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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Standard half-cell potentials

Problem Example 1
Find the standard potential of the cell

Cu(s) | Cu²⁺ || Cl^– | AgCl(s) | Ag(s)

and predict the direction of electron flow when the two electrodes are connected.
Solution: The net reaction corresponding to this cell will be

2 Ag(s) + 2 Cl^–(aq) + Cu²⁺(aq) → AgCl(s) + Cu(s)

Since this involves the reverse of the AgCl reduction, we must reverse the corresponding half-cell potential:

E_cell = (.337 – .222) v = .115 v

Since this potential is positive, tthe reaction will proceed to the right; electrons will be withdrawn from the copper electrode and flow through the external circuit into the silver electrode. Note carefully that in combining these half-cell potentials, we did not multiply E° the for the Cu²⁺/Cu couple by two. The reason for this will be explained later.

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What is an activity series, and how is it used?

An activity series is a list of substances ranked in order of relative reactivity. For example, magnesium metal can knock hydrogen ions out of solution, so it is considered more reactive than elemental hydrogen:

Mg(s) + 2 H⁺(aq) H₂(g) + Mg²⁺(aq)

Zinc can also displace hydrogen ions from solution:

Zn(s) + 2 H⁺(aq) H₂(g) + Zn²⁺(aq)

so zinc is also more active than hydrogen. But magnesium metal can remove zinc ions from solution:

Mg(s) + Zn²⁺(aq) Zn(s) + Mg²⁺(aq)

The reaction goes nearly to completion. Magnesium is more active than zinc, and the activity series including these elements would be Mg > Zn > H. The following activity series built up in a similar way. The most active metals are at the top of the table; the least active are at the bottom. Each metal is able to displace the elements below it from solution (or, using the language of electrochemistry, each metal can reduce the cations
of metals below it to their elemental forms).

The metal activity series. Most active (most strongly reducing) metals appear on top, and least active metals appear on the bottom.

displace H2 from water, steam, or acids	Li	2 Li(s) + 2 H2O() 2 LiOH(aq) + H2(g)
	K	2 K(s) + 2 H2O() 2 KOH(aq) + H2(g)
	Ca	Ca(s) + 2 H2O() Ca(OH)2(s) + H2(g)
	Na	2 Na(s) + 2 H2O() 2 NaOH(aq) + H2(g)
displace H2 from steam or acids	Mg	Mg(s) + 2 H2O(g) Mg(OH)2(s) + H2(g)
	Al	2 Al(s) + 6 H2O(g) 2 Al(OH)3(s) + 3 H2(g)
	Mn	Mn(s) + 2 H2O(g) Mn(OH)2(s) + H2(g)
	Zn	Zn(s) + 2 H2O(g) Zn(OH)2(s) + H2(g)
	Fe	Fe(s) + 2 H2O(g) Fe(OH)2(s) + H2(g)
displace H2 from acids only	Ni	Ni(s) + 2 H+(aq) Ni2+(aq) + H2(g)
	Sn	Sn(s) + 2 H+(aq) Sn2+(aq) + H2(g)
	Pb	Pb(s) + 2 H+(aq) Pb2+(aq) + H2(g)
	H2
can't displace H2	Cu
	Ag
	Pt
	Au

 
The activity series is a useful guide for predicting the products of metal displacement reactions. For example, placing a strip of zinc metal in a copper(II) sulfate solution will produce metallic copper and zinc sulfate, since zinc is above copper on the series. A strip of copper placed into a zinc sulfate solution will not produce an appreciable reaction, because copper is below zinc on the series and can't displace zinc ions from solution.
The series works well as long as the reactions being predicted occur at room temperature and in aqueous solution. It isn't difficult to find reactions that are at odds with the metal and nonmetal activity series under other conditions. There are other complications too. For example, aluminum would be expected to displace hydrogen from steam, but in fact it won't unless the aluminum oxide film on its surface is scrubbed off. Copper can't displace hydrogen from acids, but it does react with acids like nitric and sulfuric because they can act as oxidizing agents.
It might be expected that metals with lower ionization energies and lower electronegativities would be more active, since they would be expected to more easily lose electrons in a displacement reaction. But while ionization energy and electronegativity do affect a metal's ranking in the series, other factors have a strong and complex influence on relative activity. obscuring the relationship.
Activity series can be devised for nonmetals as well. Since nonmetallic elements tend to accept electrons in redox reactions, the nonmetal activity series is arranged so that the most powerful oxidizing agents are considered most active (whereas in the metal series, the most powerful reducing agents are the most active):
 

The nonmetal activity series. Most active (most strongly oxidizing) nonmetals appear on top, and least active nonmetals appear on the bottom.

F2	strongest oxidizing agent
Cl2
O2
Br2
I2
S
red P	weakest oxidizing agent

For example, the series predicts that Cl₂ will displace Br^- and I^- from solution, because Cl₂ appears above Br₂ and I₂:

Cl₂(g) + 2 Br^-(aq) 2 Cl^-(aq) + Br₂()
Cl₂(g) + 2 I^-(aq) 2 Cl^-(aq) + I₂(s)
Br₂() + 2 Cl^-(aq) no reaction
I₂(s) + 2 Cl^-(aq) no reaction

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Reference Electrodes

     In most electrochemical experiments our interest is concentrated on only one of the electrode reactions. Since all measurements must be on a complete cell involving two electrode systems, it is common practice to employ a reference electrode as the other half of the cell. The major requirements of a reference electrode are that it be easy to prepare and maintain, and that its potential be stable. The last requirement essentially means that the concentration of any ionic species involved in the electrode reaction must be held at a fixed value. The most common way of accomplishing this is to use an electrode reaction involving a saturated solution of an insoluble salt of the ion. One such system, the silver-silver chloride electrode has already been mentioned:

Ag | AgCl(s) | Cl^–(aq) || ...

Ag(s) + Cl^–(aq) →AgCl(s) + e^–

This electrode usually takes the form of a piece of silver wire coated with AgCl. The coating is done by making the silver the anode in an electrolytic cell containing HCl; the Ag⁺ ions combine with Cl^– ions as fast as they are formed at the silver surface.

The other common reference electrode is the calomel electrode; calomel is the common name for

 mercury(I) chloride.

Hg | Hg²⁺(aq) | KCl || ... Hg(l) + Cl^– → ½ HgCl₂(s) + e^–

The potentials of both of these electrodes have been very accurately determined against the hydrogen electrode. The latter is seldom used in routine electrochemical measurements because it is more difficult to prepare; the platinum surface has to be specially treated by preliminary electrolysis. Also, there is need for a supply of hydrogen gas which makes it somewhat cumbersome and hazardous.

Some notes :

Make sure you thoroughly understand the following essential ideas which have been presented above. It is especially imortant that you know the precise meanings of all the highlighted terms in the context of this topic.

• A galvanic cell (sometimes more appropriately called a voltaic cell) consists of two half-cells joined by a salt bridge or some other path that allows ions to pass between the two sides in order to maintain electroneutrality.
• The conventional way of representing an electrochemical cell of any kind is to write the oxidation half reaction on the left and the reduction on the right. Thus for the reaction

Zn(s) + Cu²⁺ → Zn²⁺ + Cu(s)

we write

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

in which the single vertical bars represent phase boundaries. The double bar denotes a liquid-liquid boundary which in laboratory cells consists of a salt bridge or in ion-permeable barrier. If the net cell reaction were written in reverse, the cell notation would become

Cu(s) | Cu²⁺(aq) || Zn ²⁺(aq) | Zn (s)

Remember: the Reduction process is always shown on the Right.

• The transfer of electrons between an electrode and the solution takes place by quantum-mechanical tunneling at the electrode surface. The energy required to displace water molecules from the hydration shell of an ion as it approaches the electrode surface constitutes an activation energy which can slow down the process. Even larger activation energies (and slower reactions) occur when a molecule such as O₂ is formed or consumed.

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Electrodes and electrode reactions.

An electrode reaction refers to the net oxidation or reduction process that takes place at an electrode. This reaction may take place in a single electron-transfer step, or as a succession of two or more steps. The substances that receive and lose electrons are called the electroactive species.

Fig. 4: Electron transfer at an anode

This process takes place within the very thin interfacial region at the electrode surface, and involves quantum-mechanical tunneling of electrons between the electrode and the electroactive species. The work required to displace the H₂O molecules in the hydration spheres of the ions constitutes part of the activation energy of the process.

In the example of the Zn/Cu cell we have been using, the electrode reaction involves a metal and its hydrated cation; we call such electrodes metal-metal ion electrodes. There are a number of other kinds of electrodes which are widely encountered in electrochemistry and analytical chemistry.

Ion-ion electrodes

Many electrode reactions involve only ionic species, such as Fe²⁺ and Fe³⁺. If neither of the electroactive species is a metal, some other metal must serve as a conduit for the supply or removal of electrons from the system. In order to avoid complications that would arise from electrode reactions involving this metal, a relatively inert substance such as platinum is commonly used. Such a half cell would be represented as

Pt(s) | Fe³⁺(aq), Fe²⁺(aq) || ...

and the half-cell reaction would be

Fe²⁺(aq) → Fe³⁺ (aq) + e^–

The reaction occurs at the surface of the electrode (Fig 4 above). The electroactive ion diffuses to the electrode surface and adsorbs (attaches) to it by van der Waals and coulombic forces. In doing so, the waters of hydration that are normally attached to any ionic species must be displaced. This process is always endothermic, sometimes to such an extent that only a small fraction of the ions be able to contact the surface closely enough to undergo electron transfer, and the reaction will be slow. The actual electron-transfer occurs by quantum-mechanical tunnelling.

Gas electrodes

Some electrode reactions involve a gaseous species such as H₂, O₂, or Cl₂. Such reactions must also be carried out on the surface of an electrochemically inert conductor such as platinum. A typical reaction of considerable commercial importance is

Cl^–(aq) → ½ Cl₂(g) + e^–

Similar reactions involving the oxidation of Br₂ or I₂ also take place at platinum surfaces.

Insoluble–salt electrodes

A typical electrode of this kind consists of a silver wire covered with a thin coating of silver chloride, which is insoluble in water. The electrode reaction consists in the oxidation and reduction of the silver:

AgCl(s) + e^– → Ag(s) + Cl^–(aq)

The half cell would be represented as

... || Cl^– (aq) | AgCl (s) | Ag (s)

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Cell description conventions

In order to make it easier to describe a given electrochemical cell, a special symbolic notation has been adopted. In this notation the cell we described above would be

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

There are several other conventions relating to cell notation and nomenclature that you are expected to know:

• The anode is where oxidation occurs, and the cathode is the site of reduction. In an actual cell, the identity of the electrodes depends on the direction in which the net cell reaction is occurring.
• If electrons flow from the left electrode to the right electrode (as depicted in the above cell notation) when the cell operates in its spontaneous direction, the potential of the right electrode will be higher than that of the left, and the cell potential will be positive.
• "Conventional current flow" is from positive to negative, which is opposite to the direction of the electron flow. This means that if the electrons are flowing from the left electrode to the right, a galvanometer placed in the external circuit would indicate a current flow from right to left.

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Transport of charge within the cell

     For the cell to operate, not only must there be an external electrical circuit between the two electrodes, but the two electrolytes (the solutions) must be in contact. The need for this can be understood by considering what would happen if the two solutions were physically separated. Positive charge (in the form of Zn²⁺) is added to the electrolyte in the left compartment, and removed (as Cu²⁺) from the right side, causing the solution in contact with the zinc to acquire a net positive charge, while a net negative charge would build up in the solution on the copper side of the cell. These violations of electroneutrality would make it more difficult (require more work) to introduce additional Zn²⁺ ions into the positively-charged electrolyte or for electrons to flow into right compartment where they are needed to reduce the Cu²⁺ ions, thus effectively stopping the reaction after only a chemically insignificant amount has taken place.
     In order to sustain the cell reaction, the charge carried by the electrons through the external circuit must be accompanied by a compensating transport of ions between the two cells. This means that we must provide a path for ions to move directly from one cell to the other. This ionic transport involves not only the electroactive species Cu²⁺ and Zn²⁺, but also the counterions, which in this example are nitrate, NO₃^-.
     Thus an excess of Cu²⁺ in the left compartment could be alleviated by the drift of these ions into the right side, or equally well by diffusion of nitrate ions to the left. More detailed studies reveal that both processes occur, and that the relative amounts of charge carried through the solution by positive and negative ions depends on their relative mobilities, which express the velocity with which the ions are able to make their way through the solution. Since negative ions tend to be larger than positive ions, the latter tend to have higher mobilities and carry the larger fraction of charge.

     In the simplest cells, the barrier between the two solutions can be a porous membrane, but for precise measurements, a more complicated arrangement, known as a salt bridge, is used. The salt bridge consists of an intermediate compartment filled with a concentrated solution of KCl and fitted with porous barriers at each end. The purpose of the salt bridge is to minimize the natural potential difference, known as the junction potential, that develops (as mentioned in the previous section) when any two phases (such as the two solutions) are in contact. This potential difference would combine with the two half-cell potentials so as introduce a degree of uncertainty into any measurement of the cell potential. With the salt bridge, we have two liquid junction potentials instead of one, but they tend to cancel each other out.

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Galvanic Cells

     This arrangement is called a Galvanic cell. A typical cell might consist of two pieces of metal, one zinc and the other copper, each immersed each in a solution containing a dissolved salt of the corresponding metal. The two solutions are separated by a porous barrier that prevents them from rapidly mixing but allows ions to diffuse through.

     If we connect the zinc and copper by means of a metallic conductor, the excess electrons that remain when Zn2+ ions emerge from the zinc in the left cell would be able to flow through the external circuit and into the right electrode, where they could be delivered to the Cu2+ ions which become "discharged", that is, converted into Cu atoms at the surface of the copper electrode. The net reaction is the oxidation of zinc by copper(II) ions:

Zn(s) + Cu2+ → Zn2+ + Cu(s)

but this time, the oxidation and reduction steps (half reactions) take place in separate locations:

left electrode: Zn(s) → Zn2+ + 2e–       oxidation

right electrode: Cu2+ + 2e–→ Cu(s)     reduction

 

Electrochemical cells allow measurement and control of a redox reaction.

     The reaction can be started and stopped by connecting or disconnecting the two electrodes. If we place a variable resistance in the circuit, we can even control the rate of the net cell reaction by simply turning a knob. By connecting a battery or other source of current to the two electrodes, we can force the reaction to proceed in its non-spontaneous, or reverse direction.

     By placing an ammeter in the external circuit, we can measure the amount of electric charge that passes through the electrodes, and thus the number of moles of reactants that get transformed into products in the cell reaction.

     Electric charge q is measured in coulombs. The amount of charge carried by one mole of electrons is known as the faraday, which we denote by F. Careful experiments have determined that 1 F = 96467 c. For most purposes, you can simply use 96,500 coulombs as the value of the faraday.

When we measure electric current, we are measuring the rate at which electric charge is transported through the circuit. A current of one ampere corresponds to the flow of one coulomb per second.

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