rich substituted quinazolin-4(3//)-ones possess pK values greater than 10. The first dissociation is from the 5-hydroxyl which affords an anion (697) which is stabilized by internal hydrogen bonding expressed by pK 7.3 (7.8). Cleavage of the halogen-carbon bond in the 2-substituent in (697) involves a rate-determining reaction of the hydroquinone monoanion and dianion species, since there is an approximate 100-fold difference in rate constants between the precursor 2-chloro (696 X = Cl) and... 

The equilibrium constant for this is deduced from the solubility product of CaCOa (K3), and the dissociation constants of carbonic acid to the first degree (A j) and to the second degree (A a) the equilibrium constant (K) is thus given by the relation K =. ... 

The first meaningful deconvolution of titration curves with theoretical description of acid-based dissociation on the surface of carbons was descried in the mid-90 s of the last century [213-215]. In this approach it is assumed that the system under study consists of acidic sites characterized ly their acidity constants, Ka. It is also assumed that the population of sites can be described by a continuous pK distribution, flpKa). The experimental data can be transformed into a proton binding isotherm, Q, representing the total amount of protonated sites, which is related to the pKa distribution by the following integral equation (cf. Fig. 10) ... 

Carbonate chemistry was incorporated into the model in order to verify realistic geochemical conditions. First, the carbonate chemistry is characterized by the dissociation of carbonic acid yielding HCO3 and COf. The two dissociation constants K,. and K2c are defined as follows... 

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