Soda ash is a white, anhydrous, powdered or granular material containing more than 99% sodium carbonate (Na2CO3) when shipped. The accepted commercial standard for soda ash is expressed in terms of the equivalent sodium oxide (Na2O) content. A 99.5% soda ash is equivalent to 58.2% Na2O (the conversion equation is: % Na2CO3x 0.585 = % Na2O).
Soda ash is an alkali that has a high pH in concentrated solutions. It can irritate the eyes, respiratory tract and skin. It should not be ingested because it can corrode the stomach lining.
Soda ash is made in three main grades — light, intermediate and dense. These differ only in physical characteristics, such as bulk density and particle size and shape (which affects flow characteristics and angle of repose). Other physical and chemical properties are common to all grades, whether in solid or liquid form. These are similar to those given for pure sodium carbonate in standard reference books and other sources from which much of the data that follow are derived, e.g., the properties in Table 2-1 and the densities, dissociation pressures and heats for formation, hydration and solution in Table 2-2 through Table 2-5.
Decomposition on HeatingAnhydrous sodium carbonate loses weight when heated due to dissociation and volatilization according to the following reaction:
Na2CO3 (solid) = Na2O (solid) + CO2 (gas). Dissociation pressure rises with increasing temperature (Table 2-3).
Hydrates of Sodium CarbonateSodium carbonate has three hydrate forms: Sodium carbonate monohydrate, heptahydrate and decahydrate.
Sodium carbonate monohydrate (Na2CO3•H2O) contains 85.48% Na2CO3 and 14.52% water of crystallization. It separates as small crystals from saturated aqueous solutions above 35.4°C (95.7°F). It can be formed by wetting soda ash with a calculated quantity of water at or above this temperature. It loses water on heating, and its solubility decreases slightly with increasing temperature. It converts to Na2CO3 upon contact with its saturated solution at 109°C (228°F).
Sodium carbonate heptahydrate (Na2CO3•7H2O) contains 45.7% Na2CO3 and 54.3% water of crystallization. It is of no commercial interest because its stability range only extends from 32.0°C to 35.4°C (89.6°F to 95.7°F).
Sodium carbonate decahydrate (Na2CO3•10H2O), commonly called "sal soda" or "washing soda," usually forms large, transparent crystals containing 37.06% Na2CO3 and 62.94% water of crystallization. It can be crystallized from saturated aqueous solutions between -2.1°C and 32.0°C (28.2°F and 89.6°F respectively) or by wetting soda ash with a calculated quantity of water in this temperature range. The crystals readily effloresce in dry air, forming a residue of lower hydrates (principally the monohydrate form).
Heat of SolutionHeat is released when anhydrous or monohydrate sodium carbonate dissolve in water. Heat is absorbed when the heptahydrate or decahydrate forms dissolve in water. The stronger the concentration, the greater the heat released or absorbed per unit of Na2CO3 dissolved. Use Fig. 2-1 to calculate the heat absorbed when diluting a sodium carbonate solution. For example, when a 25% solution is diluted to 10%, temperature decreases through the absorption of:
131.7 - 114.3 = 17.5 Btu/lb. Na2CO3.
When soda ash is dissolved in water to form a 32% saturated solution, 135 Btu/lb. of heat is released (Fig. 2-1). As additional soda ash is added to the saturated solution, monohydrate crystals are formed. This heat of formation is 54 Btu/lb (Table 2-4). If equal weights of soda ash and water are mixed, forming a 50% slurry, about 42% of the soda ash dissolves to form a saturated 32% solution. The remaining 58% soda ash forms monohydrate crystals. The total heat developed in preparing a 50% slurry is:
(0.42 x 134) + (0.58 x 54) = 88 Btu/lb. of soda ash.
When more water is added to the slurry, monohydrate dissolves to saturate the water. For example, when one gallon of water is added to soda ash slurry, about 4 pounds of soda ash will dissolve (4.7 lbs. of monohydrate). Subtracting the heat of formation from the heat of solution gives the net heat released by dissolving a saturated monohydrate slurry with 1 gallon of water:
4 x (134 - 54) = 282 Btu of heat.
SolubilitySodium carbonate, although readily soluble in water, is unusual in that it reaches maximum solubility at the relatively low temperature of 35.4°C (95.7°F). At this point, 100 parts of water dissolves 49.7 parts of Na2CO3 to yield a 33.2% solution by weight. Solubility decreases above this temperature, so there are two saturation temperatures for concentrations between about 29% and 33.2%. The phase diagram (Fig. 2-2) portrays this relationship by tracing solubility (as % Na2CO3) between -2.1 and 109°C (28.2 and 228.2°F).
Solubility data above about 105°C represent solutions held above atmospheric in order to prevent boiling, since the boiling point-concentration curve crosses the solubility curve at about 105°C. Unsaturated solutions exist in the area above and to the left of this curve. The region below and to the right of the curve contains either undissolved solids in contact with saturated solutions or solids alone (Table 2-6).
This diagram helps trace the effects of cooling, heating and changing concentrations. For example, a 20% Na2CO3 solution is unsaturated at all temperatures above 22.4°C (72.3°F), which is where the 20% concentration line crosses the saturation curve. Below this temperature, the solid phase (Na2CO310H2O) begins to form, increasing in amount as temperature falls. This phase change causes the concentration of the saturated solution in contact with the crystals to decrease, until at -2.1°C (28.2°F), the liquid phase disappears leaving only a mixture of solid Na2CO310H2O and ice.
In the above example, where a 20% Na2CO3 solution is cooled below -2.1°C so it solidifies, the richer component of the final mixture (Na2CO3 10H2O) has 37.0% Na2CO3 and the leaner component (ice) has 0% Na2CO3. The final mixture then contains:
Referring to the phase diagram (Fig. 2-2), a mixture of 40% Na2CO3 and water at 50°C contains Na2CO3 H2O crystals (85.5% Na2CO3) in equilibrium with its saturated solution (32% Na2CO3). The physical composition of the mixture is:
Cooling this mixture to 35°C changes the solid phase from Na2CO3 H2O to solid Na2CO3 7H2O containing 45.7% Na2CO3 in contact its saturated solution (33% Na2CO3). The mixture now consists of:
If the mixture is cooled below 32°C, it solidifies to a mixed solid, which contains Na2CO3 7H2O (45.7% Na2CO3) and Na2CO3 10H2O (37.0% Na2CO3) in the proportion of:
The solubility of soda ash in the presence of appreciable amounts of foreign salts, such as sodium chloride, sodium sulfate and sodium nitrate, changes how well the phase diagram in Fig. 2-2 applies. See the International Critical Tables or the Solubilities of Inorganic and Metal Organic Compounds by A Seidell, 4th Edition, 1958 (Vol. 1) and 1965 (Vol. 11) for the effects of salts on the solubility of sodium carbonate.
Solution Specific Gravity and DensitySoda ash has a solubility limit of 14.5% Na2CO3 at 15.6°C (60°F). Table 2-7 lists densities at 15.6°C and specific gravities for concentrations of sodium carbonate up to 14.0% (from the International Critical Tables5).
Higher concentrations are possible above 15.6°C. The specific gravities of saturated solutions above 15.6°C lie on a smooth curve (Fig. 2-3)6. Table 2-8 lists values at saturation for concentrations of 15% and above. Crystallization occurs when temperature falls below that shown for these solutions.
The specific gravity of sodium carbonate solutions decreases with increasing temperature. Concentration can be determined if solution specific gravity and temperature are known (Table 2-8 and Fig. 2-4).
Specific HeatFigure 2-5 shows the specific heat of sodium carbonate solutions at 20°C7 and at 30 and 76.6°C8. For example, the heat required to raise the temperature of 1000 gallons of 10% Na2CO3 solution from 68°F to 170°F is calculated from:
A x B x C = Q
where A = weight of solution = 9924 lb.
Vapor PressureTable 2-10 shows vapor pressures and boiling points for sodium carbonate solutions9 with values for saturated solutions in parentheses10. Approximate values for vapor pressures in the upper temperature ranges may be interpolated using Fig. 2-6 by extending a straight line from the temperature line through the % Na2CO3 line to the vapor pressure line. The example in Fig. 2-6 (dashed line) shows the vapor pressure of a 10% sodium carbonate solution at 90°C (194°F) to be 67,500 Pascals (506 mm of mercury).
ViscosityViscosity is important for designing pumping and piping systems and for calculating heat transmission and gas absorption in chemical processes. The viscosity of a soda ash solution is determined from Fig. 2-7 by extending a straight line from the % soda ash line through the temperature curve to the viscosity line. The example in Fig. 2-7 (dashed line) shows the viscosity of a 22% sodium carbonate solution of 24°C (75°F) to be approximately 4.0 centipoises (0.004 Pascal seconds).
Hydrogen Ion Concentration (pH)Precise and accurate determination of pH values for sodium carbonate solutions requires electrometric measurement with glass electrodes. These should be designed for use in alkaline sodium salt solutions and measure over the entire pH range. Organic color indicators are not recommended for pH measurement in sodium carbonate solutions (see Analytical Procedures section). The pH of sodium carbonate solutions can be interpolated from Fig. 2-813 by extending a straight line from the % Na2CO3 line through the reference point to the pH line as shown by the dashed line on the chart.
Electrical ConductivityElectrical conductivity instruments may be designed for the measurement and control of sodium carbonate concentrations, particularly when applied to concentrations in the lower ranges (Table 2-11).
Bulk DensityThe bulk density (weight of dry soda ash per unit volume) varies with the form of ash and the handling it receives. Typical bulk densities are shown in Table 2-12.
References1. Selected Values of Chemical Thermodynamic Properties,, U.S. Department of Commerce, National Bureau ofStandards, Circular 500, 1952, p 799.
2. National Research Council, International Critical Tables,, Vol. III, McGraw-Hill, New York, 1928, p. 24.
3. Howarth, J.T. and Turner, W.E.S., J. Soc. GlassTech.,, 14T, 394-401 (1930).
4. Ref.1, p.468.
5. Ref.2, Vol.111, p.82.
6. Roberts, L.D. and Mangold, G.B., Ind. Eng. Chem.31, 1293 (1939).
7. Ref.2, Vol. V, 1929, p.124.
8. Swallow, J.C. and Alty, S., J. Chem. Soc. (London)134, 3062 (1931).
9. Ref.2, Vol. 111, p. 372.
10. Landolt-Bornstein, Physlkalisch-chemischeTabellen,, 5th Ed., Vol. III, Springer, Berlin, 1936p 2497
11. Davis, D.S., Chemical Processing Nomographs,Chemical Publishing Company, Inc., New York,1960, p, 84.
12. Ibid., p. 167.
13. Lortie, L. and Demers, P., Can. J. Research, 18,160-167 (1940).
14. Ref.2, Vol. VI, 1929, pp. 248 & 254.
15. Kobe, K.A. and Carlson, C.J., J. Electrochem Soc.,101, 155-157 (1954).
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