paper1-1

On The Production

of Hydrogen Gas

By The Electrolysis

of Water

Johnston

Page 4

To summarize up to this point; we have seen that the energy to separate the water molecules in our solution of electrolyte and water comes from the interaction between the molecules of water and the molecules of the electrolyte and occurs BEFORE any electric current is created or added. We have seen that, for every electron that enters the cell or battery at the cathode, another electron from the solution must EXIT at the anode. If these two reactions do not happen, then all reactions stop.

This finding is the basis of Faraday's law of electrolytic conduction which states; "The amount of gasses evolved is equal to the amount of current which passes". It is obvious at this point that it cannot be any other way because conduction of electricity from one side of a cell to the other cannot take place unless H2 gas is evolved by reduction at the cathode and O2 by oxidation at the anode, releasing a quantity of energy that is EQUAL to the amount of that is "used" to evolve H2 at the cathode. Equilibrium must be maintained. My trusty chemistry book describes the process of electrolysis in the following way; "In order to maintain an electric current, it is necessary to have a complete circuit, i.e., there must be a closed loop whereby the electric charge can return to it's starting point. If the complete circuit includes as one component an electrolytic conductor, chemical reaction must occur."

"Electrons are crowded onto the cathode and drained away at the anode. The circuit is not complete unless there is some way by which electrons can be used up at the cathode and formed at the anode. Chemical changes must occur."

"At the cathode, a reduction process must occur in which some ion or molecule accepts electrons and is thereby reduced. At the anode, electrons must be released by an ion or molecule to the electrode. An oxidation process must occur."

That this is true is obvious to anyone who has ever constructed a simple electrolysis cell. The "problem" then is that, when you operate an electrolysis cell, you are, by the very nature of the reaction, releasing from the cell double the amount of energy which actually enters the cell in the form of electric current. This input energy is transformed into and "stored" as the potential chemical energy of H2(g) and as a consequence of this reaction, to maintain chemical equilibrium within the solution, an equal amount of electric current is released at the anode. Faraday explained this phenomenon by calling the gasses that are produced a secondary effect (by-product) of the electrolytic conduction process.

If we consider again the remaining reactants after a complete cycle of electrolysis, in such a cell as the example I have been using up to this point, we find that we are left with our initial SO3 molecule which has now given up it's extra oxygen atom. What does it do now? Simple, it breaks up ANOTHER H2O molecule by becoming H2SO4 ( 2H+[SO4--] ) again, just as it did initially before any current was supplied. Or, in the case of sodium sulfate ( NaSo4 or; Na+[SO4 -2] ) the products that remain after electrolyzation are the same products that began the reaction, only water disappears (is consumed).

After considering the data contained in this paper I have concluded that the reactions which occur in a primary (voltaic) cell and the reactions that occur in an electrolysis cell are identical except for the fact that in the primary cell the anode is oxidized and decomposes and in the electrolysis cell (using electrodes which are chemically inert to the action of the electrolyte) the oxygen itself is oxidized. Here then, is the explanation of the seeming "surplus" energy that we get from our electrolysis cell. Since the reactions in both the battery and the cell are identical, then, by artificially creating a potential difference between the inert electrodes with the supplied current, we are causing the cell to function as a battery with the oxidation/reduction reaction proceeding in reverse. However, since no actual current passes through the battery the supplied electricity is instead stored up in the H2(g) that is released. So therefore you effectively double your supplied energy by using that supplied energy to stimulate the same set of reactions which result in the release of energy in a primary cell.

That this is just a coincidental effect of the normal operation of the electrolysis cell can be demonstrated by considering the results that could be expected if the cathode was made of a material that could itself be reduced by combining with the H+ ions, and the anode of a material which could be oxidized by the SO4 -2 ions. If that was the case the results might look like this;

H2SO4-----------------> H2M(s) + ZnSO4(s)

electrolysis

Each individual reaction would look like this;

Cathode: 2H+ and M(s)----------->H2M(s) -135kcal

Anode: Zn(s) + SO4 -2--------------->ZnSO4 +135kcal

In this case the process and the resulting products would still follow Faraday's laws but there would probably be no opportunity to extract further energy out of the by-products of this system because the hydrogen and oxygen, which would otherwise be released, would now both be bound into other molecules. Please note though that a quantity of electricity still exits the cell at the anode which is equal to the quantity which enters at the cathode and again, this is as it must be to preserve chemical equilibrium within the solution.

As another example consider the secondary cell, such as is used for car batteries. In this type of battery there is no current produced, the input current is merely stored as potential chemical energy by causing a set of chemical reactions to proceed in one direction and later released by allowing the same set of reactions to reverse themselves. However if such a battery is fully charged and more electricity is passed through it H2O will be electrolyzed. Why? Because, after being fully charged, the chemical reactions which charged it cease and so, at that time, a secondary cell resembles, on the one hand, an electrolysis cell in that it's electrolyte will no longer attack the electrodes and instead the same set of reactions between the water, the electrolyte and the current takes place as detailed above in an electrolysis cell and NOW the secondary cell, which cannot normally produce electricity, is acting like a primary cell because, given the fact that the above set of reactions are taking place, it is producing an amount of electricity at the anode which is equal to the amount of current that is entering at the cathode by the same set of reactions, and again, only water is being consumed. Therefore it must be the splitting of the water itself which is generating the extra energy. This too follows accepted fact because we have seen that the two electrons from the H2 portion of the water molecule are first passed to the oxygen atom and then to the anode in a primary cell.

My conclusion then is that this is not a matter of "it might work" or this is "theoretically possible" it is very simply HOW it works and has always worked, since the time of Volta and Faraday. All energies being consumed and produced are fully accounted for and so the requirements for the laws that I quoted at the beginning of this paper are being met. And yet the net result is that you are getting twice as much energy out of the cell as what you put into it.

The next area which I plan to consider is resistance between the electrodes and the solution and within the solution itself. From the above information it would seem that there are several ways in which such resistance could be increased or decreased to meet the needs of particular applications and some preliminary experimental evidence from my own research seems to validate this. My next then paper will deal with the resistance of electrolysis cells and how to control their conductivity to achieve the desired resistance effect in multiple "series" cells.

END

BIBLIOGRAPHY

(1) Chemistry

By Michell J. Sienko

Associate Professor of Chemistry

Cornell university

And

Robert A. Plane

Assistant Professor of Chemistry

Cornell university

(2) Experimental Researches in Electricity

By Michael Faraday

Originally published in 1839

(3) Electricity One-Seven

By Harry Mileaf, Editor-in-Chief

(4) College Physics

By Robert L Weber, Marsh W. white and Kenneth V. Manning

All from the Pennsylvania State University