9.5.D - Sodium Hydroxide


Sodium hydroxide (NaOH), also known as lye (USA) or caustic soda, is a caustic metallic base. It is widely used in industry, mostly as a strong chemical base in the manufacture of pulp and paper, textiles, drinking water, and detergents. Worldwide production in 1998 was around 45 million tonnes. Sodium hydroxide is also the most common base used in chemical laboratories.

Sodium hydroxide is produced (along with chlorine and hydrogen) via the chloralkali process. This involves the electrolysis of an aqueous solution of sodium chloride. The sodium hydroxide builds up at the cathode, where water is reduced to hydrogen gas and hydroxide ion. To produce NaOH it is necessary to prevent reaction of the NaOH with the chlorine. This is typically done in one of three ways, of which the membrane cell process is economically the most viable.

 

Diaphragm cell process - uses a steel cathode, and reaction of NaOH with Cl2 is prevented using a porous diaphragm.

Mercury cell process - sodium metal forms as an amalgam at a mercury cathode; this sodium is then reacted with water to produce NaOH. There have been concerns about mercury releases, although modern plants claim to be safe in this regard.

Membrane cell process - similar to the diaphragm cell process, with a modified Teflon membrane to separate the cathode and anode reactions. It is less expensive than the diaphragm cell process, and it produces a higher quality of NaOH.

In this topic students:

  • Explain the difference between galvanic cells and electrolytic cells in terms of energy requirements

  • Outline the steps in the industrial production of sodium hydroxide from sodium chloride solution

  • Distinguish between the three electrolysis methods used to extract sodium hydroxide: mercury process, diaphragm process and the membrane process by describing each process and analysing technical and environmental issues



Electrolytic Cells

Galvanic cells use a spontaneous chemical reaction to drive an electric current through an external circuit or to convert chemical energy into electrical energy. These cells are important because they are the basis for the batteries that fuel modern society. However, they are not the only kind of electrochemical cell. It is also possible to construct a cell that reverses the reactions in a galvanic cell by forcing electrons to flow in the opposite direction. In these cells electrical energy is converted to chemical energy and they are called electrolytic cells. Electrolysis is used to drive an oxidation-reduction reaction in a direction in which it does not occur spontaneously.

The Electrolysis of Molten Sodium Chloride

The electrolysis of molten sodium chloride.

An idealized cell for the electrolysis of sodium chloride is shown in the diagram on the right. A source of direct current is connected to a pair of inert electrodes immersed in molten sodium chloride. Because the salt has been heated until it melts, the Na+ ions flow toward the negative electrode and the Cl- ions flow toward the positive electrode.

When Na+ ions collide with the negative electrode, the battery carries a large enough potential to force these ions to pick up electrons to form sodium metal.

Negative electrode (cathode)

Na+ + e - → Na

Cl- ions that collide with the positive electrode are oxidised to Cl2 gas, which bubbles off at this electrode.

Positive electrode (anode)

2Cl - Cl 2 + 2e-

The net effect of passing an electric current through the molten salt in this cell is to decompose sodium chloride into its elements, sodium metal and chlorine gas.

The potential required to oxidize Cl- ions to Cl2 is -1.36 volts and the potential needed to reduce Na+ ions to sodium metal is -2.71 volts. The battery used to drive this reaction must therefore have a potential of at least 4.07 volts.

This example explains why the process is called electrolysis. The suffix -lysis comes from the Greek stem meaning to loosen or split up. Electrolysis literally uses an electric current to split a compound into its elements.

The overall reaction is

2NaCl (l) 2Na (l) + Cl 2 (g)

This example also illustrates the difference between voltaic cells and electrolytic cells. Voltaic cells use the energy given off in a spontaneous reaction to do electrical work. Electrolytic cells use electrical work as source of energy to drive the reaction in the opposite direction. Electrons are being pumped towards the negative electrode (where reduction is occurring).

A downs cell used to separate molten sodium chloride into
chlorine gas and sodium metal.

The dotted vertical line in the center of the figure on the previous page represents a diaphragm that keeps the Cl2 gas produced at the anode from coming into contact with the sodium metal generated at the cathode. The function of this diaphragm can be understood by turning to a more realistic drawing of the commercial Downs cell used to electrolyze sodium chloride shown in the diagram on the right.

Chlorine gas that forms on the graphite anode inserted into the bottom of this cell bubbles through the molten sodium chloride into a funnel at the top of the cell. Sodium metal that forms at the cathode floats up through the molten sodium chloride into a sodium-collecting ring, from which it is periodically drained. The diaphragm that separates the two electrodes is a screen of iron gauze, which prevents the explosive reaction that would occur if the products of the electrolysis reaction came in contact.

The feedstock for the Downs cell is a 3:2 mixture by mass of CaCl2 and NaCl. This mixture is used because it has a melting point of 580°C, whereas pure sodium chloride has to be heated to more than 800°C before it melts.

The Electrolysis of Aqueous Sodium Chloride

The electrolysis of aqueous sodium hydroxide.

A solution of sodium chloride will produce different electrolytic products than if pure molten sodium hydroxide were used. The diagram on the right shows an idealised drawing of a cell in which an aqueous solution of sodium chloride is electrolysed.

Once again, the Na+ ions migrate toward the negative electrode and the Cl- ions migrate toward the positive electrode. But, in this cell water is present and presents itself as a possible candidate for oxidation or reduction.

At the cathode (-) the possibilities are:

Na+ + e - → Na with Eø = -2.71 V

2H2O + 2e- H2 + 2OH- with Eø = -0.83 V

Because it is much easier to reduce water than Na+ ions, the only product formed at the cathode is hydrogen gas. The dodium ions are not reduced.

There are also two substances that can be oxidised at the anode (+) which are Cl- ions and water molecules.

2Cl - Cl2 + 2e- with E ø = -1.36 V

2H2O O2 + 4H+ + 4e- with E ø = -1.23 V

The standard-state potentials for these half-reactions are so close to each other that we might expect to see a mixture of Cl2 and O2 gas collect at the anode. In practice, the only product of this reaction is Cl2.

At first glance, it would seem easier to oxidize water (Eø = -1.23 volts) than Cl- ions (Eø = -1.36 volts). It is worth noting, however, that the cell is never allowed to reach standard-state conditions. The solution is typically 25% NaCl by mass, which significantly decreases the potential required to oxidize the Cl- ion. The pH of the cell is also kept very high, which decreases the oxidation potential for water. The deciding factor is a phenomenon known as overvoltage, which is the extra voltage that must be applied to a reaction to get it to occur at the rate at which it would occur in an ideal system.

Under ideal conditions, a potential of 1.23 volts is large enough to oxidize water to O2 gas. Under real conditions, however, it can take a much larger voltage to initiate this reaction. (The overvoltage for the oxidation of water can be as large as 1 volt.) By carefully choosing the electrode to maximize the overvoltage for the oxidation of water and then carefully controlling the potential at which the cell operates, we can ensure that only chlorine is produced in this reaction.

In summary, electrolysis of aqueous solutions of sodium chloride do not give the same products as the electrolysis of molten sodium chloride. Electrolysis of molten NaCl decomposes this compound into its elements:

2NaCl (l) 2Na (l) + Cl 2 (g)

Electrolysis of aqueous NaCl solutions gives a mixture of hydrogen and chlorine gas and an aqueous sodium hydroxide solution.

Negative electrode (cathode)

2H2O + 2e- H2 + 2OH-

Positive electrode (anode)

2Cl - Cl 2 + 2e-

The overall equation is

2NaCl (aq) + 2H 2O (l) 2Na+ (aq) + 2OH- (aq) + H2 (g) + Cl2 (g)

Because the demand for chlorine is much larger than the demand for sodium, electrolysis of aqueous sodium chloride is a more important process commercially. Electrolysis of an aqueous NaCl solution has two other advantages. It produces H2 gas at the cathode, which can be collected and sold. It also produces NaOH, which is the reason that this process forms the basis of the industrial production of sodium hydroxide.



A video on the electrolysis of sodium chloride solution or brine.

 

Properties and Uses

The ionic lattice structure and
appearance of sodium hydroxide.

Sodium hydroxide (NaOH), also known as lye or caustic soda, is a caustic metallic base. It is widely used in industry, mostly as a strong chemical base in the manufacture of pulp and paper, textiles, drinking water, and detergents. Worldwide production in 1998 was around 45 million tonnes. Sodium hydroxide is also the most common base used in chemical laboratories.

Pure sodium hydroxide is a white solid, available in pellets, flakes, granules and also 50% saturated solution. It is very deliquescent and also readily absorbs carbon dioxide from the air, so it should be stored in an airtight container. It is very soluble in water with liberation of heat.

Sodium hydroxide also dissolves in ethanol and methanol, though it exhibits lower solubility in these solvents than does potassium hydroxide. It is insoluble in ether and other non-polar solvents. A sodium hydroxide solution will leave a yellow stain on fabric and paper.

Industrial Production

Industrially, sodium hydroxide is produced by electrolysing seawater or brine. Brine is the only feedstock used in this process. In 1998, total world production was around 45 million tonnes. Of this, both North America and Asia contributed around 14 million tonnes, and Europe produced around 10 million tonnes.

Brine Purification

Aqueous brine (NaCl) must be purified before it can be used for sodium hydroxide production. The calcium, magnesium and sulfate ions are removed by precipitation.

The calcium ions are removed by precipitation with aqueous sodium carbonate:

Na2CO3 (aq) + CaCl2 (aq) CaCO3 (s) + 2NaCl (aq)

The magnesium ions are removed by precipitation with aqueous sodium hydroxide:

2NaOH (aq) + MgCl2 (aq) Mg(OH)2 (s) + 2NaCl (aq)

The sulfate ions are removed by precipitation with aqueous barium chloride:

BaCl2 (aq) + Na2SO 4 (aq) BaSO4 (s) + 2NaCl (aq)

In the production sodium hydroxide, both the diaphragm and the mercury amalgam processes are increasingly being replaced by membrane technology that consumes less energy and is more environmentally benign.

The use of membranes for chlor-alkali electrolysis has introduced requirements for much stricter purity control of the brine. The presence of impurities such as Ca2+, Mg2+, Sr2+, Ba2+, Al3+, SiO2, SO42-, and I - can shorten the lifetime of the membranes or can damage the electrodes. This results in a higher consumption of energy and higher membrane replacement cost. The contaminants are brought into the system by salt, dilution water and chemicals used in the process.

Typically, the brine purification steps include: saturation, precipitation (shown above), clarification, filtration, selective ion exchange, electrolysis and dechlorination. If the salt is already of high purity, such as vacuum salt, a primary purification with precipitation-filtration is not necessary and a secondary purification with ion exchange only is sufficient.

Methods of Production

Sodium hydroxide is produced (along with chlorine and hydrogen) via the chloralkali process. This involves the electrolysis of an aqueous solution of sodium chloride. The sodium hydroxide builds up at the cathode, where water is reduced to hydrogen gas and hydroxide ion:

2NaCl (aq) + 2H2O (l) 2Na+ (aq) + 2OH- (aq) + H2 (g) + Cl2 (g)

To produce NaOH it is necessary to prevent reaction of the NaOH with the chlorine. This is typically done in one of three ways, of which the membrane cell process is economically and environmentally the most viable.

Diaphragm Cell Process

This process uses a steel cathode and reaction of NaOH with Cl2 is prevented using a porous diaphragm.

The diaphragm process processes sodium hydroxide at a concentration of 50% w/v
and with about 2% impurities.

In this the oldest process, the diaphragm was traditionally made from asbestos but now made from a substance called PMX. Both compartments contain brine. The cathode is steel mesh and the anode is graphite or titanium coated with ruthenium-titanium oxide. The voltage across the electrodes is typically 3.5 - 5 V. Production plants draw huge currents at such voltages in the order of tens of thousands of amperes. Sodium ions migrate from the anode compartment through the diaphragm to the cathode compartment to neutralise the charge from the hydroxide ions formed there. The diaphragm also allows some diffusion of hydroxide ions and chloride ions across the membrane.

The solution formed in the cathode compartment contains both NaOH and NaCl and is further purified by the crystallisation of sodium chloride. The final solution is about 50% w/v NaOH with a salt contamination of about 2%.

Anode (+) oxidation reaction:

2Cl- Cl 2 + 2e -

Cathode (-) reaction:

2H 2O + 2e- H2 + 2OH-

Overall cell reaction:

2NaCl (aq) + 2H2O (l) 2NaOH (aq) + H2 (g) + Cl2 (g)

This process produces large quantities of two high demand industrial chemicals at quite reasonable costs and acceptable levels of purity for most purposes. The disadvantages are:

  • there is always a small amount of chloride in the NaOH produced

  • there are health and environmental problems associated with small losses of asbestos in making and using the diaphragms

  • there may be hypochlorite (OCl-), a strong oxidant in the waste brine solution and this needs to be removed before brine is discharged into the environment.

Diaphragm cells are now old technology and new plants are being built which utilise the membrane process. The mercury process was developed next to increase the purity of the sodium hydroxide product.



A video showing how sodium hydroxide is made using the diaphragm process using an asbestos diaphragm.

 

Mercury Cell Process

In this process, sodium metal forms as an amalgam at a mercury cathode which is then reacted with water to produce NaOH. There have been concerns about mercury releases, although modern plants claim to be safe in this regard.

The mercury process produces sodium hydroxide at a concentration of 50% w/v
and with virtually no impurities.

In this process, liquid mercury forms the cathode and sodium ions are oxidised to sodium metal there. The sodium rapidly alloys with the mercury to form an amalgam. At the voltages used, the sodium is prevented from reacting with the brine. The sodium is later removed from the mercury by reaction with water to produce NaOH and H2. The mercury can then be recycled. This process produces much higher purity sodium hydroxide than the diaphragm process and the product contains no chloride ions. About 50% w/v solutions are formed.

Anode (+) oxidation reaction:

2Cl- Cl 2 + 2e -

Cathode (-) reaction:

Na+ + e - → Na

Overall cell reaction:

2NaCl 2Na + Cl 2

Reaction of sodium with water:

2Na + 2H2O → 2NaOH + H2

The advantages of the mercury cell is that it produces very pure sodium hydroxide and avoids using asbestos. The main disadvantage is that the cells lose mercury which can be discharged into the environment with the spent brine. Mercury is a heavy metal and discharging it into aquatic environments can lead to bioaccumulation in the food chain. Mercury is toxic, affecting the central nervous system and eventually leading to brain damage. Environmental authorities in many countries have set targets for mercury output in the order less than 1 g per tonne of NaOH released. These targets are difficult to meet and few new mercury plants are being built as a consequence.

Membrane Cell Process

This process is similar to the diaphragm cell process but uses a modified Teflon membrane to separate the cathode and anode reactions. It is less expensive than the diaphragm cell process, and it produces a higher quality of NaOH.

The membrane process processes sodium hydroxide at a concentration of 50% w/v,
with about 0.02% impurities and the lowest environmental risk.

The anode compartment contains strong brine solution and the cathode compartment contains dilute sodium hydroxide solution. The two compartments are separated by an ion selective membrane made from a modified Teflon, polytetrafluoroethylene (PTFE). The cathode is made of steel mesh and the anode is titanium coated with ruthenium - titanium oxide. The NaOH produced is very pure because of the use of the ion-selective membrane. The final solution is about 50% w/v with a salt contamination of about 0.02%.

The reactions are the same as those for the diaphragm process.

Anode (+) oxidation reaction:

2Cl- Cl 2 + 2e -

Cathode (-) reaction:

2H 2O + 2e- H2 + 2OH-

Overall cell reaction:

2NaCl (aq) + 2H2O (l) 2NaOH (aq) + H2 (g) + Cl2 (g)

In recent years all new chloro-alkali plants have been of the membrane cell type. This is an example of how advances in chemistry lead to changes in technology; the availability of new materials led to the development of a superior diaphragm, the ion exchanging PTFE membrane that overcame the disadvantages of the old asbestos diaphragm cell without introducing any new problems. Earlier the mercury cell had been introduced to avoid the disadvantages of the diaphragm cell but it introduced another problem, namely the loss of mercury from the cell into the environment. The development of the membrane cell has meant that in recent years all new chlor-alkali plants have been of the membrane-cell type. Orica, the large Australian chemicals manufacturer, recently replaced its mercury-based chlor-alkali plant on Botany Bay with a plant using membrane cells.