Electrolytic refining of aluminum

     

     Aluminum is present in most rocks and is the most abundant metallic element in the earth's crust (eight percent by weight.) However, its isolation is very difficult and expensive to accomplish by purely chemical means, as evidenced by the high E° (–1.66 v) of the Al³⁺/Al couple. For the same reason, aluminum cannot be isolated by electrolysis of aqueous solutions of its compounds, since the water would be electrolyzed preferentially. And if you have ever tried to melt a rock, you will appreciate the difficulty of electrolyzing a molten aluminum ore! Aluminum was in fact considered an exotic and costly metal until 1886, when Charles Hall (U.S.A) and Paul Hérault (France) independently developed a practical electrolytic reduction process.

The Hall-Hérault process takes advantage of the principle that the melting point of a substance is reduced by admixture with another substance with which it forms a homogeneous phase. Instead of using the pure alumina ore Al₂O₃ which melts at 2050°C, it is mixed with cryolite, which is a natural mixture of NaF and AlF₃, thus reducing the temperature required to a more manageable 1000°C. The anodes of the cell are made of carbon (actually a mixture of pitch and coal), and this plays a direct role in the process; the carbon gets oxidized (by the oxide ions left over from the reduction of Al³⁺ to CO, and the free energy of this reaction helps drive the aluminum reduction, lowering the voltage that must be applied and thus reducing the power consumption. This is important, because aluminum refining is the largest consumer of industrial electricity, accounting for about 5% of all electricity generated in North America. Since aluminum cells commonly operated at about 100,000 amperes, even a slight reduction in voltage can result in a large saving of power.
The net reaction is

2 Al₂O₃ + 3 C → 4 Al + 3 CO₂

However, large quantities of CO and of HF (from the cryolite), and hydrocarbons (from the electrodes) are formed in various side reactions, and these can be serious sources of environmental pollution.

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Mind Map.

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Difference between Electrochemical Cell and Electrolytic Cells.

Electrochemical Cell	Electrolytic Cell
It converts chemical energy into electrical energy.	It converts electrical energy into chemical energy.
It is based upon the redox reactions which are spontaneous.	The redox reactions are non-spontaneous and take place only when energy is supplied.
The chemical changes occurring in the two beakers are different.	Only one chemical compound undergoes decomposition.
Anode (-ve) - Oxidation takes place.	Anode (+ve) - Oxidation takes place.
Cathode (+ve) - Reduction takes place.	Cathode (-ve) - Reduction takes place.

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Industrial electrolytic processes.

The chloralkali industry

The electrolysis of brine is carried out on a huge scale for the industrial production of chlorine and caustic soda (sodium hydroxide). Because the reduction potential of Na⁺ is much higher than that of water, the latter substance undergoes decomposition at the cathode, yielding hydrogen gas and OH^–.

anode reactions	2 Cl– → Cl2(g) + 2 e–	-1.36 v	i
	4 OH– → O2(g) + 2 H2O + 4 e–	-0.40 v	ii
cathode reactions	Na+ + e– → Na(s)	-2.7 v	iii
	H2O + 2 e– → H2(g) + 2 OH–	+.41 v	iv

A comparison of the E°s would lead us to predict that the reduction (ii) would be favored over that of (i). This is certainly the case from a purely energetic standpoint, but as was mentioned in the section on fuel cells, electrode reactions involving O₂ are notoriously slow (that is, they are kinetically hindered), so the anodic process here is under kinetic rather than thermodynamic control. The reduction of water (iv) is energetically favored over that of Na⁺ (iii), so the net result of the electrolysis of brine is the production of Cl₂ and NaOH ("caustic"), both of which are of immense industrial importance:

2 NaCl + 2 H₂O → 2 NaOH + Cl₂(g) + H₂(g)

     Since chlorine reacts with both OH^– and H₂, it is necessary to physically separate the anode and cathode compartments. In modern plants this is accomplished by means of an ion-selective polymer membrane, but prior to 1970 a more complicated cell was used that employed a pool of mercury as the cathode. A small amount of this mercury would normally find its way into the plant's waste stream, and this has resulted in serious pollution of many major river systems and estuaries and devastation of their fisheries.

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Faraday's laws of electrolysis.

One mole of electric charge (96,500 coulombs), when passed through a cell, will discharge half a mole of a divalent metal ion such as Cu²⁺. This relation was first formulated by Faraday in 1832 in the form of two laws of electrolysis:

1. The weights of substances formed at an electrode during electrolysis are directly proportional to the quantity of electricity that passes through the electrolyte.
2. The weights of different substances formed by the passage of the same quantity of electricity are proportional to the equivalent weight of each substance.

The equivalent weight of a substance is defined as the molar mass, divided by the number of electrons required to oxidize or reduce each unit of the substance. Thus one mole of V³⁺ corresponds to three equivalents of this species, and will require three faradays of charge to deposit it as metallic vanadium.

Most stoichiometric problems involving electrolysis can be solved without explicit use of Faraday's laws. The "chemistry" in these problems is usually very elementary; the major difficulties usually stem from unfamiliarity with the basic electrical units:

• current (amperes) is the rate of charge transport; 1 amp = 1 c/sec.
• power (watts) is the rate of energy production or consumption;
  1 w = 1 J/sec = 1 volt-amp; 1 watt-sec = 1 J, 1 kw-h = 3600 J.

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Electrolysis.

     Electrolysis refers to the decomposition of a substance by an electric current. The electrolysis of sodium and potassium hydroxides, first carried out in 1808 by Sir Humphrey Davey, led to the discovery of these two metallic elements and showed that these two hydroxides which had previously been considered un-decomposable and thus elements.

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Significance of the Nernst Equation.

 

The Nernst equation tells us that a half-cell potential will change by 59 millivolts per 10-fold change in the concentration of a substance involved in a one-electron oxidation or reduction; for two-electron processes, the variation will be 28 millivolts per decade concentration change. Thus for the dissolution of metallic copper

Cu(s) → Cu²⁺ + 2e^–

the potential

E = (– 0.337) – .0295 log [Cu²⁺]

becomes more positive (the reaction has a greater tendency to take place) as the cupric ion concentration decreases. This, of course, is exactly what the Le Châtelier Principle predicts; the more dilute the product, the greater the extent of the reaction.

 

Electrodes with poise

The equation just above for the Cu/Cu²⁺ half-cell raises an interesting question: suppose you immerse a piece of copper in a solution of pure water. With Q = [Cu²⁺] = 0, the potential difference between the electrode and the solution should be infinite! Are you in danger of being electrocuted? You need not worry; without any electron transfer, there is no charge to zap you with. Of course it won't be very long before some Cu²⁺ ions appear in the solution, and if there are only a few such ions per liter, the potential reduces to only about 20 volts. More to the point, however, the system is so far from equilibrium (for example, there are not enough ions to populate the electric double layer) that the Nernst equation doesn't really give meaningful results. Such an electrode is said to be unpoised. What ionic concentration is needed to poise an electrode? I don't really know, but I would be suspicious of anything much below 10^–6 M.

The Nernst equation works only in dilute ionic solutions

Ions of opposite charge tend to associate into loosely-bound ion pairs in more concentrated solutions, thus reducing the number of ions that are free to donate or accept electrons at an electrode. For this reason, the Nernst equation cannot accurately predict half-cell potentials for solutions in which the total ionic concentration exceeds about 10^–3 M.

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