On The Production

of Hydrogen Gas

By The Electrolysis

of Water

Johnston

In the case of a battery (Fig. 1) the current at the cathode is supplied by the decomposition of the zinc anode by oxidation, where the oxygen atom which was "stolen" from the water molecule by the SO3 molecule is transfered to a zinc atom, creating an imbalance in the ion concentration in the electrolyte and, as part of the transfer, releases the two electrons which originally belonged to the two H+ ions. These two electrons then create an electrical imbalance in the wire which connects the two electrodes and the H+ ions are then able to pick up two electrons at the surface of the cathode and unite with each other to become H2 gas. This maintains chemical equilibrium in the solution and electrical neutrality in the connecting wire and allows further decomposition to occur at the anode. The anode product of this oxidation is zinc sulfate.

Please notice that the net result here is the creation of an electric current. This current is the by-product of the reaction between zinc metal (Zn+2 ion) and the SO4 -2 ion. Zinc will not react (appreciably) with water alone or with SO3. The electrons which are transfered to the zinc anode when it is oxydized are "technically" the same two electrons that the SO4 -2 ion "stole" from the two H atoms in the water molecule which was split by reacting with the SO3. So here too the energy that we are getting out of the cell is being created by chemical reactions. This is the second set of reactions. The first was when the SO3 oxidized the water molecule and ended up with two extra electrons. The next reaction is when the SO4 -2 ion reacts with zinc and releases these same two electrons. The final step in this set of reactions is when the 2 electrons travel around the wire connecting the anode and cathode and become available at the surface of the cathode as free electrons. There they create a negative electromagnetic field which draws the H+ ions to them and the H+ ions are then able to join with these free electrons and ultimately with each other to form H2 gas. So you see the two H+ ions get their original electrons back and equilibrium is maintained.

The anode reaction would be: Zn(s) + SO4 -2 --------> ZnSO4(s) + e-2 or, if you looked at it on a level of mole quantities it would be;

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

and the cathode reaction would be: 2H+ plus e-2 (from the anode)------------> H2(g) or, again, if you looked at the same equation in mole quantities it would be;

2H+ plus 135kcal (from the anode reaction) -----------> H2(g) -135kcal

This set of equations is the necessary expression of the balanced set of reactions which undeniably DO occur during the operation of a voltaic cell and fit comfortably into the previously mentioned laws of energy transfer and transformation. One side is exothermic and the other endothermic and no energy is created or destroyed. Neither is the anode actually destroyed, it is just combined by oxidation into another compound and this is as expected since no matter is ever actually created or destroyed, just transformed. As we have seen, Faraday's laws describe this reaction in exactly this way.

Figure 3

The above drawing depicts an electrolysis cell of the type that Faraday might have used. It consists of two electrodes which are made of a metal which is inert to the oxydizing power of the electrolyte in the solution into which they are placed, an electrolytic solution of water and an electrolyte, and a power source which is connected between the electrodes.

In the case of an electrolysis cell the waiting electrons (at the cathode surface) are supplied by the current from some outside power source rather than coming from the decomposition of the anode (as in the previous example of the action in a voltaic cell). This is because the anode and cathode are made of the same inert metal and so no potential difference is created between the electrodes as it does in a primary cell where the electrodes are made of different metals, ideally at opposite ends of the electromotive series. Also because, in an electrolysis cell, there is no continuous wire connecting the anode to the cathode as there is in the primary cell/battery and hence no complete circuit from electrode to electrode. Instead the circuit goes from the external power source through the cell and back to the power source. The reaction in an electrolysis cell must then proceed in the reverse direction of the reaction in a voltaic cell and be considered as a reduction/oxidation reaction. In both cases the energy to break the H2O molecule apart is being supplied by the interaction between the electrolyte and the water BEFORE any electricity is created or added.

Please note that no electrical current actually "passes through" the battery or the electrolysis cell as it would in the case of a wire (metallic) conductor. In the battery it is released at the anode and is used up at the cathode, and in the electrolysis cell enters at the cathode, to create, or better, to "reconstitute" H2 atoms and at the anode oxygen oxydizes itself to form O2(g) and releases 2 electrons. So this process conforms to the law of conservation of energy and is therefore "acceptable" This process continues until the anode is totally decomposed (in the case of the battery). All of the energy being created at the anode must re-enter at the cathode to maintain chemical equilibrium in the solution within the cell. If the wire connecting the electrodes contains too much resistance for all of the electrons to re-enter the cell then the battery "weakens" as the concentration of H+ ions becomes too high and the decomposing action at the anode eventually stops.

In either the electrolysis cell or the battery, once the supplied current at the cathode has given the H+ ions the electrons that they need to be reconstituted/set free as gas, the remaining SO4 molecule (usually written as SO4-2) is now in possession of the two electrons left behind by the H+ ions when the sulfur in the SO3 molecule was oxidized by the oxygen in the water molecule. The obvious question at this point is; in a solution of H2SO4 and water, before current is either created (battery) or supplied (electrolysis), how is it possible to see the SO4 molecule as holding two "excess" electrons "yet"? In figure 2 I showed that the absence of the two H+ ions from the SO4 created a "complete" molecule in which all of the oxygen atoms and the sulfur atom had all of the holes in their valence shells filled and so the SO4 molecule should be happy, right? Right! Good question! And indeed this is the case as any chemical reaction seeks it's own equilibrium point and, once it reaches that point, all reaction stops. If that is the case then why don't the two ions recombine to form H2SO4 again? Because the SO4 is happy being SO4 and the H+ ions (protons) are not strong enough to take their electrons back.

Note: There are inter-ionic reactions, especially when using complex electrolytes or more than one electrolyte in the same solution, or when pure (distilled) water is not used.(mj)

As you can plainly see from this example, the creation of an electrolytic solution consisting of an electrolyte and water, by simply mixing the two together, is what causes the separation/breakup of the water molecules as part of the process of creating ions! NO electrical current is required for this to happen! It is the energy of the atoms themselves which powers the chemical reactions which break the hydrogen bond to oxygen and tear the water molecules apart. Michael Faraday commented on this very power as follows; "868. What, then, follows as a necessary consequence of this whole experiment? Why, this: that the chemical action upon 32.31 parts, or one equivalent of zinc, in this simple voltaic circle, was able to evolve such quantity of electricity in the form of a current, as, passing through water, should decompose 9 parts, or one equivalent of that substance: and considering the definite relations of electricity as developed in the preceeding parts of the present paper, the results, prove that the quantity of electricity which, being naturally associated with the particles of matter, gives them their combining power, is able, when thrown into a current, to separate those particles from their state of combination: or, in other words, that the electricity which decomposes, and that which is evolved by the decomposition of, a certain quantity of matter, are alike."

"869. The harmony which this theory of the definite evolution and the equivalent definite action of electricity introduces into the associated theories of definite proportions and electro-chemical affinity, is very great. According to it, the equivalent weights of bodies are simply those quantities of them which contain equal quantities of electricity, or have naturally equal electric powers; it being the ELECTRICITY which determines the equivalent number, because it determines the combining force."(2)(page 389, Paragraphs; 868 & 869)

From this information it should be obvious that energy is required to create the ions in an electrolytic solution and that this energy must come from the substances within the solution itself. The energy given off by the oxidation of the anode in a voltaic cell is not the reciprocal of this initial energy of creating the ions in the solution. It is a totally separate reaction because, as noted previously, the ions are formed before any electrical energy is added or created. And since part of the process of creating the ions requires the use of some type of activation energy then, to preserve chemical equilibrium, this amount of required energy should be analogous (equal to) to the amount of energy released when the products of this reaction are reunited.

In other words; the reaction expressed by the equation:

H2(g) + O2(g) ---------------> H2O(l) + 135kcal

combustion

represents the other half of the chain of reactions which initially separates the components of the water molecules:

SO3 + H20-------->2H+ and SO4-- (H2SO4)

then

Zn+2 and SO4 -2 --------> ZnSO4 and e -2

then

2H+ and e -2 -------> 2H2(g)

Or the attractive force between the sulfur atom and the oxygen atom minus the attractiive force between the oxygen atom and the hydrogen atoms. So the attractive force of the sulfur for the oxygen atom must be greater than the combination energy of 135kcal per mole which is released when hydrogen and oxygen are burned together. Not necessarily a great deal more though, just enough to make the oxygen from the water molecule be more attracted to the sulfur than it is to the hydrogen. Then, conversely, this same process must be true of the reaction where the "extra" oxygen atom is transferred to a zinc atom in a battery and releases 2 electrons. The attractive force between the extra oxygen atom in the SO4 molecule and the sulfur must be less than the attractive force between the oxygen atom and a zinc atom from the anode.

This would seem contradictory because, if the attractive force between the oxygen atom and the zinc atom is enough to pull the oxygen atom away from the sulfur (SO4) atom then it should be enough to pull the oxygen atom directly out of the water molecule. This is however not the case with zinc metal. So the intermediary step provided by the sulfur pulling the oxygen atom out of the water molecule must, in some way, be necessary.

This is true of many metals but not all of them. Sodium metal for example, when placed in water, will be oxidized to NaOH (Na+[OH-]) without the presence of an electrolyte, with an accompanying release of hydrogen gas taking place. "For example, all the alkaline-earth metals have the ability to react with water to release hydrogen by the reaction;

M(s) + 2H2O-------> M++ and H2(g) and 2OH-" (1)(Section 19.1, Page 383)

 

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