11. The use of common salt, sodium chloride, as an electrolyte
Common salt dissolved in water has been used as an electrolyte to carry out electro-etching. It is cheap, easy to handle and can be used to etch copper, zinc, aluminium and steel (iron). However, the process becomes a little more complex as now gases - hydrogen, oxygen and chlorine – may be produced at the electrodes and “by-products” formed within the electrolyte. Since hydrogen and chlorine gases and some of the by products produced may be hazardous when present in sufficient quantity and concentration it is important to have some idea of what is taking place in the electro-etching cell. Fortunately some guidance may be obtained from the related techniques of electro-chlorination of drinking water and electro-coagulation/flocculation of wastewaters.
11.1 Ions in the aqueous sodium chloride electrolyte available for reactions
Sodium chloride readily ionises to form Na+ and Cl− ions while water, although a very poor electrical conductor, ionises one molecule in every ten million, splitting into a positive hydrogen ion H+ and a negative hydroxyl ion OH−. Provided there are no dissolved impurities, such a solution should be neutral, pH = 7.
Commonly, at the cathode, irrespective of its composition, the preferred primary reaction is the evolution of hydrogen over the deposition of sodium since the latter is a more non-spontaneous reaction .
At the anode both the chlorine ion and the hydroxyl ion or water molecules could give up their electrons and evolve as chlorine or oxygen gases respectively, see appendix A section A2.4. What actually happens depends on the nature of the metal used for the electrode: an inert electrode for electrolysis or a “dissolving “ electrode for etching.
11.2 Inert electrodes
The characteristic of inert electrodes is that they do not take part in the electrochemical or chemical reactions that occur in the cell: their role is to provide a pathway for electrons to flow out of the electrolyte into the external circuit at the anode and to flow into the electrolyte from the external circuit at the cathode. Typically, an inert electrode may be made from titanium metal covered with a thin platinum coating.
The electrons at the anode electrode are provided by the discharging of negative ions from the electrolyte at the surface of the electrode to form molecules of oxygen or chlorine gas.
Yellow cuprous oxide produced in
sodium chloride electrolyte
after electro-etching a copper plate
For the electrolysis of sodium chloride solution (brine) it is found that chlorine is produced at high salt concentrations, oxygen at low concentrations and both at intermediate values To maintain electro-neutrality in the electrolyte, sodium and hydroxyl ions in the solution react to form sodium hydroxide (caustic soda). Further, if chlorine gas is evolved at the anode and mixed with the electrolyte from the cathode region , it can react with the sodium hydroxide to form sodium hypochlorite (sodium chlorate (I)), a constituent of household bleach. It is informative to estimate the quantities of the various molecules formed for an electrical current and time typical for an etching operation, see below.
Assessment of the amount of hydrogen H2, chlorine Cl2 and sodium hydroxide NaOH produced in the electrolysis of brine.
Established equations for the reactions occurring at the anode, cathode and in the electrolyte for the electrolysis of brine can be found in standard texts. These enable the quantities of hydrogen, oxygen and sodium hydroxide produced when 1 mole of electrons passes through the system - the cell and the external circuit.
To find the amounts produced when a current of ,say, 1 ampere passes for one hour it is only necessary to find the number of moles of electrons , contained in this amount of charge.
It can be deduced that one ampere of electrical current flowing for one hour produces:
3.7 10-2 g hydrogen , equal to 0.42 L in gaseous form @STP
1.3 g chlorine, equal to 0.42 L in gaseous form @ STP;
1.48 g sodium hydroxideSince the quantities produced are proportional to the current (amperes) multiplied by time (hours), values for other currents and times can be found on a pro-rata basis.
By extending this procedure for the situation where all the chlorine gas produced mixes with the sodium hydroxide to produce sodium hypochlorite NaOCl, it can be shown that the corresponding quantity for 1 ampere and 1 hour is 1.4g sodium hypochlorite.
For 1 litre of electrolyte, and assuming that the sodium hydroxide or sodium hypochlorite is well mixed within it, the corresponding molar concentrations would be approximately 4.10-2 M and 2 .10-2M respectively.
Standard temperature and pressure: 1 atmosphere, 0 ⁰C
11.3 “Active” Electrodes for etching –copper, zinc aluminium and iron (mild steel)
11.3.1 Difference between ”active” and inert anode electrodes
The anode will be made from whatever metal we wish to etch; it will be covered with a resist into which an image has been drawn which exposes the metal surface at these places to the electrolyte. In contrast to using an inert anode, an anode made from our etching metal, say zinc, does take part in the electrochemical reactions where it is in contact with the electrolyte: its atoms give up electrons to the external circuit and “dissolve” into the electrolyte as positive ions. Remember, the rate at which this dissolution or etching takes place depends only on current and etching time: in practice the reduced surface area available for etching due to the resist may influence the actual values of current which can be obtained but nothing else.
11.3.2 Discharge of gases at the anode (etched plate)
What about the chlorine and hydroxyl ions or water molecules discharging at the anode? In principle these non-spontaneous gaseous reactions at the anode can still occur but now they are in competition with the dissolution of the metal anode itself - aluminium, iron and zinc metal plates (spontaneous reactions) and copper (less non-spontaneous than the gaseous reactions). Consequently, dissolution of the metal anodes rather than evolution of chlorine or oxygen is the preferred reaction. The evolution of chlorine gas may be negligible and certainly smaller in comparison with brine electrolysis (correspondingly the formation of sodium hypochlorite, if any, should be smaller).
The amount of hydrogen which is produced at the cathode is the same as for brine electrolysis. Additionally, for aluminium plates, a parasitic reaction (independent of the current) between the aluminium anode and the water of the electrolyte may result in some additional hydrogen.
A cautious approach, that provides an upper estimate of the amount of chlorine which could be produced, is to assume that the quantity is the same as for the brine electrolysis, although this should not be the case. Information obtained from the field of electro-coagulation/flocculation for wastewater treatment – technology which uses the insoluble metal hydroxides formed by the dissolution of aluminium anodes to trap waste – suggests that excessive voltages are required before chlorine or oxygen gases are produced external to the cell. In a stirred electrolyte, the chlorine molecules can be consumed in the formation of sodium hypochlorite as already discussed. What exactly is meant by “excessive” is not stated but working voltages appear to be typically greater than 5 volts. Available electro-chlorination systems are able to retain any chlorine molecules produced within the electrolyte where they react to form hypochlorite, (see appendix B)
Neither hydrogen nor oxygen dissolve readily in water and so bubbles of both would be seen if they are produced. In contrast, at room temperatures, chlorine can be absorbed by water, around 4g/litre depending on the concentration of the sodium chloride solution. Furthermore, as described any chlorine gas formed at the anode region may react with the sodium hydroxide. As result, even if chlorine gas is produced, for a period of time it can be retained within the electrolyte solution; clearly the larger the volume of the electrolyte, the more effective this will be. The presence of chlorine may be detected by the nose at low concentrations , not that this can be relied on as a means of sensing its presence.
11.3.3 By products produced in the electrolyte
In addition to the sodium hydroxide, other products may be formed in the electrolyte. Secondary reactions between the “dissolved” metallic ions from the anode and the various ions in the electrolyte may follow a number of intermediate routes involving the formation of more complex ions but typically the end result is the formation of insoluble metal hydroxides, metal oxides or metal chlorides depending on the current used, etching time and initial condition of the electrolytic solution. For example, the etching of copper in sodium chloride solution can lead to a rich display of electrolyte colours as the secondary copper compounds change from copper(II) chloride CuCl2- a soluble blue/green foam -to copper hydroxide Cu(OH)2 – an insoluble pale blue gelatinous solid –to cuprous oxide Cu2O – insoluble yellow or brick red particulates depending on the particle size.
Some of these by-products in concentrated form are classified as “harmful”; whether or not they represent a real safety or disposal issue will largely depend on the concentration produced. A useful guide in such matters, “Student Safety Sheets”, can be obtained from CLEAPSS® at www.cleapss.org.uk
Compared to the electrolysis of brine with inert electrodes, the formation of sodium hydroxide may be less due to the competition from the dissolving metallic ions for the hydroxyl ions; the possible formation of sodium hypochlorite should be smaller given that there may be negligible chlorine produced. Again, a cautious upper estimate may use the values found from the electrolysis of brine. It is seen that the amounts formed in a single etching operation are small; their possible cumulative effect after many etchings should be considered. The presence of these compounds has consequences for the safe and environmentally acceptable disposal of the spent electrolyte.
If required, an estimate of the possible by-products formed from the dissolving metallic ions may be obtained without attempting to follow the detailed chemistry. First, the amount of “dissolved ”metal ions can be found from electrode reactions, and then for the assumed end product, the quantity can be found using its chemical formula. For example, one ampere passing for one hour produces 1.2 g of copper Cu. If the final product is copper chloride CuCl2, the estimated amount is 2.5g; if it is copper hydroxide Cu(OH)2 the quantity is 1.9g. Similarly the corresponding masses for etching aluminium would be: 0.3g of aluminium Al and 0.9g of aluminium hydroxide Al(OH)3.
11.4 A view on the practical use of sodium chloride solution as an electrolyte
The use of sodium chloride electrolyte for zinc, copper, iron and steel provides no obvious etching advantage over the simpler metal/same metal salt electrolyte combination. The etching of aluminium is somewhat different since there is no equivalent process using a same metal salt electrolyte system. In this case, it would appear that no matter what approach is used the composition of the electrolyte will change with use and eventually require replacement and disposal. The use of passive etching with saline sulphate/Bordeaux etch should not be forgotten and the possibility of using other salts, for example sodium carbonate, may be worth investigating.
Zinc plate after electro- etching
in sodium chloride electrolyte
For a teaching or open access studio it can be imagined that there would be sufficient demand for dedicated electro-etching units for each metal so the availability of a single electrolyte is of reduced benefit. However, the low cost, ease of handling, and availability of common salt in comparison with metal sulphates may be appealing for the individual etcher.
The generation of hydrogen, the possibility, however small, of producing chlorine gas and the forming of by-products requiring handling and disposal, indicate that a safety assessment should be made. Recently, some strong concerns have been expressed about using sodium chloride as an electrolyte. The issues involved are discussed in appendix B. For the moment the use of sodium chloride should only be contemplated by the experienced and informed electro-etcher who wishes to explore and evaluate the approach and is capable of making a risk assessment for their particular sodium chloride etching system and carrying out tests in a safe manner.