Water Electrolysis

Electrolysis cells are opposite of the galvanic cells. In galvanic cells, the cells provide power to drive electrolysis reactions while for the electrolysis cells energy is needed to cause the chemical reactions to occur, which under normal situations can not take place spontaneously.  The electrolysis process is very crucial in many applications such as purifying, plating and metals refining. For example, the well-known method for the use of electrolysis is the production of aluminum from bauxite(its ores). In some few years to come, oil will be a scarce product, and it is also the major contributor to environmental pollution. Therefore, this calls for an urgent need for a clean alternative source of energy.  Hydrogen has taken the center of discussion as the suitable source of energy to replace oil use in cars.

At the moment, hydrogen is used as fuel in rocket industry, in which hydrogen powers the engines of space shuttles. A significant amount of hydrogen required for the space shuttle is produced from the natural gas which is also a non-renewable source of energy. However, the large quantities of hydrogen that can meet the world hydrogen demand can only be obtained from the electrolysis. The production of hydrogen via electrolysis is possible because the electrolysis will require a cheap source of electricity, like nuclear or solar power.

For the electrolysis of water to occur, it should be conductive than pure water([HO₃⁺] and [^-OH] of clean water are only 10^-7 M).The more conducive of water for electrolysis is achieved by adding a salt or acid. For demonstration in this article, 0.10M KNO₃(basic salt) is used to make water conductive.  A 9-volt battery is employed to provide power for the chemical reactions to occur. The two graphite electrodes, to which one is attached negative terminal of the battery and the other on positive terminal provide a non-reactive surface for the electrolysis reactions to take place. The electrolysis reactions that occur at anode and cathode are as follow:

2H₂O=O₂(g) +4H⁺ +4e^-   ……………………………………………(Oxidation) equation 1 

2 x (2 H₂O + 2e^- = H₂(g) + 2 OH^-)…………………..………………..Reduction) equation 2

Overall reactions(Equation 1 and 2):

6H₂O=2H₂(g) + O₂(g) +4H+ + 4^-OH

Net reaction:

2 H₂O= 2 H₂(g) + O₂(g)       (Whitten et al.)

Hence, the net electrolysis reaction is the production of the hydrogen gas and oxygen gas. Hydrogen gas is liberated at the cathode while oxygen gas is produced at the anode. As the hydrogen gas is produced at the cathode, the solution around the base of the electrode becomes basic. On the other hand, when oxygen is liberated at the anode, the solution at the base of the electrode becomes acidic. The K⁺ and NO₃^-  do not take part in the reaction because they are less readily reduced and oxidized as the water(H₂O). However, KNO₃ plays a role in the conduction of the current through the solution. 

How to carry out water electrolysis (Experiment)

As stated earlier, 9-volt battery, two graphite pencils(as the lead) and 0.10M potassium nitrate are used. Place 10ml of 0.1M potassium nitrate in the 50ml beaker, add eight drops of Yamada indicator. Connect the graphite, one graphite to the positive terminal and the other to the negative terminal of the battery. Dip the two graphite(electrodes) into the solution(0.1M potassium nitrate) in the beaker. Observe the color changes at each electrode. At the anode electrode, the color change of the Yamada is red while the color change at the cathode is purple. Alternatively, one can use litmus paper, by dipping the litmus paper at the base of each electrode and observing the color change. For instance, at the cathode, the litmus paper turns to blue while at the anode litmus paper changes to red. The gasses(hydrogen and oxygen) can be collected and applied in various areas.