Tartaric acid (TA) is a byproduct of wine production. This organic acid and its salts are used in foods such as fruit jellies, preserves, jams, baked goods, and confections. TA is hardly metabolized and degraded by yeast and spoilage bacteria providing microbiological stability to foods that contain it. In addition to its use as antimicrobial and acidulant, TA and its salts are used as emulsifiers, leavening, and anticacking agents. TA is also used in the beverage industry and has non-food uses in textile coloring, galvanizing, and mirror production.
The increasing popularity in wine consumption in recent years has resulted in the increase of waste from wine making practices. One liter of white wine generates the same amount of water pollution as 3 people in one day. Waste-waters from wine production contain biodegradable compounds and fruit suspended solids;  their treatment is of great importance because their high pollutant activity. Moreover, waste treatment is of economic interest because the organic compounds present in waste from the wine making process can have value as additives, ingredients, and substrates in the food and pharmaceutical industries.
|Type of Waste||Name||Derived from||Treatment|
|Anaerobic digestion, ozonation, thermopilic anaerobic digestion, aerobic biodegradation, sequencing batch reactor, electrodyalisis, and wet oxidation|
|Skin, stalks, and seeds||Combustion, solid-state fermentation, incineration, composting, and pyrolysis.|
Tartaric Acid separated from grape juice
Tartaric Acid crystalsThe table above shows sources of waste in wine making as well as methods to treat them. A method devised by Rivas, et al. (2006) to treat distilled lees involves the reaction of tartaric acid with calcium-ions to form calcium tartrate. This salt of limited solubility is then redissolved with HCL to obtain TA. Upon removal of TA, the distilled lees can be used as nutrients for lactic acid bacteria (Lactobacilllus pentosus) for production of lactic acid.
In the method described above, addition of calcium ions prompts the precipitation of calcium tartrate. Once tartaric acid is removed from the solution as tartrate, it will partially dissociate into calcium and tartrate ions establishing the following equilibrium with the solid salt:with the corresponding equilibrium constant If we designate as x the concentration of each of the ions, the concentration of calcium-ions at equilibrium is From which The concentration of calcium-ions at equilibrium is 8.77 x 10–4 mol dm–3. If now we increase the concentration of tartrate-ions, the equilibrium will shift to the left according to the Le Chatelier’s principle. More calcium tartrate will precipitate, decreasing the concentration of calcium ions. The decrease in concentration obtained in this way is often referred to as the common-ion effect. Similarly, if an excess of calcium-ions is added to the solution, the concentration of tartrate-ion will decrease. Since in this process we are concerned about removing as much tartaric acid in the form of tartrate as possible, addition of calcium ions in excess will minimize dissociation of calcium tartrate increasing its the yield. The solubility product can be used to calculate how much the calcium-ion concentration is decreased by the common-ion effect. Suppose we mix 10 cm3 of a saturated solution of calcium tartrate with 10 cm3 of concentrated sodium tartrate (4 M C4H6O6). Because of the twofold dilution, the concentration of tartrate will be 2 mol dm–3. Feeding this value into equation 7 from the solubility product section, we then have the result or so that We have thus lowered the calcium-ion concentration from an initial value of 8.77 x 10–4 mol dm–3 ) to a final value of 3.85 × 10–7 mol dm–3, a decrease of about a factor of 2000!.
(1)In this case Q has a value below the solubility product, 7.7 × 10–7 mol2 dm–6. In order for equilibrium between the ions and a precipitate to be established, either the calcium-ion concentration or the tartrate-ion concentration or both must be increased until the value of Q is exactly equal to the value of the solubility product. The opposite situation, in which Q is larger than Ksp, corresponds to concentrations which are too large for the solution to be at equilibrium. When this is the case, precipitation occurs, lowering the concentration of both the lead and chloride ions, until Q is exactly equal to the solubility product. To determine in the general case whether a precipitate will form, we set up an ion-product expression Q which has the same form as the solubility product, except that the stoichiometric concentrations rather than the equilibrium concentrations are used. Then if
while if no precipitation occurs
EXAMPLE 1 Decide whether CaC2O4, calcium oxalate, will precipitate or not when (a)100 cm3 of 0.02 M CaCl2 and 100 cm3 of 0.02 M Na2C2O4 are mixed, and also when (b) 100 cm3 of 0.0001 M CaCl2 and 1000 cm3 of 0.0001 M Na2C2O4 are mixed. Ksp = 2.32 × 10–9 mol2 dm–6.
Solutiona) After mixing, the concentration of each species is halved. We thus have so that the ion-product Q is given by or
Since Q is larger than Ksp (2.32 × 10–9 mol2 dm–6), precipitation will occur.b) In the second case the concentration of each ion becomes and
thusSince Q is now less than Ksp, no precipitation will occur.
EXAMPLE 2 Calculate the mass of CaC2O4 precipitated when 100 cm3 of 0.0200 M CaCl2 and 100 cm3 of 0.0200 M Na2C2O4 are mixed together.
Solution In in part a of the previous example we determined that precipitation does actually occur. In order to find how much calcium oxalate is precipitated, we must concentrate on the amount of each species. Since 100 cm3 of 0.02 M CaCl2 was used, we have
similarly If we designate the amount of CaC2O4 that precipitates as x mmoles, we can set up the following table
|Species||Ca2+ (aq)||C2O42– (aq)|
|Initial amount (mmol)|
|Amount reacted (mmol)|
|Equilibrium amount (mmol)|
|Equilibrium concentration (mmol cm–3)|
or Rearranging, or so that Since 1.904 mmol of CaC2O4 precipitated, its mass is Note: Since the solubility product of CaC2O4 is very small, about 95% of the calcium oxalate originally formed precipitates.