- Electrolysis reaction
- Half-cell reactions
- Process
- Techniques
- Electrolysis with alkaline water
- Electrolysis with polymer electrolytic membrane
- Electrolysis with solid oxides
- What is the electrolysis of water for?
- Hydrogen production and its uses
- As a debugging method
- As an oxygen supply
- Home experiment
- Home variables
- References
The electrolysis of water is the decomposition of water into its elemental components by applying an electrical current. As they proceed, hydrogen and molecular oxygen, H 2 and O 2, are formed on two inert surfaces. These two surfaces are better known by the name of electrodes.
Theoretically, the volume of H 2 formed should be twice the volume of O 2. Why? Because the water molecule has an H / O ratio equal to 2, that is, two H for each oxygen. This relationship is directly verified with its chemical formula, H 2 O. However, many experimental factors influence the volumes obtained.
Source: Antti T. Nissinen via Flickr
If the electrolysis is carried out inside tubes immersed in water (upper image), the lower column of water corresponds to hydrogen, since there is a greater amount of gas exerting pressure on the surface of the liquid. The bubbles surround the electrodes and end up rising after overcoming the vapor pressure of the water.
Note that the tubes are separated from each other in such a way that there is a low migration of gases from one electrode to the other. At low scales, this does not represent an imminent risk; but on industrial scales, the gaseous mixture of H 2 and O 2 is highly dangerous and explosive.
For this reason, the electrochemical cells where the electrolysis of water is carried out are very expensive; They need a design and elements that guarantee that gases never mix, a profitable current supply, high concentrations of electrolytes, special electrodes (electrocatalysts), and mechanisms to store the H 2 produced.
Electrocatalysts represent friction and at the same time wings for the profitability of water electrolysis. Some consist of noble metal oxides, such as platinum and iridium, whose prices are very high. It is at this point especially where researchers join forces to design efficient, stable and cheap electrodes.
The reason for these efforts is to accelerate the formation of O 2, which occurs at lower rates compared to H 2. This slowing down by the electrode where O 2 is formed brings as a general consequence the application of a potential much higher than necessary (overpotential); which is equal, to a lower performance and higher expenses.
Electrolysis reaction
The electrolysis of water involves many complex aspects. However, in general terms, its basis rests on a simple global reaction:
2H 2 O (l) => 2H 2 (g) + O 2 (g)
As seen in the equation, two water molecules are involved: one usually must be reduced, or gain electrons, while the other must oxidize or lose electrons.
H 2 is a product of the reduction of water, because the gain of electrons promotes that the H + protons can bind covalently, and that the oxygen is transformed into OH -. Therefore, H 2 is produced at the cathode, which is the electrode where the reduction occurs.
While O 2 comes from the oxidation of water, due to which it loses the electrons that allow it to bind to hydrogen, and consequently releases H + protons. O 2 is produced at the anode, the electrode where oxidation occurs; And unlike the other electrode, the pH around the anode is acidic and not basic.
Half-cell reactions
This can be summarized with the following chemical equations for half-cell reactions:
2H 2 O + 2e - => H 2 + 2OH - (Cathode, basic)
2H 2 O => O 2 + 4H + + 4e - (Anode, acid)
However, water cannot lose more electrons (4e -) than the other water molecule gains at the cathode (2e -); therefore, the first equation must be multiplied by 2, and then subtracted with the second equation to get the net equation:
2 (2H 2 O + 2e - => H 2 + 2OH -)
2H 2 O => O 2 + 4H + + 4e -
6H 2 O => 2H 2 + O 2 + 4H + + 4OH -
But 4H + and 4OH - form 4H 2 O, so they eliminate four of the six H 2 O molecules, leaving two; and the result is the global reaction just outlined.
Half-cell reactions change with pH values, techniques, and also have associated reduction or oxidation potentials, which determine how much current needs to be supplied for the electrolysis of water to proceed spontaneously.
Process
Source: Ivan Akira, from Wikimedia Commons
A Hoffman voltameter is shown in the image above. The cylinders are filled with water and the selected electrolytes through the middle nozzle. The role of these electrolytes is to increase the conductivity of the water, since under normal conditions there are very few H 3 O + and OH ions - products of their self-ionization.
The two electrodes are usually made of platinum, although in the image they were replaced by carbon electrodes. Both are connected to a battery, with which a potential difference (ΔV) is applied that promotes the oxidation of water (formation of O 2).
The electrons travel the entire circuit until they reach the other electrode, where the water wins them over and becomes H 2 and OH -. At this point, the anode and cathode have already been defined, which can be differentiated by the height of the water columns; the one with the lowest height corresponds to the cathode, where H 2 is formed.
In the upper part of the cylinders, there are keys that allow the gases generated to be released. The presence of H 2 can be carefully checked by reacting it with a flame, the combustion of which produces gaseous water.
Techniques
Water electrolysis techniques vary depending on the amount of H 2 and O 2 to be generated. Both gases are very dangerous if they are mixed together, which is why electrolytic cells involve complex designs to minimize the increase in gaseous pressures and their diffusion through the aqueous medium.
Also, techniques vary depending on the cell, the electrolyte added to the water, and the electrodes themselves. On the other hand, some imply that the reaction is carried out at higher temperatures, reducing electricity consumption, and others use enormous pressures to keep the H 2 stored.
Among all the techniques, the following three can be mentioned:
Electrolysis with alkaline water
Electrolysis is carried out with basic solutions of the alkali metals (KOH or NaOH). With this technique the reactions occur:
4H 2 O (l) + 4e - => 2H 2 (g) + 4OH - (aq)
4OH - (aq) => O 2 (g) + 2H 2 O (l) + 4e -
As can be seen, both at the cathode and at the anode, water has a basic pH; and in addition, the OH - migrate towards the anode where they are oxidized to O 2.
Electrolysis with polymer electrolytic membrane
In this technique a solid polymer is used that serves as a membrane that is permeable for H +, but impermeable for gases. This ensures greater safety during electrolysis.
The half-cell reactions for this case are:
4H + (aq) + 4e - => 2H 2 (g)
2H 2 O (l) => O 2 (g) + 4H + (aq) + 4e -
H + ions migrate from the anode to the cathode, where they are reduced to become H 2.
Electrolysis with solid oxides
Very different from the other techniques, this one uses oxides as electrolytes, which at high temperatures (600-900ºC) function as a transport medium for the O 2- anion.
The reactions are:
2H 2 O (g) + 4e - => 2H 2 (g) + 2O 2-
2O 2- => O 2 (g) + 4e -
Note that this time it is the oxide anions, O 2-, that travel to the anode.
What is the electrolysis of water for?
The electrolysis of water produces H 2 (g) and O 2 (g). Approximately 5% of the hydrogen gas produced in the world is made through the electrolysis of water.
H 2 is a by-product of the electrolysis of aqueous NaCl solutions. The presence of salt facilitates electrolysis by increasing the electrical conductivity of the water.
The overall reaction that takes place is:
2NaCl + 2H 2 O => Cl 2 + H 2 + 2NaOH
To understand the enormous importance of this reaction, some of the uses of gaseous products will be mentioned; Because at the end of the day, these are the ones that drive the development of new methods to achieve the electrolysis of water in a more efficient and green way.
Of all of them, the most desired is to function as cells that energetically replace the use of burning fossil fuels.
Hydrogen production and its uses
-Hydrogen produced in electrolysis can be used in the chemical industry acting in addiction reactions, in hydrogenation processes or as a reducing agent in reduction processes.
-It is also essential in some actions of commercial importance, such as: the production of hydrochloric acid, hydrogen peroxide, hydroxylamines, etc. It is involved in the synthesis of ammonia through a catalytic reaction with nitrogen.
-In combination with oxygen, it produces flames with a high caloric content, with temperatures ranging between 3,000 and 3,500 K. These temperatures can be used for cutting and welding in the metal industry, for growth of synthetic crystals, quartz production, etc..
-Water treatment: excessively high nitrate content in water can be reduced by its elimination in bioreactors, in which bacteria use hydrogen as an energy source
-Hydrogen is involved in the synthesis of plastics, polyester and nylon. In addition, it is part of the production of glass, increasing combustion during baking.
-Reacts with the oxides and chloride of many metals, among them: silver, copper, lead, bismuth and mercury to produce pure metals.
-And additionally, it is used as fuel in chromatographic analysis with a flame detector.
As a debugging method
The electrolysis of sodium chloride solutions is used for the purification of swimming pool water. During electrolysis, hydrogen is produced at the cathode and chlorine (Cl 2) at the anode. Electrolysis is referred to in this case as a salt chlorinator.
Chlorine dissolves in water to form hypochlorous acid and sodium hypochlorite. Hypochlorous acid and sodium hypochlorite sterilize water.
As an oxygen supply
The electrolysis of water is also used to generate oxygen on the International Space Station, which serves to maintain an oxygen atmosphere at the station.
Hydrogen can be used in a fuel cell, a method of storing energy, and use the water that is generated in the cell for consumption by astronauts.
Home experiment
Water electrolysis experiments have been carried out at laboratory scales with Hoffman voltmeters, or another assembly that allows to contain all the necessary elements of an electrochemical cell.
Of all possible assemblies and equipment, the simplest may be a large transparent water container, which will serve as a cell. In addition to this, any metal or electrically conductive surface must also be on hand to function as electrodes; one for the cathode, and the other for the anode.
For this purpose even pencils with sharp graphite tips at both ends can be useful. And finally, a small battery and some cables that connect it to the improvised electrodes.
If not carried out in a transparent container, the formation of gaseous bubbles would not be appreciated.
Home variables
Although the electrolysis of water is a subject that contains many intriguing and hopeful aspects for those seeking alternative energy sources, the home experiment can be boring for children and other bystanders.
Therefore, sufficient voltage can be applied to generate the formation of H 2 and O 2 by alternating certain variables and noting the changes.
The first one is the variation of the pH of the water, using either vinegar to acidify the water, or Na 2 CO 3 to slightly basify it. A change in the number of bubbles observed must occur.
Additionally, the same experiment could be repeated with hot and cold water. In this way, the effect of temperature on the reaction would then be considered.
Finally, to make the data collection a little less colorless, you can use a very dilute solution of purple cabbage juice. This juice is an acid-base indicator of natural origin.
Adding it to the container with the electrodes inserted, it will be noted that at the anode the water will turn pink (acid), while at the cathode, the color will be yellow (basic).
References
- Wikipedia. (2018). Electrolysis of water. Recovered from: en.wikipedia.org
- Chaplin M. (November 16, 2018). Electrolysis of water. Water structure and science. Recovered from: 1.lsbu.ac.uk
- Energy Efficiency & Renewable Energy. (sf). Hydrogen production: electrolysis. Recovered from: energy.gov
- Phys.org. (February 14, 2018). High-efficiency, low-cost catalyst for water electrolysis. Recovered from: phys.org
- Chemistry LibreTexts. (June 18, 2015). Electrolysis of water. Recovered from: chem.libretexts.org
- Xiang C., M. Papadantonakisab K., and S. Lewis N. (2016). Principles and implementations of electrolysis systems for water splitting. The Royal Society of Chemistry.
- Regents of the University of Minnesota. (2018). Electrolysis of Water 2. University of Minnesota. Recovered from: chem.umn.edu