Hal 47-53.docx

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pH of the solution 14 – 2.9 or 11.1. Thus a ditute solution of a solute carbonate is fairly strongly alkaline Hydrolysis of CuSO4. The hydrolysis reaction is

For which an equilibrium constant can be set up :

In this expression, multiplication of both numerator and denominator of K by (OH-) makes it possible to evaluate the constant as a product of the reciprocoal of the constant for Eq. (2-20) and the constant for the ionization of water [Eq. (2-5)] :

If no other source of Cu++ or H+ is present in the solution, we may set (CuOH+) = (H+) and (Cu++) = 10-4 – ( H+), whence

And therefore An approximate solution is (H+) = 10-5.7 M ; hence the pH of 10-4M CuSO4 is about 5.7. For this kind of problem is preferable to use the stepwise ionization of the hydroxide [Eq. (2-20) rather than Eq. (2-22)]. If the hydrolysis had been set up to show the formation of Cu(OH)2 appears in copper sulfate solutions.

General Equation fo Hydrolysis From these examples come the generalization that any hydrolysis constant may be obtained by dividing the ionization constant of water (raised to a power if the coefficient of H+ or OH- is greater than 1) by the ionization constant of the weak acid or weak base that is formed in the hydrolysis reaction

One may generalize also tha simple solutions of any soluble sulfide or soluble carbonate will necessarily be alkaline and that simple solutions of salts of the common heavy metals [for example, FeCl3, Pb(NO3)2, NiSO4] will neccessariy be acidic, because of hydrolysis reactions analogous to eqs. (2-24) and (2-26). We shall find these generalizayions useful in discussing the many geologic process in which hydrolysis plays an important role.

2-8. ESTIMATING IONIC CONCENTRATIONS In a solution of given pH, if dissolved carbonate is known to be present, would it exist chiefly as H2CO3, HCO3̅, or CO3̅? If the solution contains ferric iron, is it mainly Fe3+, FeOH++, or Fe(OH)2+? Would dissolved zinc be present as positive ions (Zn++ and ZnOH+) or as negative ions [Zn(OH)3̅ and Zn(OH)4̿]? This is a kind of question often encounterd in geochemistry and easily answered if equilibrium constants are known. The distribution of carbonate species will serve as a convent example. We know to begin with, in a general qualitative way, that dissolved carbonate must exist chiefly as H2CO3 in acid solutions, as CO3̅ in basic solutions, and as HCO3̅ in some intermediate range. To fix the limits, we write the equations for the two ionization contants [Rqs. (2-9) and (2-10)] in the form

From the expressions, the concentrations of HCO3̅ and H2CO3 must be equal when (H+) has a numerical value equal to K1, and the concentrations of CO3̿ and HCO3̅ are equal when (H+) = K2. Hence we can say immediately that H2CO3 is the dominant carbonate species in all solutions with pH less than 6.4 [or (H+) greater than 10-6.4M], HCO3̅ is dominant in the pH range 6.4 to 10.3, and CO3̿ is dominant at pH’s above 10.3. These rules concentrated it may be or what other solutes may be present.

Sippose now that we have given also a total analytical concentration of carbonate, say 0.001M. In a solution whose pH is 6.4 the concentrations of both H2CO3 and H2CO3 must be half this number, or 0.0005M, and the concentration of CO3̿ and HCO3̅ are both 0.0005M and (H2CO3) is effectively 0.001M ; at pH’s well above 10.3, (CO3̿) must be 0.001M. It is often desirable to know the concentrations of all three carbonate species in a given solution, even though one or two may be very minor. For example, what are the concentrations of CO3̿ and H2CO3 in a solution containing 0.001M total dissolved carbonate and having a pH of 8.0? This is in the intermediate range, where most of the dissolved carbonate exist as HCO3̅, so that this ion may be assigned a concentration of approximately 0.001M. Then the equation for K2 [Eq. (2-10)] becomes

And the equation for K1[Eq. (2-9)]

If (H+) is 10-8M, these equations give (CO3̿) = 10-5.8M. and (H2CO3) = 10-4.6M. To generalize these result, it is convenient to rewrite the two equation s in logarithmic form:

These may be Simplifed to

If now log(CO3̿) is plotted against pH, it should give a straight line with unit positive slope, and log (H2CO3) should give a straight line with unit negative slope. These relations hold over the pH range in which HCO3̅ has the approximate concentration 0.001M, in other words from roughly 7.0 to 9.5. Similar equations can be set up for other pH ranges, giving combined plot,

Shown in fig. 2-2. The diagram is drawn for a total carbonate concentration of 0.001M., but it can be used for any desired concentrations of each of the three carbonate species at any pH can be read as intersections of the appropriate lines with a vertical line through the pH value. A similar diagram for sulfide solutions is given in Fig. 2-3. The two acids H2CO3 and H2S are the most important weak acid in geologic environments, and an understanding of the relations of the two acids and their ions, as summarized in Figs 2-2 and 2-3, is essential to geochemical work with natural solutions. Similar diagrams may be constructed for other weak acids, and similar reasoning may be applied to other solutes that can exist in different species.

2-9. CARBONATE EQILIBIRA

As another illustration of the use of equilibrium reasoning, we consider next the relations between carbonate acid and carbonate minerals. These relations determine the conditions under

Which carbonate rocks are formed or dissolved and likewise the conditions of formation of carbonate gangue minerals in veins. The discussion in this section will be limited to qualitative reasoning , and to a single carbonate—the carbonate of calcium, which is by far the most abundant carbonate in nature. In the next chapter we shall make the treatment more quantitative and extend it to other common carbonate minerals. A strong acid dissolves calcium carbonate by the familiar reaction

If the concentration of acid is low, a more accurate equation would be

Showing that H+ takes CO3̿ away from Ca++ to from the very weak (little ionized) acid HCO3̅. These reactions would take place in nature, for example, where strongly acid solutions from the weathering of pyrite encounter limestone. The reactions can be reserved by any process that uses up H+; for example, if a base is added

Quite evidently the solubility of CaCO3 is determined in large part by the pH of its environment. By determining equilibrium constants for the above reactions, we could express this dependence quantitavely, but first it will be useful to get a feeling for the qualitatively relationships. Under natural conditions the dissolving of calcium carbonate is a little more complicated, because the acids involved are usually weak rather than strong. When limestone dissolves in carbonic acid, for example, the overall process may be summarized by the equation

Note that the two HCO3̅ ions are from different sources: one is simply left over from the onization of H2CO3, and the other is formed by the reaction of H+ from the acid with CaCO3, as shown by Eq. (2-29). Equation (2-31) is the essential reaction for an understanding of carbonate behavior in nature. The forward reaction shows what happens when limestone weathers, when limestone is dissolved to form caves, or when marble is dissolved by ore-bearing solutions in the walls of a fissure. The reverse process represents the precipitation of calcium carbonate in the sea, as a cementing material in sedimentary rocks, or where droplets evaporate at the tip of a stalactite. The effect of pH on solubility is shown as well by Eq. (2-31) as by the simpler equations preciding it. At low pH, where most dissolved carbonate exist as H2CO3 ( Fig. 2-2), the forward reaction is favored; at high pH, the reverse reactiom leading to precipitatiom is favored, because OH- reacts preferentially with the stronger acid, H2CO3, rather than with the very weak HCO3̅. The equatiom shows also that the position of equilibrium (and hence the solubility) depends on the pressure of CO2 above the solution, since this pressure helps to determine the concentration of dissolved H2CO3:

Any process that increases the amount of CO2 available to the solution makes more CaCO3 dissolve; anything that decreases the amount of CO2 causes CaCO3 to precipitate. Some of the important naturalproccess that affect the solubility of CaCO3 by changing the position of the equilibrium Eq. (2-31) are described in the following paragraphs.

Temperature Changes The solubility of CaCO3 in pure water decreases somewhat as the temperature rises. This opposite to the behavior of most salts; the general result of increasing temperature is to give higher solubilities, but a number of carbonates and sulfates are exceptions. In additional to this effect, the solubility of CaCO3 in natural waters decreases at higher temperatures because CO2, like any other gas, is less soluble in hot water that in cold water . In general, the solubility of carbonates is much more influenced by this change in solubility of CO2 than by the temperature coefficient of the solubility itself. As an illustration of the effect of temperature, CaCO3 dissolves at great depths in the ocean, where the water is perennially cold, but precipitates near the surface, especially in the tropics, where the water is warm.

Changes in Pressure The effect of pressure by itself, independent of its effect on CO2, is to increase the solubility of CaCO2 slightly. Where pressure are very large, the effect of pressure may be substantial; in the deep parts of the ocean, for example, pressure in near-surface environments, however, is the change in amount of dissolved CO2 when the pressure of the gas changes in the surrounding atmosphere. Theoretically even day-to-day barometric changes should have a detectable effect on solubility, and the local production of CO2 in abnormal ammounts, say by a forest fire; an industruial plant, or a volcanic eruption, could cause a marked increase temporarily. But circulation of the atmosphere is so effective in keeping the partial pressure of CO2 uniform that this factor is probably less important than the others.

Organic Activity Many organism use calcium carbonate in the construction of their shells. Just how ther accomplish this is not certaim, but they flourish in greatest numbers in water approximately saturated with CaCO3, where only a minor chage in pH is needed