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Electrolysis is one of the acknowledged means of generating chemical products from their native state. For example, metallic copper is produced by electrolyzing an aqueous solution of copper sulfate, prepared by leaching the copper bearing ores with sulfuric acid. Or, one can prepare chlorine gas and sodium hydroxide solution by electrolyzing an aqueous solution of sodium chloride, which exists in nature in a solid form as rock salt and also available as solar or vacuum evaporated salt. The solution of sodium chloride (common table salt) is often called "brine."
The primary products of electrolysis are chlorine gas, hydrogen gas, and sodium hydroxide solution (commonly called "caustic soda" or simply "caustic"). However, if the electrolyte is maintained at a pH of 6.5 or 10, one can form chlorate or hypochlorite from the electrogenerated chlorine and caustic. This is the basis for the electrolytic production of sodium chlorate or sodium hypochlorite (commonly known as "bleach").
The first observation of a possible application for chlorine was its bleaching effect on vegetable matter. In 1774, Carl Wilhelm investigated the reactivity of the greenish-yellow gas generated during the reaction involving the oxidation of hydrochloric acid by a manganese dioxide ore (pyrolusite). In 1785, Berthollet tried unsuccessfully to use elemental chlorine for textile bleaching to replace solar bleaching. Elemental chlorine caused discomfort to the workers, corroded metal parts, and softened the fabrics. The first use of chlorine in the form of potassium hypochlorite was for bleaching, and dates back to 1789. It was in 1808 that Davy characterized this greenish-yellow gas as an element and named it "chlorine."
The development of chemical bleaching with chlorine and the discovery of calcium hypochlorite bleaching powder as a practical mode of transporting chlorine was of great significance. These technologies made a marked impact on the textile bleaching operations in Great Britain and Europe, who were in the middle of the industrial revolution with expanding production, and hence, the demand for textiles. The invention of the power loom provided the capability to produce textiles on a large scale. However, solar bleaching, by spreading the cloth in open fields for months, became increasingly expensive in view of the soaring land values. The chlorine bleaching process not only shortened the operations from months to few days, but also freed vast areas of land for more productive use. Based on the greatly improved efficiency of textile bleaching, the pulp and paper industry also began using bleaching powder. Between 1756 and 1932, the use of chlorine in the pulp making industry increased. Chlorine, in the form of hypochlorites, removed the color or color producing materials from the cellulose fibers, without undue degradation of the fibers.
The first use of chlorine for disinfection dates back to 1823, when it was used in hospitals. Chlorine water was employed in obstetric wards to prevent puerperal fever in 1826, and fumigation with chlorine was practiced during the great European cholera epidemic. Following the discovery that bacteria were responsible for the transmission of certain diseases, several investigators studied chlorination of both sewage and potable water in 1890's in an attempt to destroy these bacteria. By 1912, the use of chlorine for water treatment had become a common practice. There was significant reduction in the incidence of water borne diseases, such as typhoid. For example, from October to December 1909, 549 cases of "winter typhoid" were reported in Montreal, Canada. After chlorination of drinking water was begun in 1910, only 170 cases were reported for the same 4-month period. Thus, virtually all the chlorine manufactured during the 19th century was consumed by these two industries. The major turning event for the growth of the chlorine industry was its use in 1912 for water purification during the Niagara Falls typhoid epidemic. It should be noted that bleaching powder was used in 1897 to clean the polluted mains during a typhoid break in England.
|Fig. 1. Chlorine end uses.|
Presently, the primary uses of chlorine are in the pulp and paper manufacturing operations for bleaching to produce a high quality whitened material and in water treatment operations as a disinfectant (Figure 1). The other uses of chlorine include the production of organic and inorganic chemicals. The largest volume organic chemical manufactured that involves chlorine is polyvinyl chloride (PVC). PVC is a very versatile thermoplastic, used in a wide variety of daily products. The major use of chlorine in the production of inorganic chemicals is for titanium dioxide (a widely used pigment), manufactured from naturally occurring ores (ilmenite or rutile).
|Fig. 2. Caustic soda end uses.|
The major use of caustic for making inorganic chemicals is in the production of hypochlorite for household and industrial bleaching purposes. Also, its use in the pulp and paper industry is in the production of sodium sulfide and sodium hydrosulfide for mechanical pulping. It is also used in the food processing applications, which include skin removal of potatoes, tomatoes etc, for further processing.
|Fig. 3. Chlorate end uses.|
The electrosynthesis of sodium chlorate dates back to 1802, when von Hisinger and Berzelius prepared sodium chlorate by the electrolysis of sodium chloride solution. The first chlorate cell patent was issued to Watt in 1851. The first chlorate plant was built in 1886 in Villers-St. Sepulchre in Switzerland, where chlorate was electrochemically produced in cells made of wood and equipped with a diaphragm. The energy consumption was about 15,000 kWh/ton potassium chlorate. This may be compared to an energy consumption for a crystal product of about 5,000- 6,000 kWh/ton with modern technologies.
Figure 3 describes the end-use profile of sodium chlorate in 1998. About 93% of sodium chlorate is used for production for bleaching in the pulp and paper industry. The remainder is utilized in the agricultural industry as a cotton defoliant or herbicide (weed killer), as an oxidizer in uranium milling, and in the production of ammonium perchlorate used in rocket propulsion. (It is worth noting that perchlorates are also produced by an electrolytic process, where chlorate is anodically oxidized to perchlorate.) These uses of sodium chlorate have remained unchanged over the past 20 years, although the relative demands have changed. World capacity of sodium chlorate was estimated as about 2.8 million short tons during 1998; the North American share was about 1.95 million tons.
|Fig. 4. Chlorine cell technology in the U.S.|
The chlor-alkali (also called "chlorine-caustic") industry is one of the largest electrochemical technologies in the world. It is an energy intensive process and is the second largest consumer of electricity (2400 billion kWh) among electrolytic industries. In 2006, about 84% of the total world chlorine capacity of about 59 million metric tons was produced electrolytically using diaphragm and membrane cells, while about 13% was made using mercury cells (Figure 4).
Chlorine is produced by the electrolysis of sodium chloride (common table salt) solution, often called "brine." Thus, when sodium chloride is dissolved in water, it dissociates into sodium cations and chloride anions. The chloride ions are oxidized at the anode to form chlorine gas and water molecules are reduced at the cathode to form hydroxyl anions and hydrogen gas. The sodium ions in the solution and the hydroxyl ions produced at the cathode constitute the components of sodium hydroxide formed during the electrolysis of sodium chloride. (The chemical reactions occurring in the cell are shown in the Appendix.)
Chlorine is produced electrolytically using three types of electrolytic cells. The main difference in these technologies lies in the manner by which the chlorine gas and the sodium hydroxide are prevented from mixing with each other to ensure generation of pure products. Thus, in diaphragm cells, brine from the anode compartment flows through the separator to the cathode compartment, the separator material being either asbestos or polymer-modified asbestos composite deposited on a foraminous cathode. In membrane cells, on the other hand, an ion-exchange membrane is used as a separator. Anolyte-catholyte separation is achieved in the diaphragm and membrane cells using separators and ion-exchange membranes, respectively, whereas mercury cells contain no diaphragm or membrane and the mercury itself acts as a separator. The anode in all technologies is titanium metal coated with an electrocatalytic layer of mixed oxides. All modern cells (since the 1970's) use these so-called “dimensionally stable anodes" (DSA). Earlier cells used carbon based anodes. The cathode is typically steel in diaphragm cells, nickel in membrane cells, and mercury in mercury cells. These cell technologies are schematically depicted in Figures 5-7 and are described below.
|Fig. 5. Schematic of a mercury cell.|
The mercury cell has steel bottoms with rubber-coated steel sides, as well as end boxes for brine and mercury feed and exit streams with a flexible rubber or rubber-coated steel cover. Adjustable metal anodes hang from the top, and mercury (which forms the cathode of the cell) flows on the inclined bottom. The current flows from the steel bottom to the flowing mercury.
Saturated brine fed from the end box is electrolyzed at the anode to produce the chlorine gas, which flows from the top portion of the trough and then exits. The sodium ion generated reacts with the mercury to form sodium amalgam (an alloy of mercury and sodium), which flows out of the end box to a vertical cylindrical tank. About 0.25% to 0.5% sodium amalgam is produced in the cell. The sodium amalgam reacts with water in the decomposer, packed with graphite particles and produces caustic soda and hydrogen. Hydrogen, saturated with water vapor, exits from the top along with the mercury vapors. The caustic soda then flows out of the decomposer as 50% caustic. The unreacted brine flows out of the exit end box. Some cells are designed with chlorine and anolyte outlets from the end box, which are separated in the depleted brine tank. The mercury from the decomposer is pumped back to the cell.
|Fig. 6. Schematic of a diaphragm cell.|
The diaphragm cell is a rectangular box with metal anodes supported from the bottom with copper-base plates, which carries a positive current. The cathodes are metal screens or punch plates connected from one end to the other end of the rectangular tank. Asbestos, dispersed as a slurry in a bath, is vacuum deposited onto the cathodes, forming a diaphragm. Saturated brine enters the anode compartment and the chlorine gas liberated at the anode during electrolysis, exits from the anode compartment. It is saturated with water vapor at a partial pressure of water over the anolyte. The sodium ions are transported from the anode compartment to the cathode compartment, by the flow of the solution and by electromigration, where they combine with the hydroxyl ions generated at the cathode during the formation of the hydrogen from the water molecules. The diaphragm resists the back migration of the hydroxyl ions, which would otherwise react with the chlorine in the anode compartment. In the cathode compartment, the concentration of the sodium hydroxide is ~12%, and the salt concentration is ~14%. There is also some sodium chlorate formed in the anode compartment, dependent upon the pH of the anolyte.
|Fig. 7. Schematic of a membrane cell.|
In a membrane cell, an ion-exchange membrane separates the anode and cathode compartments. The separator is generally a bi-layer membrane made of perfluorocarboxylic and perfluorosulfonic acid-based films, sandwiched between the anode and the cathode. The saturated brine is fed to the anode compartment where chlorine is liberated at the anode, and the sodium ion migrates to the cathode compartment. Unlike in the diaphragm cells, only the sodium ions and some water migrate through the membrane. The unreacted sodium chloride and other inert ions remain in the anolyte. About 30-32% caustic soda is fed to the cathode compartment, where sodium ions react with hydroxyl ions produced during the course of the hydrogen gas evolution from the water molecules. This forms caustic, which increases the concentration of caustic solution to ~35%. The hydrogen gas, saturated with water, exits from the catholyte compartment. Only part of the caustic soda product is withdrawn from the cathode compartment. The remaining caustic is diluted to ~32% and returned to the cathode compartment.
Thus, all three basic cell technologies generate chlorine at the anode, and hydrogen along with sodium hydroxide (caustic soda) in the cathode compartment (or in a separate reactor for mercury cells, see Figure 5). The distinguishing difference between the technologies lies in the manner by which the anolyte and the catholyte streams are prevented from mixing with each other. Separation is achieved in a diaphragm cell by a separator, and in a membrane cell by an ion-exchange membrane. In mercury cells, the cathode itself acts as a separator by forming an alloy of sodium and mercury (sodium amalgam) which is subsequently reacted with water to form sodium hydroxide and hydrogen in a separate reactor.
A comparison of the performance characteristics of these three technologies is presented in the Appendix together with schematic process diagram for each of the cell technologies. The primary technology that is presently being used for future expansions or replacements of existing circuits is the membrane cell technology. The major membrane cell technology suppliers, include: Uhde GmbH, Asahi Chemicals, and Chlorine Engineers. DeNora Tech is the sole supplier of diaphragm cell technology. It is highly unlikely that anyone will build a new mercury- or diaphragm-cell plant in the future. Figures 8 and 9 illustrate cell rooms with diaphragm and membrane chlor-alkali cells.
|Fig. 8. Chlor-alkali cell room with MDC-55 diaphragm cells (Courtesy of Occidental Chemical Corporation).|
|Fig. 9. Chlor-alkali cell room with BL-2.7 membrane cells (Courtesy of Uhde GmbH).|
Chlorine is first cooled to 60oF (16oC) and passed through demisters to remove the water droplets and the particulates of salt and sodium sulfate. The cooled gas goes to sulfuric acid circulating towers, which are operated in series. Commonly, three towers are used for the removal of moisture. The dried chlorine then goes through demisters before it is compressed and liquefied at low temperatures. The non-condensed gas, called snift gas, is used for producing hypochlorite or hydrochloric acid. If there is no market for hydrochloric acid, the snift gas is neutralized with caustic soda or lime (calcium hydroxide) to form hypochlorite. The hypochlorite is either sold as bleach or decomposed to form salt and oxygen.
The membrane cell caustic is concentrated in a multiple effect falling film evaporator, which increases the caustic soda concentration to 50% with a high steam economy. Caustic soda from membrane cells generally has 30-ppm sodium chloride and 5-10 ppm sodium chlorate.
The catholyte from the diaphragm cells contains ~12% sodium hydroxide, ~14% sodium chloride, 0.25%-0.3% sodium sulfate, and 100-500 ppm sodium chlorate. The catholyte is evaporated in a multi-effect evaporator. Most of the salt from the catholyte will precipitate during the concentration of the caustic soda to 50% sodium hydroxide. The 50% caustic soda product will contain about 1% sodium chloride. The 50% caustic also has a high chlorate concentration (~0.1%) compared to the caustic from membrane or mercury cells (~10 ppm). The salt, separated from the caustic during evaporation, is used to re-saturate the brine fed to the cell.
An additional single-effect evaporator is needed to produce 73% caustic soda. Anhydrous (dry) caustic soda is produced in a rising film evaporator, operating at 725oF (385oC) and at a few inches (one inch =2.54 cm) of water vacuum.
The solution-mined brine or the solid salt dissolved in the salt dissolver is treated in a reactor with sodium carbonate and caustic soda to precipitate calcium carbonate and magnesium hydroxide (see the Appendix). These precipitates are settled in a settler. The underflow carries the solid slurry, which is pumped to a filter to remove it as sludge, or sometimes, it is disposed off along with the rest of the liquid effluent from the plant. The calcium carbonate precipitates are heavy, and drag with it the hydroxides of aluminum, magnesium, strontium, etc. The overflow from the settler, which carries ~10-50 ppm of suspended solids, is filtered. For the mercury and the diaphragm cell process, this brine is adequate, and can be fed to the electrolyzers.
In the all cell processes, the filtered brine is heated and passed through a bed of salt in a saturator in order to increase the salt concentration before feeding it to the electrolyzers. In some plants, the brine feed is acidified to improve the cell current efficiency. The acidification reduces the alkalinity, which would otherwise react with the chlorine in the anolyte compartment forming chlorate.
The membrane cell process brine specifications are more stringent than that of the mercury and diaphragm processes, and calls for impurities to be at the parts-per-billion (ppb) level. This is accomplished by filtering the brine in a pre-coat type secondary filter. An ion-exchange resin is used to remove the calcium, magnesium, barium, and iron impurities. It is also possible to remove aluminum by ion exchange. Often, aluminum and silica are removed by adding magnesium chloride in the brine exiting from the salt dissolver.
The depleted brine from the membrane and mercury cell processes carries dissolved chlorine. This brine is acidified to reduce the chlorine solubility, and then dechlorinated in a vacuum brine dechlorinator. The dechlorinated brine is returned to the brine wells for solution mining or to the salt dissolver. If the membrane and diaphragm processes coexist at a given location, the dechlorinated brine can be sent to a saturator for resaturation before being sent to the diaphragm cells.
The major raw material is sodium chloride, either very pure, such as solar rock salt, or partially purified evaporated salt. The salt is stored and dissolved in lixiviators to produce a saturated sodium chloride solution. This solution is purified by removing calcium, magnesium, fluoride, sulfate, and iron as insoluble compounds, through the addition of sodium carbonate, sodium hydroxide, sodium phosphate, and barium chloride.
The impurities or precipitates are removed in a pressure leaf filter with diatomaceous earth as a filter precoat and filter aid. This filter cake, containing approximately 35% water, is the only solid waste stream from the process. A polishing filtration stage and an ion-exchange system follow pressure leaf filtration.
The chemistry and electrochemistry of chlorate formation dictates that an efficient and economical cell should embody several distinct zones. In the electrolysis zone, the electrolytic reactions take place along with the hydrolysis of chlorine. As the chemical chlorate formation proceeds very slowly, a relatively large volume of chemical reaction zone is needed. A cooling zone is also required to remove the excess heat generated from the reaction and control the operating temperature. The cooling zone may be located within the chemical reactor or in an external heat exchanger. Hydrogen gas generated at the cathode must be released from the cell liquor. This hydrogen release takes place in the electrolysis cell, a separate vessel, or the chemical reactor.
A continuous stream of cell liquor flows from the electrolysis system to the "hypo removal" system, where the sodium hypochlorite concentration is reduced to low levels simply by heating the cell liquor to about 185-200oF (85-95oC) under careful pH control. Final traces of hypochlorite can be completely removed by treatment with a reducing agent (such as sodium sulfite or hydrogen peroxide).
|Fig. 10. Chlorate cell room with M25 Chemetics cells (Courtesy of Aker Chemetics).|
Sodium chlorate is usually recovered from cell liquors by concentration, followed by cooling to facilitate crystallization. Hot cell liquor, following hypo removal, is fed continuously into the circulation leg of a draft tube baffle evaporator/crystallizer. Crystal slurry is withdrawn from the bottom of the crystallizer section. The crystals are separated from the mother liquor and washed with water in a pusher centrifuge. They are thoroughly washed to remove sodium dichromate (an additive to the cell solution to increase current efficiency) from the chlorate crystals. Sodium dichromate contains chromium in the hexavalent state, which is a recognized human carcinogen. A white sodium chlorate crystal, containing about 1 to 1.5% moisture, is obtained from the centrifuge. Mother liquor from the centrifuge is mixed with fresh purified brine and recycled to the electrolytic cells.
Approximately 98% of the sodium chlorate capacity in North America is produced directly in sodium chlorate cells. The remaining 2% is produced "chemically" by the reaction of chlorine and caustic (see the Appendix for details).
In recent years, sodium chlorate technology sales have been dominated by the following three suppliers: Technip in France, Chemetics International in Canada, and Huron Technology in Canada. There are many other sodium chlorate technologies in operation, such as DeNora, Eltech, OxyChem, Oulu Oy, and Atochem. None of these are considered to be available for license. Figure 10 depicts a chlorate manufacturing facility with M25 Chemetics cells.
There are several manufacturers of seawater electrolysis cells in the market. The best known cells include Seachlor made by DeNora (producing 1000-2500 ppm active chlorine) or Salinec made by Exceltec International Corporation (generating 200-300 ppm active chlorine).
|Fig. 11. U.S. chlorine and caustic soda capacity and production.|
During the 1950's and 1960's, the demand for chlorine grew at a rate of 8%/yr and the plants were operating at greater than 90% capacity. The demand was for chlorinated derivatives and intermediates such as pesticides (DDT) used in agriculture and solvents, mainly chlorinated ethanes, which replaced flammable hydrocarbons in many cleaning and degreasing applications. Use of chlorinated methanes as intermediates increased in the manufacture of organosilicones, in tetramethyl lead gasoline additives and for fluorocarbons used as aerosol propellants and refrigerants. Polyvinyl chloride (PVC) plastics grew by 14%/yr as did ethylene and propylene oxide, which were produced by processes using chlorinated intermediates.
In 1970's, chlorine growth slowed to 2%/yr because of environmental concerns bringing restrictions on the use of pesticides such as DDT, kepone, dieldrin, and endrin. Carcinogenic characteristics of trichloroethylene, polychlorinated biphenyls (PCBs) and similar compounds also contributed to the declined growth. In 1978, the U.S. Environmental Protection Agency banned the use of fluorocarbon propellants for aerosols because of fears related to depletion of the ozone in the upper atmosphere. The passage of clean water legislation also had an adverse impact on a variety of end-uses. Thus, the paper industry started implementing changes in bleaching technology by increased use of sodium chlorate, oxygen, and hydrogen peroxide as replacements for chlorine bleaching. During this period, many chemical processes that used chlorine, particularly ethylene oxide and propylene oxide, were also converted to non-chlorine consuming processes.
Chlor-alkali producers, ignoring the potential impact of new non-chlorine based technologies and the various environmental concerns, continuing to project growth rates of 4-6%/yr. These projections are based on chlorine demand from exports, particularly to the Far East. Anticipating significant growth in exports, 15,000 tons/day of new capacity was added through the early 1980's.
In the 1980's, the environmental constraints impacted the downstream use of chlorine and operating costs increased because of the energy crisis or the cost of electricity. In addition, the exports declined because of the new ethylene dichloride (EDC) plants coming on stream overseas. As a result, the demand declined and the industry operated at only a 64% capacity. Overcapacity, slow growth, and high energy costs forced chlor-alkali producers to mothball or put on a standby mode a large number of production facilities, accounting for about 1.2 million metric tons (MT). By the end of 2006, restructuring decreased the operating capacity to 13 million MT/yr at an effective operating rate of ~89%.
Thus, the major factors that influenced the chlor-alkali industry are the environmental issues related to the use of products such as DDT or aerosols and the development of non-chlorine based technologies. There is yet another problem that confronts chlor-alkali producers, that is, the out-of-phase demand for chlorine and caustic soda. Chlorine markets follow the economy closely, since a large portion of the PVC market (its largest application) is in the housing and automotive industries.
Caustic soda, on the other hand, does not respond as readily to economic changes because of the diverse nature of its markets, such as pulp and paper and chemical processing. Another advantage for caustic soda is that it can more easily be stored which helps flatten out variable demand. These fluctuating demands for chlorine and caustic soda, resulting from the overall changes in the economy, generally lead to production cut backs and increased prices for either chlorine or caustic soda.
In the late 1980's, the chlorine industry recovered from earlier declines in consumption and enjoyed banner production years. In 1987 and 1988, annual increases reached 4 to 5% due to the strong economy. This was characterized by the increased demand for PVC and pulp and paper products and by increased exports of chlorine derivatives.
It should be noted here that most of the chlorine is traded globally as EDC, vinyl chloride monomer (VCM) and primary forms of PVC and that very little in its elemental form. The U.S. alone accounted for almost 50% of this trade in 1992. It is because of this market that chlor-alkali has seen moderate growth rate of 1-2% through 1990's. The world demand for chlorine is projected to grow at an annual rate of 2% through the year 2011. The annual capacity of chlorine will grow to 54 million MT in 2010 from 45 million MT in 2001. Low cost energy regions, such as Middle East, are projected to have higher annual growth rates of 3 to 3.5%.
Since 1990's, the use of elemental chlorine free (ECF) pulp bleaching involving chlorine dioxide from sodium chlorate and hydrogen peroxide has grown from 3% of the total bleached pulp production to ~55% in 1997. In contrast, the total chlorine-free (TCF) bleaching employing a combination of ozone, hydrogen peroxide, and oxygen has decreased ~18% in 1997 from ~40% in 1990. It may be noted that TCF uses greater amounts of hydrogen peroxide compared to that used in ECF technology. Although the current pulp inventories have declined in 1998, the final adoption of the Environmental Protection Agency's cluster rules in April 1998 allowing the substitution of chlorine dioxide for chlorine over total chlorine-free bleaching, has spurred the demand for sodium chlorate which is projected to grow at a rate of 5% per year through the next ten years.
Even with all these constraints, the chlor-alkali industry is projected to grow at a rate of 1 to 3% depending on pessimistic or optimistic reasoning. Much of this will be dictated on how effectively the industry responds to the concerns of the environmentalists and the government agencies.
|||2Cl- ==> Cl2 + 2e-||(anodic reaction)|
|||2Na+ + 2Hg + 2e- ==> 2Na (in Hg)||(cathodic reaction)|
|||2Cl- + 2Na+ + 2Hg ==> Cl2 + 2Na (in Hg)||(overall cell reaction)|
|||2Na (in Hg) + 2H2O ==> H2 +2NaOH + Hg||(decomposer reaction)|
|||2NaCl + 2H2O ==> Cl2 +2NaOH + H2||(overall process reaction)|
|||2Cl- ==> Cl2 + 2e-||(anodic reaction)|
|||2H2O + 2e- ==> 2OH- + H2||(cathodic reaction)|
|||2Cl- + 2H2O ==> Cl2 + H2 + 2OH-||(overall ionic reaction)|
|||2NaCl + 2H2O ==> Cl2 +2NaOH + H2||(overall reaction)|
|||Cl2 + 2NaOH ==> NaOCl + NaCl + H2O||(side reaction)|
|||3NaOCl ==> NaClO3 + 2NaCl||(side reaction)|
Reaction  will contaminate the caustic product with chlorate.
|||CaSO4 + Na2CO3 ==> CaCO3 +NaSO4||(CaCO3 precipitates)|
|||MgCl2 + 2NaOH ==> Mg(OH)2 + 2NaCl||(Mg(OH)2 precipitates)|
|Operating current density ( kA/m2)||8 - 13||0.9 - 2.6||3 - 5|
|Cell voltage (V)||3.9 - 4.2||2.9 - 3.5||3.0 - 3.6|
|NaOH strength (wt%)||50||12||33-35|
|Energy consumption ( kWh/MT Cl2) at a current density of (kA/m2)||3360 (10)||2720 (1.7)||2650 (5)|
|Steam consumption (kWh/MT Cl2) for concentration to 50% NaOH||0||610||180|
|Fig. 12. Mercury cell process flow sheet.|
|Fig. 13. Diaphragm cell process flow sheet.|
|Fig. 14. Membrane cell process flow sheet.|
Sodium hypochlorite/chlorate manufacturing process
|||2Cl- ==> Cl2 + 2e-||(anodic reaction)|
|||2H2O + 2e- ==> 2OH- + H2||(cathodic reaction)|
|||Cl2 + 2OH- ==> OCl- + Cl- + H2O||(hypochlorite formation)|
|||3OCl- ==> ClO3- + 2Cl-||(chlorate formation)|
|||NaCl + H2O ==> NaOCl + H2||(overall hypochlorite reaction)|
|||NaCl + 3H2O ==> NaClO3 + 3H2||(overall chlorate reaction)|
|||3Cl2 + 6NaOH ==> NaClO3 + 5NaCl + 3H2O||(chemical chlorate formation)|
Hypochlorite formation is promoted by the use of weak brine, basic solution, and low cell temperatures.
Chlorate formation is promoted by the use of saturated brine, acidic solution, and temperatures close to the boiling point of the solution.
|Fig. 15. Chlorate cell process flow sheet.|
Current density distribution in electrochemical cells
Extracting metals from sulfide ores
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