Concentration cells and absurdities of modern science

Material revised and improved by Lipsa Dorelia.

Background and actual explanation

A concentration cell is a galvanic cell with half-cells of identical composition but differing concentrations.

The most common example is represented by two copper electrodes inserted in two solutions of copper sulfate at different concentrations.

The connection between half-cells is insured by a salt bridge or by an osmotic membrane and at copper electrodes a certain voltage is measured.

In case of osmotic membrane, there is the possibility for ions in solution to pass through, in order to equilibrate the concentrations, so after a certain interval of time both solutions arrive at same concentration and the difference of potential becomes null.

In case of a salt bridge, this furnishes ions to solutions and the same events happen. In concentrated half cell, copper is reduced and the remaining anion pass into bridge, and correspondently in half diluted cell, copper is oxidized and pass in solution and the bridge furnishes an anion.

Figure 1 Concentration Cell details

As shown in the fig. 1, the dilution of the cathode half-cell is achieved by reducing Cu2+ to Cu metal and plating that metal onto the Cu electrode. In the anode half-cell, the Cu anode is oxidized to Cu²⁺ and, thus, dissolved into the solution, making the anode cell more concentrated.

In order to calculate the cell potential for up specified concentrations (0.01M and 1M) at 25^oC, Nernst Equation is used:

E = E° - (0.059/n) log Q = +0.058 V

The concentration cells are an interesting case of cells, where no net chemical reaction occurs. The numbers of copper and sulphate ions in system does not change; it is the distribution of these species in the cell that provides the driving force.

From thermodynamic point of view, two solutions with different concentrations will diffuse one into another in order to arrive at a uniform concentration.

Putting copper electrodes into the solutions offers an alternative method of achieving the same end. One electrode can release a Cu2+ ion into the dilute solution; the electrons so produced then travel through the wires to the other electrode, and a Cu2+ ion is removed from the concentrated solution, reduced to a copper atom and plated out on the electrode. Both these events are simultaneously.

Why the actual explanation is erroneous

To date, the subject of concentration cell is reminded in the scientific texts without a detailed description.

The actual accepted explanation is not able to give a coherent explanation of ,,cathode and anode” phenomena. Why copper ions are reduced in case of concentrated solution and copper is oxidized in diluted solution?

Why there are no the opposite phenomena, more precisely Cu²⁺ from diluted solution is reduced and in case of concentrated solution Cu metal is oxidized?

Are directions of chemical reactions dictated by gradient concentration in the last time in chemistry?

Why for concentration cell oxidation at one electrode and a reduction at the other electrode are dictated by solution concentration?

From low level chemistry it is well known that an oxidation process takes place in presence of a reducing agent.

For example Cu metal can be oxidized by an oxidizing agent like oxygen or nitric acid. On the other hand a reduction of Cu²⁺ needs a reducing agent like hydrogen, etc.

Equations for the reactions are

2 Cu(s ) + O₂(g ) --> 2 CuO(s )

CuO(s ) + H₂(g ) --> Cu(s ) + H₂O(g )

By contrary to this well known comportment of copper, a simple diffusion of ions (without any chemical process) in case of a diffusion cell is able to produce an oxidation of Cu to one electrode and reduction of Cu²⁺ to other electrode.

A chemical reaction is a process whereby the chemical properties of a substance are changed by a rearrangement its atoms. The change produced by a chemical reaction is completely different from a purely physical change, which does not affect the fundamental properties of the substance itself.

Probably a new discipline of physical chemistry is appearing according to actual scientists, where chemical reactions are produced by simple diffusion of an inactive agent.

In case of a salt bridge between these half cells a more absurd situation is encountered. An inactive chemical compound in salt bridge is able to produce oxidation at one electrode and reduction to other electrode.

Here, the case of concentration cell with salt bridge is analyzed in detail.

In actual electrochemistry and physics it is admitted that salt bridge deliver the charged particle in order to have a neutrality of compartments in a cell.

In case of a concentration cell with a KCl salt bridge, it is figure out in fig. 2. the mechanism of this process.

Figure 2. Concentration cell with a KCl salt bridge

In the fig. 2 a limited number of KCl species are figured in order to have a simple and intuitive imagine of entire process.

According to actual interpretation, in concentrated CuSO₄ solution compartment (1M), copper deposits on the electrode.

Cu²⁺ + 2e Cu

The sulfate anion catch potassium from cell and K₂SO₄ is formed into solution.

SO₄^2- + 2K⁺ K₂SO₄

In diluted CuSO4 solution compartment (0,01M), copper is released from electrode.

Cu Cu²⁺ + 2e^-

The bridge delivers the chloride ion and CuCl₂ is formed in solution:

Cu²⁺ + 2Cl^- CuCl₂

According to this interpretation, in time the concentration of salt (KCl) in the bridge must decrease and in every compartment new compounds are formed as is figured out in fig. 3

Figure 3. Concentration cell with a KCl salt bridge - detail

As it can be seen, in order to have a functional concentration cell, in actual representation, a complex chain of reactions between Cu or CuSO₄ and KCl take place in half cells, but this reaction is impossible (endothermic) in principle.

What should be the difference of potential for an endothermic reaction?

According to actual interpretation, in one half cell (concentrated initial solution) after a certain time of cell working, potassium species and no chloride species must be detected.

In the other half cell, chloride species and no potassium must be detected.

It is very simple, with actual instrumentation, to be checked this conclusion of actual interpretation.

Let’s pass over the up presented objection, because other problems appear for actual interpretation.

The most important observation against actual interpretation must regard the impossibility of sulfate ion equilibration between compartments in case of a slat bridge.

It is accepted that an equilibration of activity (concentration) of sulfate ion represent ,,the engine” of a concentration cell. But the actual explanation does not give any possibility for a sulfate ion to pass from a compartment to another.

It can be imagined a process of sulfate pumping from a compartment to another according to following representation.

The starting point is the same cell with a limited quantity of salt in the bridge as in fig. 4.

Figure 4. Concentration cell with a KCl salt bridge- detail

The copper in concentrated compartment is deposited on electrode. The sulfate ion pushes the chloride ion from molecule to molecule until a chloride ion falls into the diluted CuSO₄ solution compartment. Here, a copper from electrode is released and CuCl₂ is formed.

After a time of working, the bridge is filled with K₂SO₄ as represented in fig. 5 and the sulfate is pumped from a compartment to another.

Does this mechanism can work in reality?

From a chemical point of view this mechanism is not feasible.

Even it is supposed that sulfate ion has enough energy to remove a chloride ion from a KCl molecule, it will be the first case in chemistry when sulfate pushes an entire chain of chloride atoms.

In case of high quantity of KCl in the bridge, the cell is working for long time with a circuit of chloride from bridge to solution. In this case it is possible to have after a week the cell working, without any transfer of sulfate from a compartment to another. Does the intention of sulfate to move from a compartment to another really produce a countable energy?

In case of low quantity of KCl alt in the bridge, after a certain time interval, the bridge should contain K₂SO₄ and no chloride.

Figure 5. Concentration cell with a KCl salt bridge -detail

Concentration Cell – Experimental Part

Experiment 1

A copper concentration cell is device that plates out copper from a concentrated Cu²⁺(aq) solution and dissolves copper from an electrode into a relatively dilute solution of copper ion. The purpose of this experiment is to measure the mass transfer in case of a concentration cell.

Materials

ü 1 M copper sulfate solution,

ü 0.01 M copper sulfate solution

ü two PVC containers ( as in next figure - section), with copper electrodes, passing taps and the possibility of connection through a membrane filter or a thicker layer.

In the experiment the containers was 2 liter volume and the filter membrane used was a nylon 0,22 micrometer pore size as in fig. 6. One container was filled with 1 M solution and the other with a 0.01M of CuSO4. Other barrier layer are foreseen to be used in the future (maybe someone suggests one as described in actual scientific texts able to permit a unidirectional circulation only for sulfate species !!!). The taps of system were opened and the potential was measured in order to check the actual prediction (Nernst equation); further the system was leaved for a week in order to equilibrate the concentrations in both compartments.

After this time both electrodes were weighed with a precision balance and used again for another week, filling the containers with fresh solutions. At the end of second week again both electrodes are again weighed and the procedure of forming a new concentration cell was repeated. The procedure was repeated for the third week again.

During these 3 weeks of concentrations equilibration there was no transfer of copper electrodes mass from one to another. On the electrode immersed in concentrated CuSO₄ solution, in time the blue CuSO₄ adhere to the electrode and the mass of this electrode is increased a little bit (about 1 g). If the electrode is washed with clear water the final mass of electrode is the same like initial one.

The increase of electrode mass found in the diluted solution is about 0.3 g. Again after simple washing the mass of this electrode remain invariant. The increase of mass electrodes is caused by CuSO₄ which adhere to the metal proportional with its concentration in solution and not due to a chemical process.

Figure 6. Concentration cell details

In the experiment were measured both electrodes in order to avoid some experimental errors. A mass transport of Cu should give us a correlation between an eventually decrease of an electrode mass and increase of another one.

What are the prediction of actual physical chemistry for up presented experiment and how up presented results fit to actual theory?

It is known from thermodynamic that the process

solute (concentrated) → solute (dilute) is accompanied by a fall in free energy.

A concentration cell use this energy and some chemical processes take place according to:

cathode: Cu²⁺(1 M) + 2e– → Cu(s)

anode: Cu(s) → Cu²⁺(.01 M) + 2e–

The net reaction regard a transfer of copper mass from anode (diluted container) to cathode (concentrated container).

What is the theoretical mass of transported copper?

If anode generates a Cu cation, another sulfate anion must pass the membrane in order to have the neutrality of solution and simultaneously a copper atom is deposited on the cathode. The final concentration of both solutions is 0.5M (the initial concentration of anode compartment is neglected).

In order to have a 0.5 M sulfate concentration in anode compartment, the quantity of copper released in 2 liters of solution must be equal with 1 mole, which means 63,5 g.

On the other hand the mass of cathode should appear increased with 63,5 g.

As initial mass of electrodes was about 25 g, normally after one week one electrode must be dissolved into solution and the mass of second electrode should be increased with correspondent quantity.

As the solution was changed once at week for three times, theoretically, the total quantity of copper displaced must be about 190 g. Even in a case of a fluffy deposited copper, it is impossible to not observe this quantity at solution change.

The experimental reality is completely different. Except the CuSO₄ material which adheres to one electrode, no increased mass of copper is observed for one electrode and no decrease mass for the other electrode.

The up presented calculus was made in considering the concentration cell working with a yield of 100 %. Of course there is a Fick diffusion between compartments which diminishes the copper transfer. In case of a concentration cell working with a yield of 0.1 % (which is absurd because the cells are recognized to have a better yield then another thermodynamic processes), at a single concentration equilibration should appear a mass transfer of 0.1*63.5 =6,35 g.

Experiment 2

Material:

ü 1 M copper sulfate solution,

ü 0.01 M copper sulfate solution

ü two small containers with copper electrodes,

ü salt bridge made by saturating a piece of cotton in a CaCl₂ solution and another one made with agar agar and CaCl₂. The use of CaCl₂ in the salt bridge instead of KCl is motivated only by the simpler analytical possibilities of Ca determination.

Two concentrations cells are made using the up presented salt bridges and their potentials are measured. After this operation, the Cu electrodes are put in short circuit with a simple metallic conductor in order to hurry up the electrochemical processes as in fig. 7.

Figure 7. CaCl₂ bridge concentration cell

After 3 days of working, the metallic conductor is removed and the difference of potential between electrodes is measured. Contrary to expectations the cells present the same difference of potential like initial.

After measurement again the electrodes are short circuit and leaved for another 5 days. After this period again the potential difference is measuread.

In case of agar agar CaCl₂ bridge, the difference of potential remains constant after 8 days of cell working. The measurements of Cl and Ca concentrations in both compartments, in case of agar agar bridge, indicate a level of their concentration under the chemical limit of detection (precipitation with AgNO₃ and H₂SO₄ after both time intervals).

In the cotton piece bridge at the same time interval there is a diminishing of cell potential with approx. 0.02 V after 3 days and with approx. 0.07 V after another 5 days. Both Ca and Cl species were detected in the anode and cathode compartments. A normal analytical procedure was followed for measuring their concentrations and it was found that:

Anode [Ca]=[Cl]/2

Cathode [Ca] =[Cl]/2

The concentrations of Ca and Cl species at anode are a little bit higher then at cathode region (probably due to the small different level of immersion of bridge into solution).

Let’s analyze the theoretical consideration in the frame of actual physics for the up presented experiment.

Consider the following concentration cell:

Cu(s) | CuSO4(1 M) || CuSO4(.01 M) | Cu(s)

with a salt bridge made by KCl.

This kind of cell, provide experimentally a potential difference closed by the Nernst equation prediction.

Nernst equation:

E = E° - (0.059/n) log Q = 0 - 0.029 log 0.01 = +0.058 V

The experimental measured difference was E’ = 0.048 in case of agar CaCl₂ bridge and 0.046 in case of cotton CaCl₂, a little bit smaller then predicted. It seems that theory fit very well with practice. The difference can be very easy attributed to the contact potential existent in the circuit.

But is this appearance based on a consistent reality? How is this kind of cell working properly?

It is known and accepted by actual physics, the fact that an electrochemical cell needs an exothermic chemical reaction or a transport of an ionic species from a high concentration domain to a low concentration one as fundament of a potential difference generation.

Anyone can test the impossibility of a free (spontaneous) reaction between CuSO₄, Cu and KCl , in a bi or tri component system reaction.

ISOTOPE CHANGE AND ELECTROCHEMISTRY

The mass transfer and what’s happened at interface metal solution can be very easy verified with isotopes techniques.

Having a metal electrode in a solution that contains ions of that metal, a potential difference between the metal and the solution appears according to actual interpretation due to the following equilibrium:

M Mⁿ⁺ + ne^-

When the metal strip contain only one isotope and the solution of its salt contain another isotope after a period of time there will be a process of isotopic change between metal and solution. Consequently is very easy for example to use a radioisotope in the metal strip and non radioisotope for salt solution and after a time to measure the isotopic exchange (the lows of isotopic exchange are well known).

Of course the experiment can be performed with non radioactive isotopes, but in this case a isotope ratio mass spectrometer is necessary in order to observe the isotopic pattern change in solution and in electrode.

In proposed theory, the experiment will fail to give a expected result – more precisely no isotope exchange is observed between metal and its salt solution.