Chloralkali process
The chloralkali process is an industrial electrolytic method that decomposes aqueous sodium chloride (brine) into chlorine gas at the anode, hydrogen gas at the cathode, and sodium hydroxide (caustic soda) as a co-product.[1][2] The process, which produces chlorine and sodium hydroxide in roughly equivalent molar amounts, underpins the manufacture of essential chemicals for water treatment, plastics like polyvinyl chloride, pulp and paper production, and soaps.[3] Commercialized in the late 19th century, it employs three primary cell technologies—mercury cathode, diaphragm, and membrane cells—with the latter two avoiding direct mercury use.[3][4] While highly efficient for bulk production, the mercury cell variant, once dominant, caused widespread environmental contamination through mercury emissions and losses into water and soil, prompting global phase-outs in favor of cleaner membrane technology since the 1970s.[4][5] Membrane cells now predominate due to superior energy efficiency, product purity, and minimal hazardous waste, though legacy mercury pollution from older facilities persists in some ecosystems.[5][6]Fundamentals of the Process
Chemical and Electrochemical Principles
The chloralkali process relies on the electrolysis of saturated aqueous sodium chloride solution, known as brine, to simultaneously produce chlorine gas, sodium hydroxide, and hydrogen gas. In this electrolytic decomposition, an electric current drives the non-spontaneous redox reactions within an electrochemical cell, where chloride ions are oxidized at the anode and water is reduced at the cathode. The process exploits the selective ion transport in brine, with sodium ions migrating toward the cathode to balance the generated hydroxide ions, forming sodium hydroxide without direct production of sodium metal.[7][8] At the anode, the primary reaction is the oxidation of chloride ions: $2Cl^- \rightarrow Cl_2 + 2e^-, which occurs at a standard electrode potential of approximately +1.36 V versus the standard hydrogen electrode. This chlorine evolution reaction (CER) predominates over oxygen evolution from water due to the high chloride concentration in brine, typically 300 g/L NaCl, which shifts the thermodynamics in favor of Cl₂ production. At the cathode, water reduction prevails over sodium ion reduction because of the negative standard potential for Na⁺/Na (-2.71 V), yielding: $2H_2O + 2e^- \rightarrow H_2 + 2OH^-, with a potential of -0.83 V at pH 14. The net cell reaction is thus $2NaCl + 2H_2O \rightarrow 2NaOH + Cl_2 + H_2, requiring a theoretical minimum voltage of about 2.19 V under standard conditions.[9][10][3] Thermodynamically, the process is governed by the Gibbs free energy change, ΔG° ≈ +237 kJ/mol for the overall reaction per mole of Cl₂, corresponding to a minimum energy input of 1654 kWh per metric ton of Cl₂ produced at 25°C. In practice, actual cell voltages range from 3.0 to 4.5 V due to overpotentials—particularly the high anodic overpotential for CER on dimensionally stable anodes (typically RuO₂-IrO₂ coated titanium, 0.2-0.4 V)—cathodic hydrogen evolution overpotentials (0.1-0.2 V on nickel cathodes), and ohmic losses from electrolyte resistance and separators (0.5-1.0 V). These inefficiencies arise from kinetic barriers in multi-step electron transfer and bubble formation, which increase resistance and reduce current efficiency to 90-95% for Cl₂.[10][11][12] The electrochemical principles also involve Faraday's laws, where the theoretical yield is 1.128 kg Cl₂ per kAh passed, but side reactions like hypochlorite formation (from Cl₂ reacting with OH⁻) or oxygen evolution reduce efficiency unless mitigated by cell design and operating conditions such as temperature (80-90°C) and current density (2-4 kA/m²). Ion-selective barriers prevent mixing of anolyte and catholyte, ensuring product purity: anode compartment yields >99% Cl₂, while cathode yields 30-50% NaOH solution.[9][8]Inputs, Outputs, and Stoichiometry
The chloralkali process requires as primary inputs a purified aqueous solution of sodium chloride, commonly termed brine, with a typical concentration of 300 grams of NaCl per liter of solution, alongside deionized water to maintain electrolyte balance and direct electrical current supplied at voltages of 3 to 4.5 volts per cell depending on the technology employed.[13] The brine serves as the source of chloride ions for oxidation at the anode, while water provides the protons and hydroxide ions involved in the cathodic reaction.[3] The principal outputs are chlorine gas (Cl₂) generated at the anode, hydrogen gas (H₂) evolved at the cathode, and an aqueous solution of sodium hydroxide (NaOH), also known as caustic soda, produced in the catholyte compartment.[13] In modern membrane cell operations, the NaOH output achieves concentrations up to 33% by weight, with chlorine gas purity exceeding 99.5% after drying and compression. Hydrogen gas is typically collected at over 99% purity and utilized as a fuel source or feedstock in other processes.[3] Stoichiometrically, the process adheres to the overall balanced equation $2NaCl + 2H_2O \rightarrow Cl_2 + H_2 + 2NaOH, derived from the anodic oxidation of chloride ions ($2Cl^- \rightarrow Cl_2 + 2e^-) and cathodic reduction of water ($2H_2O + 2e^- \rightarrow H_2 + 2OH^-), with sodium ions migrating to balance the catholyte.[13] This reaction requires theoretically two moles of electrons per mole of chlorine produced, corresponding to Faraday's laws of electrolysis, where one faraday (96,485 coulombs) liberates one equivalent of product; practical current efficiencies range from 90% to 95% due to minor side reactions like oxygen evolution at the anode.[3] The molar ratio of outputs is 1:1:2 for Cl₂:H₂:NaOH, ensuring balanced production when operating at theoretical conditions.[14]Historical Development
Early Discoveries and Non-Electrolytic Methods
The discovery of chlorine occurred in 1774 when Swedish chemist Carl Wilhelm Scheele produced the gas by reacting hydrochloric acid with manganese dioxide (pyrolusite).[15] Scheele's greenish-yellow gas was later recognized as a distinct element in 1810 by Humphry Davy, who named it "chlorine" from the Greek word for greenish-yellow.[16] In 1785, French chemist Claude-Louis Berthollet demonstrated chlorine's bleaching properties by dissolving it in alkaline solutions to form hypochlorites, enabling early applications in textile whitening without electrolytic means.[17] Industrial-scale chlorine production initially relied on non-electrolytic oxidation of hydrochloric acid, a byproduct of the Leblanc process for soda ash. The Leblanc process, patented by Nicolas Leblanc in 1791, converted sodium chloride and sulfuric acid into sodium carbonate (soda ash) via intermediate sodium sulfate, generating hydrochloric acid as waste: $2 \text{NaCl} + \text{H}_2\text{SO}_4 \rightarrow \text{Na}_2\text{SO}_4 + 2 \text{HCl}, followed by reduction with carbon and calcium carbonate to yield \text{Na}_2\text{CO}_3.[18] This acid was oxidized using manganese dioxide: \text{MnO}_2 + 4 \text{HCl} \rightarrow \text{MnCl}_2 + \text{Cl}_2 + 2 \text{H}_2\text{O}, but the process was inefficient due to manganese loss until Walter Weldon's 1866 improvement recycled manganese chloride back to dioxide via lime treatment, reducing costs and enabling wider adoption for bleaching powder production.[19] Further advancement came with Henry Deacon's 1868 process, which catalytically oxidized hydrochloric acid with atmospheric oxygen at approximately 450°C using cupric chloride: $4 \text{HCl} + \text{O}_2 \rightarrow 2 \text{Cl}_2 + 2 \text{H}_2\text{O}, bypassing manganese entirely and producing elemental chlorine more economically for industrial use.[20] By 1900, the Weldon and Deacon methods together supported annual chlorine output sufficient for about 150,000 tons of bleaching powder in England alone.[19] Sodium hydroxide (caustic soda) production predated chlorine's industrial scale and occurred separately via causticization of soda ash with slaked lime: \text{Na}_2\text{CO}_3 + \text{Ca(OH)}_2 \rightarrow 2 \text{NaOH} + \text{CaCO}_3. This batch process, using soda ash from the Leblanc method, became industrial standard by 1853, yielding purified NaOH after filtration and evaporation for applications in soap, paper, and textiles.[21] Until the late 19th century, these non-electrolytic routes decoupled chlorine and caustic soda manufacture, with no integrated process linking brine directly to both products, as hydrochloric acid from Leblanc fueled chlorine but not hydroxide synthesis.[3]Emergence of Electrolytic Production
The electrolytic decomposition of brine to produce chlorine gas and sodium hydroxide (caustic soda) was first demonstrated experimentally in 1800 by William Cruikshank, who electrolyzed a sodium chloride solution to generate chlorine at the anode, though this remained a laboratory curiosity without practical application due to inefficient power sources and lack of product separation.[19] Practical advancements began in 1851 when Charles Watt secured a British patent for an electrolytic cell designed to yield chlorine, caustic soda, and sodium hypochlorite from brine, marking the initial conceptualization of a coupled production process; however, high energy consumption from early dynamos and inadequate diaphragm materials rendered it uneconomical compared to chemical synthesis methods like the Leblanc process.[22][23] Industrial viability emerged in the 1890s amid improvements in electrical generation, particularly with alternating current transmission enabling cheaper power, and innovations in cell design to prevent anode-cathode product mixing. Hamilton Castner patented a mercury cathode cell in 1892, in which sodium amalgamated with mercury at the cathode, facilitating separation of chlorine and hydrogen gases while allowing caustic soda recovery via decomposition of the amalgam, thus achieving purer outputs than prior unseparated systems.[24] Concurrently, Austrian engineer Carl Kellner developed a variant incorporating a rocking mercury cell to enhance circulation and efficiency, leading to the joint Castner-Kellner process licensed for commercial use.[25] The first full-scale electrolytic chloralkali plant commenced operations in 1892 at Rumford Falls, Maine, employing an early mercury-based design to produce approximately 1 ton of chlorine per day, signaling the shift from batch chemical methods to continuous electrochemical production capable of meeting rising demand for disinfectants, bleaches, and alkalies in textiles and soap manufacturing.[26] By the mid-1890s, similar facilities proliferated in the United States and Europe, with Niagara Falls becoming a hub due to abundant hydroelectric power; for instance, a plant there utilized 68 Townsend cells operating at 2 kA each to generate low-hypochlorite caustic soda, underscoring rapid scale-up driven by energy cost reductions from 50 cents per kWh in the 1880s to under 2 cents by 1900.[19] These developments supplanted non-electrolytic routes, as electrolytic processes offered higher yields—up to 95% current efficiency in early mercury cells—and co-production of valuable hydrogen, though initial challenges included mercury handling and graphite anode corrosion, later addressed by material refinements.[22]Major Technological Shifts and Scale-Up
The mercury cell process, commercialized in 1892 through the Castner-Kellner design, represented the first major technological shift enabling large-scale chloralkali production. Unlike earlier diaphragm cells introduced in 1885, which allowed partial mixing of chlorine and hydroxide products leading to contamination, the mercury cathode formed a sodium amalgam that prevented direct contact between anode and cathode compartments. This innovation permitted continuous operation and higher-purity caustic soda output, transitioning the industry from small-batch electrolytic setups to facilities capable of producing hundreds of tons annually by the early 1900s.[27][28] Mid-20th-century advancements further drove scale-up, including the development of dimensionally stable anodes in the 1960s, which lowered cell voltage by reducing oxygen evolution and extended electrode life from months to years. These improvements, combined with optimized brine purification and larger cell configurations, supported the expansion of production capacities amid rising demand for chlorine in PVC manufacturing and water treatment post-World War II. By the 1970s, mercury and diaphragm cells dominated, with global chlorine production reaching tens of millions of tons, reflecting the cumulative effects of these engineering refinements on energy efficiency and throughput.[3] The emergence of membrane cell technology in the 1970s marked a transformative shift, motivated by the oil crises' emphasis on energy reduction and regulatory pressures against mercury pollution. Ion-exchange membranes, pioneered with perfluorosulfonic acid types like Nafion by DuPont around 1962 but scaled for chloralkali in the early 1970s, selectively permitted sodium ion transport while minimizing hydroxide back-migration, achieving current efficiencies above 95% and energy use 25-30% lower than mercury cells. Adoption accelerated in the 1980s, with new plants favoring membranes for their environmental compliance and operational reliability, enabling modular designs with hundreds of cells per electrolyzer string and individual facilities exceeding 1 million tons of chlorine capacity annually by the 2000s. This evolution has underpinned sustained industry growth, prioritizing causal factors like reduced operational costs and regulatory imperatives over legacy methods.[29][30][31]Process Technologies
Mercury Cell Technology
The mercury cell process, developed in the late 19th century, utilizes liquid mercury as the cathode in the electrolysis of brine to produce chlorine gas, sodium hydroxide, and hydrogen. In the electrolytic cell, a saturated sodium chloride solution flows over a horizontal layer of mercury, while anodes, typically graphite or dimensionally stable anodes, are suspended above. At the anode, chloride ions are oxidized to chlorine gas according to the reaction \ce{2Cl^- -> Cl2 + 2e^-}, which is collected, cooled, dried, and compressed for storage. At the mercury cathode, sodium ions are reduced and amalgamate with mercury: \ce{Na^+ + e^- + Hg -> Na(Hg)}, forming a sodium-mercury amalgam that flows continuously to a separate decomposer vessel.[32][33] In the decomposer, the amalgam reacts with water under controlled conditions, typically with a graphite packing to facilitate the reaction: \ce{2Na(Hg) + 2H2O -> 2NaOH + H2 + 2Hg}, regenerating the mercury for recirculation and yielding a 50% sodium hydroxide solution with low salt content, alongside hydrogen gas. This two-stage separation ensures high-purity caustic soda without the need for extensive post-processing evaporation or purification, a key operational advantage over diaphragm cells. The process requires approximately 3,200–3,400 kWh of electricity per metric ton of chlorine produced, higher than modern membrane cells due to overpotential losses at the mercury cathode. Mercury consumption is minimal, around 10–50 g per metric ton of chlorine, but cumulative losses through vaporization, effluent, and product contamination have historically totaled several hundred kilograms per plant annually.[32][3][34] Introduced commercially in the 1890s following innovations by Hamilton Castner, the mercury cell became the dominant chloralkali technology by the early 20th century, supplanting earlier diaphragm methods due to its superior product quality and efficiency in producing concentrated, impurity-free caustic soda suitable for rayon and chemical synthesis applications. Plants operated with cell rooms featuring multiple horizontal cells in series, enabling large-scale production; by the mid-20th century, it accounted for the majority of global capacity. Advantages included excellent anode-cathode separation minimizing hypochlorite formation and consistent output of 48–50% NaOH, reducing energy for concentration compared to weaker solutions from other cells. However, drawbacks encompassed higher capital and maintenance costs from mercury handling, elevated energy demands, and environmental risks from mercury's toxicity, including bioaccumulation leading to neurological damage in ecosystems and human populations.[21][3][34] Mercury emissions, primarily via air (stack gases), water (brine purges and decomposer effluents), and trace contamination in products, prompted regulatory scrutiny starting in the 1970s amid evidence of widespread pollution, such as elevated mercury levels in sediments near facilities. The process's reliance on toxic mercury conflicted with emerging environmental standards, leading to voluntary industry commitments and mandates for conversion. In Europe, the chloralkali sector pledged in 2001 to phase out mercury cells by 2020, achieving complete elimination via conversion to membrane technology or closures. Globally, the Minamata Convention on Mercury requires phase-out of mercury cell production by 2025, with most facilities transitioning to non-mercurial alternatives; as of 2019, only two U.S. plants remained operational, subject to EPA rules prohibiting emissions by 2025–2028 through shutdown or conversion. Residual mercury management, including decontamination of equipment and waste treatment, remains a challenge, with decontamination processes recovering over 99% of cell mercury but requiring specialized handling to prevent releases.[34][35][36]Diaphragm Cell Technology
The diaphragm cell process in chloralkali electrolysis employs a permeable separator to divide the anode compartment, where chlorine gas evolves, from the cathode compartment, where hydrogen gas and sodium hydroxide form from saturated sodium chloride brine. During operation, brine flows downward through the diaphragm into the cathode area, minimizing back-migration of hydroxide ions while permitting sodium ions to pass; chlorine gas bubbles collect at the anode (typically dimensionally stable titanium coated with mixed metal oxides), while at the iron or steel cathode, water reduction produces hydrogen and dilute caustic liquor containing 10-12% NaOH contaminated with residual salt. The resulting chlorine stream requires purification to remove oxygen and moisture via compression, cooling, and caustic scrubbing.[37][3] Traditionally, the diaphragm consists of asbestos fibers packed onto a perforated cathode plate to form a porous barrier approximately 1-3 mm thick, which controls ion transport but allows some anolyte-catholyte mixing, necessitating downstream evaporation of the caustic effluent to achieve 50% NaOH concentration using multiple-effect evaporators that consume significant steam (around 1.1-1.3 tons per ton of 100% NaOH). Due to asbestos's carcinogenic properties upon inhalation, leading to asbestosis and mesothelioma, modern variants employ non-asbestos materials such as fluoropolymer-based separators like polytetrafluoroethylene (PTFE) microfibril mats (e.g., Tephram) or proprietary non-carcinogenic composites, which maintain permeability while reducing health risks and improving durability. These alternatives emerged in the 1990s and 2000s to comply with regulatory phase-outs of asbestos in industrial applications.[38][39] Compared to mercury cells, diaphragm cells require lower electrical energy (approximately 2,500-2,800 kWh per metric ton of Cl2) due to the absence of amalgam decomposition and simpler brine systems, and they avoid mercury pollution, though total process energy remains higher than membrane cells owing to evaporation demands (adding 20-30% equivalent energy via steam). Capital costs are moderate, with simpler cell construction, but disadvantages include lower caustic purity (11-12% salt in initial liquor versus <50 ppm in membrane processes), higher maintenance from diaphragm degradation (lifespan 1-3 years), and fugitive chlorine emissions if not managed. Historically introduced in the early 20th century following mercury cell dominance, diaphragm technology peaked in the mid-20th century, accounting for about 67% of U.S. chlorine production by the 1970s, but has declined globally to under 20% by 2020 as membrane cells offer superior efficiency and product quality without asbestos.[3][40]Membrane Cell Technology
The membrane cell technology in the chloralkali process utilizes an ion-selective membrane to separate the anode and cathode compartments, enabling the production of chlorine gas at the anode, hydrogen gas and sodium hydroxide at the cathode, while minimizing mixing of products.[41] The membrane, typically a perfluorinated ion-exchange material such as Nafion with sulfonic or carboxylic functional groups, permits sodium ions and a limited amount of water to migrate from the anolyte (brine) to the catholyte, preventing hydroxide ions from passing in the opposite direction.[42] This selective permeability results in a caustic soda solution of approximately 30-33% concentration directly from the cell, with low salt contamination below 100 ppm NaCl.[3] Electrolysis occurs in a brine solution at the anode where chloride ions are oxidized to chlorine gas, while at the cathode water is reduced to hydrogen gas and hydroxide ions that combine with permeated sodium ions to form NaOH.[41] The process requires purified brine to avoid membrane fouling, and the produced chlorine may contain trace oxygen, necessitating liquefaction and purification steps.[41] Energy consumption is approximately 2,530-2,600 kWh per ton of chlorine, lower than mercury cells by about 26% and diaphragm cells.[3][43] Developed in the 1960s and first commercialized in the 1970s, membrane technology gained prominence due to its environmental advantages, avoiding mercury pollution and asbestos use associated with older methods.[43][44] By the late 1970s, initial U.S. commercial plants were operational, and it has since become the dominant process globally, driven by regulations phasing out mercury cells and its superior energy efficiency and product purity.[3][44] The technology requires less steam for caustic concentration—under one tonne per tonne of NaOH—further enhancing operational efficiency.[41]Alternative and Historical Variants
Early electrolytic chloralkali processes utilized undivided cells, where direct contact between anode and cathode compartments allowed chlorine generated at the anode to react with sodium hydroxide formed at the cathode, yielding sodium hypochlorite for bleaching applications rather than separate chlorine and caustic soda products.[45] These configurations, dating to the mid-19th century, suffered from low efficiency due to back-migration and unwanted reactions, limiting scalability until partitioned designs emerged.[19] The Castner–Kellner process, patented in 1892, marked a historical variant of mercury cell technology, employing a horizontal trough with flowing mercury as the cathode to form sodium amalgam, which was subsequently decomposed in a separate denuder to produce caustic soda and hydrogen while minimizing direct contact between chlorine and alkali.[46] This innovation improved product purity over earlier mercury setups by externalizing amalgam decomposition, facilitating commercial adoption in Europe and the United States by the early 20th century.[47] Modern alternatives include anion exchange membrane (AEM) electrolyzers, which conduct hydroxide ions and enable non-precious metal catalysts, potentially reducing capital costs by 30-50% compared to cation exchange membrane systems, though current energy efficiencies lag behind established methods.[48] Decoupled processes, demonstrated in laboratory settings since 2018, separate chlorine evolution from hydrogen and sodium hydroxide production using redox mediators like Na0.44MnO2 electrodes, eliminating membrane needs and enabling flexible operation with renewable energy inputs, albeit at smaller scales without widespread industrial deployment.[49][50]Operational Components
Electrode Materials and Design
In the chloralkali process, anodes are primarily responsible for the chlorine evolution reaction (CER), where chloride ions oxidize to form Cl₂ gas, and modern installations predominantly use dimensionally stable anodes (DSAs) comprising a titanium base substrate coated with mixed metal oxides such as ruthenium dioxide (RuO₂), iridium dioxide (IrO₂), and titanium dioxide (TiO₂).[51][52] These coatings, applied via thermal decomposition, confer corrosion resistance in the acidic anolyte (pH ~2-4) and reduce CER overpotential by 200-300 mV compared to legacy graphite anodes, which degraded via exfoliation and increased cell voltage over time.[50][53] DSA designs typically feature expanded titanium mesh or rod arrays to maximize geometric surface area (up to 500-1000 m²/m³) while facilitating gas bubble detachment and minimizing mass transport limitations, thereby achieving current efficiencies exceeding 95% and service lives of 5-10 years under industrial currents of 3-6 kA/m².[51][54] Cathodes support the hydrogen evolution reaction (HER) and hydroxide ion generation, with materials selected to withstand alkaline conditions (pH 12-14) and minimize HER overpotential for energy efficiency. In mercury cells, a flowing liquid mercury cathode amalgamates sodium atoms formed via Na⁺ reduction, preventing direct NaOH production but enabling high-purity output; mercury's high hydrogen overpotential (∼0.7 V) suppresses competing HER, though this design has been phased out due to environmental concerns.[8] Diaphragm cells employ perforated nickel or mild steel cathodes, often uncoated, which tolerate brine crossover but exhibit higher HER overpotentials (∼0.2-0.3 V) and require asbestos diaphragms for separation.[55] Membrane cells favor activated nickel cathodes with catalytic layers, such as Raney nickel or precious-metal-free coatings (e.g., Ni-Mo or Ni-P alloys), reducing overpotential to ∼0.1 V and enabling zero-gap configurations where the cathode presses against the ion-exchange membrane to cut ohmic losses by 20-30%.[50][56] Cathode geometries mirror anodes with mesh structures to enhance electrolyte contact and H₂ bubble release, supporting current densities up to 6 kA/m² with efficiencies >99% for H₂ production.[57] Electrode design optimizations focus on inter-electrode spacing, typically reduced to <3 mm in advanced cells to lower IR drop (∼0.1-0.2 V savings), and surface texturing to mitigate bubble coverage, which can block 20-30% of active sites if unaddressed.[56][11] Coatings are engineered for uniform thickness (5-10 μm) and adhesion via multiple firing cycles at 400-500°C, ensuring stability against deactivation mechanisms like Ru dissolution or phase segregation under anodic potentials of ∼1.3-1.4 V vs. SHE.[58] These advancements have decreased overall cell voltage from ∼4.5 V in early designs to ∼3.0-3.2 V today, correlating with 20-25% energy reductions per ton of Cl₂ produced.[11][50]Brine Purification and Preparation
The preparation of brine for the chloralkali process involves dissolving sodium chloride—sourced from rock salt, vacuum-evaporated salt, or solar-evaporated sea salt—in purified water or recycled dilute brine to achieve saturation at approximately 300–330 g/L NaCl, depending on the cell technology used (higher for membrane cells at up to 445 g/L).[3] This step incorporates heating and agitation to facilitate dissolution while minimizing insoluble impurities from the salt source.[3] Recycled brine from the electrolysis cells is often blended in to optimize salt recovery and reduce fresh water usage.[59] Purification is essential to eliminate impurities that cause electrode scaling, membrane fouling, elevated cell voltage, and diminished current efficiency; common contaminants include calcium (Ca²⁺), magnesium (Mg²⁺), sulfate (SO₄²⁻), iron (Fe³⁺), aluminum (Al³⁺), and trace organics or silica.[3] Primary purification targets bulk removal through chemical precipitation: sodium carbonate (Na₂CO₃) is added to form insoluble calcium carbonate (CaCO₃), while sodium hydroxide (NaOH) or calcium hydroxide (Ca(OH)₂) precipitates magnesium as magnesium hydroxide (Mg(OH)₂); iron and aluminum are similarly removed as hydroxides.[3][59] Sulfate is addressed by adding calcium chloride (CaCl₂) to produce calcium sulfate sludge, often with sodium hypochlorite (NaOCl) if ammonia is detected to oxidize organics.[3] The mixture is then clarified in settling tanks and filtered via sand beds, pressure leaf filters, or candle filters to separate the brine mud—a sludge of precipitated salts generating about 30 kg per 1,000 kg of chlorine produced, varying with raw salt quality.[3] For diaphragm and mercury cells, primary treatment suffices with hardness limits of <5 ppm Ca²⁺ and <0.5 ppm Mg²⁺, alongside sulfate below 5 g/L.[3] Membrane cells demand secondary purification due to their sensitivity, employing chelating ion-exchange resins or nanofiltration to polish the brine to ultra-low levels: combined Ca²⁺ and Mg²⁺ below 20 ppb, and sulfate controlled to 4.7–6.8 g/L to prevent membrane degradation.[3][60][61] Resins are regenerated periodically with hydrochloric acid (HCl) and NaOH, and pH is adjusted to 10–11 to optimize precipitation while avoiding over-acidification that could harm downstream components.[59] Final steps include reheating, salt resaturation if needed, and quality checks via online analyzers for real-time impurity monitoring.[59]| Impurity | Diaphragm/Mercury Cell Limit | Membrane Cell Limit |
|---|---|---|
| Calcium (Ca²⁺) | <5 ppm | <20 ppb |
| Magnesium (Mg²⁺) | <0.5 ppm | <20 ppb |
| Sulfate (SO₄²⁻) | <5 g/L | <4.7–6.8 g/L |
Cell Design and Energy Requirements
The chloralkali process employs electrolytic cells configured as either monopolar or bipolar assemblies, where multiple elementary cells are arranged in series or parallel to optimize voltage distribution and current flow for industrial-scale operation. Monopolar designs connect cells in parallel, facilitating easier maintenance but requiring more electrical connections and higher rectification capacity, whereas bipolar designs stack cells in series, minimizing inter-cell wiring and reducing overall energy losses through lower voltage per cell but complicating individual cell replacement. In both configurations, the cell structure includes an anode compartment for brine electrolysis producing chlorine gas, a cathode compartment for hydrogen evolution and caustic formation, and a separator—mercury cathode, porous diaphragm, or selective ion-exchange membrane—to prevent product mixing while allowing ion transport. Electrode spacing is minimized in modern zero-gap designs, particularly in membrane cells, to reduce ohmic resistance and associated energy dissipation.[63] Energy requirements in chloralkali cells arise from the thermodynamic decomposition potential of approximately 2.2 volts (accounting for standard electrode potentials of 1.36 V for chlorine evolution and -0.83 V for hydrogen evolution), augmented by kinetic overpotentials and ohmic losses. Anodic overpotential for chlorine evolution on dimensionally stable anodes (typically ruthenium-iridium oxide-coated titanium) ranges from 0.1 to 0.2 V at current densities of 2-4 kA/m², while cathodic overpotential for hydrogen evolution on nickel cathodes is about 0.2-0.3 V; these contribute minimally in optimized systems but increase with impurities or fouling. The dominant energy sink is ohmic drop (IR), encompassing electrolyte resistance (dependent on brine conductivity, typically 200-300 mS/cm at 30-90% saturation), separator resistance (lowest in thin perfluorosulfonic acid membranes at ~0.1 Ω cm²), and bubble-induced effects that elevate effective resistance by up to 20% if not mitigated by gas disengagement zones. Total cell voltage thus operates at 3.0-3.5 V in efficient membrane cells, with current efficiencies exceeding 95% for both chlorine and caustic production, reflecting Faraday efficiencies limited by side reactions like oxygen evolution (<2%) or back-migration of hydroxyl ions.[64][11] Specific energy consumption varies by cell technology due to differences in separator resistance and process integration. Membrane cells achieve the lowest values at 2,200-2,500 kWh per metric ton of Cl₂, benefiting from low-resistance cation-exchange membranes (e.g., Nafion-type) that permit selective Na⁺ transport while minimizing water and OH⁻ crossover, enabling operation at higher caustic concentrations (30-35 wt%) without additional evaporation energy. Diaphragm cells consume 2,800-2,900 kWh/t Cl₂, as porous asbestos or polymeric diaphragms introduce higher IR drop (~0.3-0.5 Ω cm²) and require dilution of the anolyte to prevent mixing, followed by energy-intensive caustic concentration via evaporation (adding ~0.5-1.0 MWh/t NaOH equivalent). Mercury cells, now largely phased out, demanded 3,200-3,400 kWh/t Cl₂ owing to amalgam transport inefficiencies and higher cathodic overpotentials, though they offered high purity products. These figures assume direct current efficiencies of 90-98% and exclude auxiliary power for pumps, compression, and purification, which add 10-20% to total site energy use.[63][59]| Cell Technology | Typical Cell Voltage (V) | Energy Consumption (kWh/t Cl₂) | Current Efficiency (%) |
|---|---|---|---|
| Membrane | 3.0-3.2 | 2,200-2,500 | 95-98 |
| Diaphragm | 3.3-3.5 | 2,800-2,900 | 92-96 |
| Mercury | 3.8-4.2 | 3,200-3,400 | 90-95 |