Chloralkali process

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Old drawing of a chloralkali process plant (Edgewood, Maryland) Old drawing of a Chlorine-Caustic Soda Plant (Edgewood, Maryland).JPG
Old drawing of a chloralkali process plant (Edgewood, Maryland)

The chloralkali process (also chlor-alkali and chlor alkali ) is an industrial process for the electrolysis of sodium chloride (NaCl) solutions. It is the technology used to produce chlorine and sodium hydroxide (caustic soda), [1] which are commodity chemicals required by industry. Thirty five million tons of chlorine were prepared by this process in 1987. [2] The chlorine and sodium hydroxide produced in this process are widely used in the chemical industry.

Contents

Usually the process is conducted on a brine (an aqueous solution of NaCl), in which case sodium hydroxide (NaOH), hydrogen, and chlorine result. When using calcium chloride or potassium chloride, the products contain calcium or potassium instead of sodium. Related processes are known that use molten NaCl to give chlorine and sodium metal or condensed hydrogen chloride to give hydrogen and chlorine.

The process has a high energy consumption, for example around 2,500 kWh (9,000 MJ) of electricity per tonne of sodium hydroxide produced. Because the process yields equivalent amounts of chlorine and sodium hydroxide (two moles of sodium hydroxide per mole of chlorine), it is necessary to find a use for these products in the same proportion. For every mole of chlorine produced, one mole of hydrogen is produced. Much of this hydrogen is used to produce hydrochloric acid, ammonia, hydrogen peroxide, or is burned for power and/or steam production. [3]

History

The chloralkali process has been in use since the 19th century and is a primary industry in the United States, Western Europe, and Japan. [4] [5] It has become the principal source of chlorine during the 20th century. [6] The diaphragm cell process and the mercury cell process have been used for over 100 years but are environmentally unfriendly through their use of asbestos and mercury, respectively. The membrane cell process, which was only developed in the past 60 years, is a superior method with its improved energy efficiency and lack of harmful chemicals. [5]

Although the first formation of chlorine by the electrolysis of brine was attributed to chemist William Cruikshank in 1800, it was 90 years later that the electrolytic method was used successfully on a commercial scale. Industrial scale production began in 1892. [7] In 1833, Faraday formulated the laws that governed the electrolysis of aqueous solutions, and patents were issued to Cook and Watt in 1851 and to Stanley in 1853 for the electrolytic production of chlorine from brine. [7]

Cell room of a chlor-alkali plant ca. 1920 Cell room of a chlorine-caustic soda plant.JPG
Cell room of a chlor-alkali plant ca. 1920

Process systems

Three production methods are in use. While the mercury cell method produces chlorine-free sodium hydroxide, the use of several tonnes of mercury leads to serious environmental problems. In a normal production cycle a few hundred pounds of mercury per year are emitted, which accumulate in the environment. Additionally, the chlorine and sodium hydroxide produced via the mercury-cell chloralkali process are themselves contaminated with trace amounts of mercury. The membrane and diaphragm method use no mercury, but the sodium hydroxide contains chlorine, which must be removed.

Membrane cell

The most common chloralkali process involves the electrolysis of aqueous sodium chloride (a brine) in a membrane cell. A membrane, such as one made from Nafion, Flemion or Aciplex, is used to prevent the reaction between the chlorine and hydroxide ions.

Basic membrane cell used in the electrolysis of brine. At the anode (A), chloride (Cl ) is oxidized to chlorine. The ion-selective membrane (B) allows the counterion Na+ to freely flow across, but prevents anions such as hydroxide (OH ) and chloride from diffusing across. At the cathode (C), water is reduced to hydroxide and hydrogen gas. The net process is the electrolysis of an aqueous solution of NaCl into industrially useful products sodium hydroxide (NaOH) and chlorine gas. Chloralkali membrane.svg
Basic membrane cell used in the electrolysis of brine. At the anode (A), chloride (Cl ) is oxidized to chlorine. The ion-selective membrane (B) allows the counterion Na+ to freely flow across, but prevents anions such as hydroxide (OH ) and chloride from diffusing across. At the cathode (C), water is reduced to hydroxide and hydrogen gas. The net process is the electrolysis of an aqueous solution of NaCl into industrially useful products sodium hydroxide (NaOH) and chlorine gas.

Saturated brine is passed into the first chamber of the cell where the chloride ions are oxidised at the anode, losing electrons to become chlorine gas (A in figure):

2Cl− → Cl
2 + 2e−

At the cathode, positive hydrogen ions pulled from water molecules are reduced by the electrons provided by the electrolytic current, to hydrogen gas, releasing hydroxide ions into the solution (C in figure):

2H
2O + 2e− → H2 + 2OH−

The ion-permeable ion-exchange membrane at the center of the cell allows the sodium ions (Na+) to pass to the second chamber where they react with the hydroxide ions to produce caustic soda (NaOH) (B in figure). [1] The overall reaction for the electrolysis of brine is thus:

2NaCl + 2H
2O → Cl
2 + H
2 + 2NaOH

Diaphragm cell

In the diaphragm cell process, there are two compartments separated by a permeable diaphragm, often made of asbestos fibers. Brine is introduced into the anode compartment and flows into the cathode compartment. Similarly to the membrane cell, chloride ions are oxidized at the anode to produce chlorine, and at the cathode, water is split into caustic soda and hydrogen. The diaphragm prevents the reaction of the caustic soda with the chlorine. A diluted caustic brine leaves the cell. The caustic soda must usually be concentrated to 50% and the salt removed. This is done using an evaporative process with about three tonnes of steam per tonne of caustic soda. The salt separated from the caustic brine can be used to saturate diluted brine. The chlorine contains oxygen and must often be purified by liquefaction and evaporation.

Mercury cell

Mercury cell for chloralkali process HgNaOHElectrolysis.png
Mercury cell for chloralkali process

In the mercury-cell process, also known as the Castner–Kellner process, a saturated brine solution floats on top of a thin layer of mercury. The mercury is the cathode, where sodium is produced and forms an amalgam with the mercury. The amalgam is continuously drawn out of the cell and reacted with water which decomposes the amalgam into sodium hydroxide, hydrogen and mercury. The mercury is recycled into the electrolytic cell. Chlorine is produced at the anode and bubbles out of the cell. Mercury cells are being phased out due to concerns about mercury poisoning from mercury cell pollution such as occurred in Canada (see Ontario Minamata disease) and Japan (see Minamata disease).

Unpartitioned cell

The initial overall reaction produces hydroxide and also hydrogen and chlorine gases: [8]

2 NaCl + 2 H2O → 2 NaOH + H2 + Cl2

Without a membrane, the OH− ions produced at the cathode are free to diffuse throughout the electrolyte. As the electrolyte becomes more basic due to the production of OH−, less Cl2 emerges from the solution as it begins to disproportionate to form chloride and hypochlorite ions at the anode:

Cl2 + 2 NaOH → NaCl + NaClO + H2O

The more opportunity the Cl2 has to interact with NaOH in the solution, the less Cl2 emerges at the surface of the solution and the faster the production of hypochlorite progresses. This depends on factors such as solution temperature, the amount of time the Cl2 molecule is in contact with the solution, and concentration of NaOH.

Likewise, as hypochlorite increases in concentration, chlorates are produced from them:

3 NaClO → NaClO3 + 2 NaCl

This reaction is accelerated at temperatures above about 60 °C. Other reactions occur, such as the self-ionization of water and the decomposition of hypochlorite at the cathode, the rate of the latter depends on factors such as diffusion and the surface area of the cathode in contact with the electrolyte. [9]

If current is interrupted while the cathode is submerged, cathodes that are attacked by hypochlorites, such as those made from stainless steel, will dissolve in unpartitioned cells.

If producing hydrogen and oxygen gases is not a priority, the addition of 0.18% sodium or potassium chromate to the electrolyte will improve the efficiency of producing the other products. [9]

Electrodes

Due to the corrosive nature of chlorine production, the anode (where the chlorine is formed) must be non-reactive and has been made from materials such as platinum metal, [10] graphite (called plumbago in Faraday's time), [10] or platinized titanium. [11] A mixed metal oxide clad titanium anode (also called a dimensionally stable anode) is the industrial standard today. Historically, platinum, magnetite, lead dioxide, [12] manganese dioxide, and ferrosilicon (13–15% silicon [13] ) have also been used as anodes. [14] Platinum alloyed with iridium is more resistant to corrosion from chlorine than pure platinum. [14] [15] Unclad titanium cannot be used as an anode because it anodizes, forming a non-conductive oxide and passivates. Graphite will slowly disintegrate due to internal electrolytic gas production from the porous nature of the material and carbon dioxide forming due to carbon oxidation, causing fine particles of graphite to be suspended in the electrolyte that can be removed by filtration. The cathode (where hydroxide forms) can be made from unalloyed titanium, graphite, or a more easily oxidized metal such as stainless steel or nickel.

Manufacturer associations

The interests of chloralkali product manufacturers are represented at regional, national and international levels by associations such as Euro Chlor and The World Chlorine Council.

See also

Related Research Articles

<span class="mw-page-title-main">Chlorine</span> Chemical element, symbol Cl and atomic number 17

Chlorine is a chemical element; it has symbol Cl and atomic number 17. The second-lightest of the halogens, it appears between fluorine and bromine in the periodic table and its properties are mostly intermediate between them. Chlorine is a yellow-green gas at room temperature. It is an extremely reactive element and a strong oxidising agent: among the elements, it has the highest electron affinity and the third-highest electronegativity on the revised Pauling scale, behind only oxygen and fluorine.

<span class="mw-page-title-main">Electrochemistry</span> Branch of chemistry

Electrochemistry is the branch of physical chemistry concerned with the relationship between electrical potential difference and identifiable chemical change. These reactions involve electrons moving via an electronically-conducting phase between electrodes separated by an ionically conducting and electronically insulating electrolyte.

<span class="mw-page-title-main">Electrolysis</span> Technique in chemistry and manufacturing

In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell. The voltage that is needed for electrolysis to occur is called the decomposition potential. The word "lysis" means to separate or break, so in terms, electrolysis would mean "breakdown via electricity".

The term chloride refers either to a chloride ion, which is a negatively charged chlorine atom, or a non-charged chlorine atom covalently bonded to the rest of the molecule by a single bond. Many inorganic chlorides are salts. Many organic compounds are chlorides. The pronunciation of the word "chloride" is.

<span class="mw-page-title-main">Brine</span> Concentrated solution of salt in water

Brine is water with a high-concentration solution of salt. In diverse contexts, brine may refer to the salt solutions ranging from about 3.5% up to about 26%. Brine forms naturally due to evaporation of ground saline water but it is also generated in the mining of sodium chloride. Brine is used for food processing and cooking, for de-icing of roads and other structures, and in a number of technological processes. It is also a by-product of many industrial processes, such as desalination, so it requires wastewater treatment for proper disposal or further utilization.

<span class="mw-page-title-main">Sodium chloride</span> Chemical compound with formula NaCl

Sodium chloride, commonly known as table salt, is an ionic compound with the chemical formula NaCl, representing a 1:1 ratio of sodium and chloride ions. Sodium chloride is the salt most responsible for the salinity of seawater and of the extracellular fluid of many multicellular organisms. In its edible form, it is commonly used as a condiment and food preservative. Large quantities of sodium chloride are used in many industrial processes, and it is a major source of sodium and chlorine compounds used as feedstocks for further chemical syntheses. Another major application of sodium chloride is deicing of roadways in sub-freezing weather.

<span class="mw-page-title-main">Sodium hypochlorite</span> Chemical compound (known in solution as bleach)

Sodium hypochlorite, commonly known in a dilute solution as (chlorine) bleach, is an alkaline inorganic chemical compound with the formula NaOCl, consisting of a sodium cation and a hypochlorite anion. It may also be viewed as the sodium salt of hypochlorous acid. The anhydrous compound is unstable and may decompose explosively. It can be crystallized as a pentahydrate NaOCl·5H
2O, a pale greenish-yellow solid which is not explosive and is stable if kept refrigerated.

<span class="mw-page-title-main">Potassium hydroxide</span> Inorganic compound (KOH)

Potassium hydroxide is an inorganic compound with the formula KOH, and is commonly called caustic potash.

<span class="mw-page-title-main">Electrolytic cell</span> Cell that uses electrical energy to drive a non-spontaneous redox reaction

An electrolytic cell is an electrochemical cell that utilizes an external source of electrical energy to force a chemical reaction that would otherwise not occur. The external energy source is a voltage applied between the cell's two electrodes; an anode and a cathode, which are immersed in an electrolyte solution. This is in contrast to a galvanic cell, which itself is a source of electrical energy and the foundation of a battery. The net reaction taking place in a galvanic cell is a spontaneous reaction, i.e., the Gibbs free energy remains -ve, while the net reaction taking place in an electrolytic cell is the reverse of this spontaneous reaction, i.e., the Gibbs free energy is +ve.

<span class="mw-page-title-main">Hypochlorite</span> Ion

In chemistry, hypochlorite, or chloroxide is an anion with the chemical formula ClO−. It combines with a number of cations to form hypochlorite salts. Common examples include sodium hypochlorite and calcium hypochlorite. The Cl-O distance in ClO− is 1.69 Å.

<span class="mw-page-title-main">Sodium chlorate</span> Chemical compound

Sodium chlorate is an inorganic compound with the chemical formula NaClO3. It is a white crystalline powder that is readily soluble in water. It is hygroscopic. It decomposes above 300 °C to release oxygen and leaves sodium chloride. Several hundred million tons are produced annually, mainly for applications in bleaching pulp to produce high brightness paper.

<span class="mw-page-title-main">Electrolysis of water</span> Electricity-induced chemical reaction

Electrolysis of water is using electricity to split water into oxygen and hydrogen gas by electrolysis. Hydrogen gas released in this way can be used as hydrogen fuel, but must be kept apart from the oxygen as the mixture would be extremely explosive. Separately pressurised into convenient 'tanks' or 'gas bottles', hydrogen can be used for oxyhydrogen welding and other applications, as the hydrogen / oxygen flame can reach approximately 2,800°C.

Salt water chlorination is a process that uses dissolved salt for the chlorination of swimming pools and hot tubs. The chlorine generator uses electrolysis in the presence of dissolved salt to produce chlorine gas or its dissolved forms, hypochlorous acid and sodium hypochlorite, which are already commonly used as sanitizing agents in pools. Hydrogen is produced as byproduct too.

<span class="mw-page-title-main">Electrolysed water</span> Chemical mixture in water solution

Electrolysed water is produced by the electrolysis of ordinary tap water containing dissolved sodium chloride. The electrolysis of such salt solutions produces a solution of hypochlorous acid and sodium hydroxide. The resulting water can be used as a disinfectant.

<span class="mw-page-title-main">Downs cell</span>

Downs' process is an electrochemical method for the commercial preparation of metallic sodium, in which molten NaCl is electrolyzed in a special apparatus called the Downs cell. The Downs cell was invented in 1923 by the American chemist James Cloyd Downs (1885–1957).

Calcium chlorate is the calcium salt of chloric acid, with the chemical formula Ca(ClO3)2. Like other chlorates, it is a strong oxidizer.

The Castner–Kellner process is a method of electrolysis on an aqueous alkali chloride solution to produce the corresponding alkali hydroxide, invented by American Hamilton Castner and Austrian Carl Kellner in the 1890s. Due to lower energy cost and fewer environmental concerns, the Castner–Kellner process is being replaced gradually with membrane electrolysis.

Chlorine gas can be produced by extracting from natural materials, including the electrolysis of a sodium chloride solution (brine) and other ways.

Electrochlorination is the process of producing hypochlorite by passing electric current through salt water. This disinfects the water and makes it safe for human use, such as for drinking water or swimming pools.

A mixed oxidant solution (MOS) is a type of disinfectant that has many uses including disinfecting, sterilizing, and eliminating pathogenic microorganisms in water. An MOS may have advantages such as a higher disinfecting power, stable residual chlorine in water, elimination of biofilm, and safety. The main components of an MOS are chlorine and its derivatives, which are produced by electrolysis of sodium chloride. It may also contain high amounts of hydroxy radicals, chlorine dioxide, dissolved ozone, hydrogen peroxide and oxygen from which the name "mixed oxidant" is derived.

References

  1. 1 2 Fengmin Du; David M Warsinger; Tamanna I Urmi; et al. (2018). "Sodium hydroxide production from seawater desalination brine: process design and energy efficiency". Environmental Science & Technology. 52 (10): 5949–5958. Bibcode:2018EnST...52.5949D. doi:10.1021/acs.est.8b01195. hdl: 1721.1/123096 . PMID   29669210.
  2. ↑ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN   978-0-08-037941-8.
  3. ↑ R. Norris Shreve; Joseph Brink (1977). Chemical Process Industries (4th ed.). p. 219. ASIN   B000OFVCCG.
  4. ↑ Crook, Jedidiah; Mousavi, Aliyar (2016-07-02). "The chlor-alkali process: A review of history and pollution". Environmental Forensics. 17 (3): 211–217. Bibcode:2016EnvFo..17..211C. doi:10.1080/15275922.2016.1177755. ISSN   1527-5922. S2CID   99354861.
  5. 1 2 "The chlor-alkali process: A review of history and pollution". ResearchGate. Retrieved 2020-10-05.
  6. ↑ "Chlor-alkali process". Encyclopedia Britannica. Retrieved 2020-10-05.
  7. 1 2 O'Brien, Thomas F.; Bommaraju, Tilak V.; Hine, Fumio, eds. (2005). "History of the Chlor-Alkali Industry". Handbook of Chlor-Alkali Technology. Boston, MA: Springer. pp. 17–36. doi:10.1007/0-306-48624-5_2. ISBN   978-0-306-48624-1 . Retrieved 2020-10-05.
  8. ↑ Tilley, R.J.D. (2004). Understanding solids: the science of materials. John Wiley and Sons. pp. 281–. Bibcode:2004usts.book.....T. ISBN   978-0-470-85276-7 . Retrieved 22 October 2011.
  9. 1 2 Thompson, M. de Kay (1911). Applied Electrochemistry. The MacMillan Company. pp.  89-90.
  10. 1 2 Faraday, Michael (1849). Experimental Researches In Electricity. Vol. 1. London: The University of London.
  11. ↑ Landolt, D.; Ibl, N. (1972). "Anodic chlorate formation on platinized titanium". Journal of Applied Electrochemistry. Chapman and Hall Ltd. 2 (3): 201–210. doi:10.1007/BF02354977. S2CID   95515683.
  12. ↑ Munichandraiah, N.; Sathyanarayana, S. (1988). "Insoluble anode of α-lead dioxide coated on titanium for electrosynthesis of sodium perchlorate". Journal of Applied Electrochemistry. Chapman and Hall Ltd. 18 (2): 314–316. doi:10.1007/BF01009281. S2CID   96759724.
  13. ↑ Dinan, Charles (1927-10-15). The Corrosion of Durion Anodes (BSc). Massachusetts Institute of Technology. p. 4. hdl:1721.1/87815 . Retrieved 2019-09-25.
  14. 1 2 Hale, Arthur (1918). The Applications of Electrolysis in Chemical Industry. Longmans, Green, and Co. p.  13 . Retrieved 2019-09-15.
  15. ↑ Denso, P. (1902). "Untersuchungen über die Widerstandsfähigkeit von Platiniridium-Anoden bei der Alkalichlorid-Elektrolyse". Zeitschrift für Elektrochemie. 8 (10): 149–150. doi:10.1002/bbpc.19020081004.

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