Химия и фармакология

Matters are different in compound substances. The chemical bonds between the atoms of different elements are not symmetrical; polar Bonds are generally the rule in molecules of compounds. This nonuniformity in the distribution of the electrons is the greatest in ionic compounds...



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Oxidation of Elements

When an element is in the free state—forms an elementary substance—the motion of the electrons about all the atoms of this substance occurs identically. This holds for all elementary substances regardless of their structure. For example, in a hydrogen molecule, the electrons travel about both atoms to an equal extent—the molecule Hg is non-polar. For crystals with a covalent bond, the chemical bonds between the atoms are also symmetrical relative to the joined atoms. For metals, the distribution of both the bound and the free electrons is also uniform on an average.

Matters are different in compound substances. The chemical bonds between the atoms of different elements are not symmetrical; polar Bonds are generally the rule in molecules of compounds. This nonuniformity in the distribution of the electrons is the greatest in ionic compounds—in the formation of substances with an ionic bond, the valence electrons pass virtually completely from an atom of one element to an atom of another one.

An element whose electrons pass to atoms of another element (completely with an ionic bond or partly with a polar one) is said to bepositively oxidised. An element to whose atoms electrons from atoms of another element pass is negatively oxidised(or reduced).

The number of electrons that have passed from one atom of a given positively oxidised element or to one atom of a given negatively oxidised or reduced element is called the oxidation number (oroxidation state) of the element.

In elementary substances, the oxidation number of an element is always zero. In compounds, some elements always display the same oxidation number, but for most elements it differs in different compounds.

The elements having a constant oxidation number are the alkali metals (+1), the alkaline-earth metals (-1-2), and fluorine (—1). hydrogen in most compounds is characterised by an oxidation number of +1, while in metal hydrides (Vol. 2, Sec. 2) and in some other compounds it is —1. The oxidation number of oxygen, as a rule, is —2. The most important exceptions here are the peroxide compounds, where it is —1, and oxygen difluoride OF2, in which the oxidation number of oxygen is +2. For elements with a changing oxidation number, its value is always simple to find knowing the formula of a given compound and taking into consideration that the sum of the oxidation numbers of all the atoms in a molecule is zero.

Let us determine as an example the oxidation number of carbon in CO, CO2, CH4, C2H6, and C2H5OH. We shall denote it byx.Hence, remembering that the oxidation number of hydrogen is +1 and that of oxygen —2, we get

To find the oxidation number of elements in compounds, the table of electronegativities of elements (Table 6) can be used. It must be borne in mind here that when a chemical bond forms, the electrons are displaced to the atom of the element with the higher electronegativity. For instance, the relative electronegativity of phosphorus is 2.2, and that of iodine 2.6. Consequently, in the compound PI 3, the shared electrons are displaced to the iodine atoms, and the oxidation numbers of phosphorus and iodine are +3 and —1, respectively. In Fig, however, the oxidation numbers of nitrogen and iodine are —3 and +1 because the electronegativity of nitrogen (3.07) is higher than that of iodine.

Oxidation-Reduction Reactions

All chemical reactions can be divided into two groups. In those of the first group, the oxidation state of all the elements in the reactants remains constant, while in those of the second group, the oxidation state of one or several elements changes.

We can exemplify the reactions of the first group by the neutralisation reaction

HC1 + NaOH = NaCI + H2O

An example of a reaction of the second group is the reaction of a metal with an acid:

Zn + 2HC1 = ZnCL2 + H2

If in the neutralisation reaction no element changes its oxidation number, in the second example the oxidation number of the zinc changes from 0 to +2, and of the hydrogen from +1 to 0.

Reactions as a result of which the oxidation numbers of elements change are called oxidation-reduction (redox) reactions.

Oxidation-reduction reactions are of very great importance in biological systems. Photosynthesis, breathing, digestion—all these are chains of oxidation-reduction reactions. In engineering, the significance of redox reactions is also very great. For example, the entire metallurgical industry is based on oxidation-reduction processes during which metals are recovered from natural compounds.

A simple example of an oxidation-reduction reaction is the reaction of formation of an ionic compound from elementary substances, for example, the reaction of sodium with chlorine:

2Na + Cl2 = 2NaCl

This reaction, like any heterogeneous one, proceeds in several steps. In one of them, the sodium atoms transform into positively charged ions; the oxidation number of sodium changes from 0 to 4-1:

Na = Na+ +e-

Such a process—the losing of electrons attended by an increase in the oxidation number of an element—is called oxidation.

The electrons lost by the sodium are gained by the chlorine atoms, which transform into negatively charged ions; the oxidation number of the chlorine changes from 0 to —1:

Cl2 + 2e- == 2C1-

The gaining of electrons attended by a decrease in the oxidation number of an element is called redaction.

Hence, in the reaction considered above, the sodium is oxidised, and the chlorine is reduced.

A substance containing an element that is oxidised is called a reducing agent, and a substance containing an element that is reduced is called an oxidising agent.

Consequently, in the above example, sodium is a reducing agent, and chlorine an oxidising one.

Inspection of the oxidation-reduction equations shows that one molecule of chlorine when reduced gains two electrons, while the oxidation of one sodium atom is attended by the loss of one electron. The total number of electrons in a system in chemical reactions does not change:the number of electrons lost by the molecules {atoms, ions) of the reducing agent equals the number of electrons gained by the molecules (atoms, ions) of the oxidising agent. Therefore, one molecule of chlorine can oxidise two sodium atoms.

Compiling Equations of Oxidation-Reduction Reactions

In Sec. 94, we considered a very simple example of a redox reaction — the formation of a compound from two elementary substances. Usually, the equations of redox reactions are more complicated and it is often a quite difficult task to balance them. We shall consider a few examples.

Example 1. The reaction between hydrogen iodide and concentrated sulphuric acid. This reaction proceeds as follows:

If we calculate the oxidation number of each element in the reactants and in the products, we shall see that it changes both in the iodine and in the sulphur. In the iodine of HI it is —1, and in free iodine it is 0. The oxidation number of sulphur, on the other hand, changes from +6 (in H2S04) to —2 (in H2S). Hence, the oxidation number of the iodine increases, and of the sulphur decreases. Consequently, the iodine is oxidised, and the sulphur is reduced.

The equation of the iodine oxidation has a simple form:

2I- == l2 + 2e- (oxidation)

The equation of sulphur reduction is more complicated because both the reactant (H2SO4 or S042-) and the product (H2S) contain other elements in addition to sulphur. In compiling this equation, we shall proceed from the fact that the reaction takes place in an. acid aqueous solution, and the ion S042- transforms into the molecule H2S:

S042--® H2S

The four oxygen atoms liberated in this process should combine with hydrogen into four molecules of water. Eight hydrogen ions are needed for this. In addition, two hydrogen ions are needed to form a molecule of H2S. Hence, ten hydrogen ions should react with the SO42- ion:

SO42- + 1OH+®H2S + 4H2O

The total charge on the ions in the left-hand side of this formula is eight elementary positive charges, while the right-hand side contains only uncharged particles. Since the total charge does not change in the course of the process, eight electrons also participate in the reduction process:

SO42- + 10H+ + 8e-® H2S + 4H2O (reduction)

In the above example, the ratio of the number of electrons participating in the reduction process to that liberated in oxidation is 4:1. To obtain the overall equation of the reaction, we have to take this ratio into account when summating the equations of the reduction and oxidation processes, i.e. multiply the reduction equation by four. It is customary practice to indicate the required multipliers to the right of a vertical line when writing the equations:

The reaction equation obtained can also be written in the molecular form:

H2SO4+8HI =4I2+H2S + H2O

Example 2. Reaction of aluminium with potassium nitrate in a basic solution. The skeleton equation is:

KN03 + Al + KOH + H2O® NH3 + KAlO2

Here the oxidation number changes in the nitrogen and aluminium. Metallic aluminium (its oxidation number is zero) transforms into the ion A10, in which the oxidation number of the aluminium is +3. To compile the oxidation equation, we shall proceed from the scheme:

Al® AlO2-

In a basic solution, the OH" ion is the source of the oxygen needed for this process to occur. Four hydroxide ions are required to combine one aluminium atom into the AlO2- ion:

Al + 40H-® AlO2- + 2H20

The left-hand side of the formula contains four negative charges, and the right-hand side only one. Hence, in the course of the process, three electrons are given up:

Al + 40H- == AlO2- + 2H2O + 3e- (oxidation)

To obtain the reduction equation, we shall proceed from the scheme:


Here in the course of the process, the nitrogen atoms lose oxygen atoms and combine with hydrogen atoms. In a basic solution, this is possible with the participation of water molecules. Three water molecules are needed to combine with three oxygen atoms, and three more water molecules to form a molecule of NH3:

NO3 + 6H20 -- NH3 + 90H-

The total charge of the right-hand side of the formula is nine negative charges, and of the left-hand side—one. Consequently, eight electrons participate in the process:

NO3- + 6H2O + 8e- = NH3+ 90H- (reduction)

The ratio between the number of electrons liberated in oxidation and the number of electrons gained in reduction is 3:8 in the given example. Hence, to obtain the overall equation of the reaction, we must summate the equations of the oxidation and reduction processes, multiplying the first of them by 8 and the second by 3:

Example 3. catalytic oxidation of ammonia. This reaction is used in the production of nitric acid (see Vol. 2, Sec. 29). It is conducted at a temperature of about 750 °C. The skeleton equation is

The condensation of water vapour (steam) at 750 °C is impossible. We shall therefore not write the equations of the oxidation and reduction processes as we did for reactions proceeding in an aqueous solution—with the participation of water molecules, hydrogen or hydroxide ions. We shall only count the number of electrons participating in the oxidation and reduction. We shall take into account that the increase in the oxidation number of an element equals the number of lost electrons, and the decrease equals the number of gained electrons.

According to the skeleton equation, the oxidation number changes both in the nitrogen and in the oxygen. In the former, it grows from —3 to +2, and in the latter, it diminishes from 0 to -2. Let us write these changes as formulas, indicating the oxidation numbers of the relevant elements as superscripts. To avoid confusion with the charge of an ion, we shall use Roman numerals for this purpose

The ratio of the number of electrons gained in reduction to the number of electrons lost in oxidation is 4:5. Hence, five molecules of oxygen can oxidise four molecules of ammonia:

The reaction equations in the above three examples were compiled in a definite sequence. It can also be adhered to in other cases when compiling the equations of redox reactions. The sequence of operations is as follows:

1. Draw up a skeleton equation indicating the reactants and products.

2. Determine the oxidation numbers of the elements in the substances of the right-hand and left-hand sides of the equation; indicate the elements whose oxidation number changes.

3. Draw up the reduction and oxidation equations; find the ratio of the number of electrons gained in reduction to that lost in oxidation.

4. Summate the oxidation and reduction equations with account taken of the ratio of the number of electrons found in paragraph 3.

Most Important Oxidising and Reducing Agents

What substances can display the properties of oxidising agents, and what—of reducing agents? We have already mentioned that an oxidising agent contains an element whose oxidation number decreases, while a reducing agent contains an element whose oxidation number grows in the course of a reaction. Consequently, oxidising agents will include first of all compounds with the higher, and reducing agents will include compounds with the lower oxidation numbers featuring a given element.

Metals display only a positive oxidation state in their compounds, and their minimum oxidation number is zero. In other words, they have the minimum oxidation number only in the free state. Indeed, all free metals, although to a different extent, are capable of exhibiting only reducing properties. The reducing agents used in practice include aluminium, magnesium, sodium, potassium, and zinc. If a metal can have several oxidation numbers, those of its compounds in which it displays the lowest of them are also reducing agents, as a rule. Examples are the compounds of iron(II), tin(II), chromium(II), and copper(I).

Those compounds of metals can be oxidising agents in which the oxidation number is high and either equals the number of the group which the metal belongs to or is close to it. Practical use has been found, in particular, by an ammonia solution of silver oxide, an ammonia solution of copper(II) sulphate, mercury(II) chloride, lead dioxide PbO2, iron(III) chloride, potassium chromate and dichromate (K2CrO4 and K2Cr2O7), potassium permanganate KMnO4, and manganese dioxide MnO2.

Non-metals exhibit both positive and negative oxidation states. It is natural that compounds containing non-metals in their higher positive oxidation states can be oxidising agents, and compounds in which a non-metal displays a negative oxidation state can be reducing agents.

The most important reducing agents are hydrogen, carbon, and carbon monooxide.

Non-metals of the upper part of groups VI and VII of the periodic table are strong oxidising agents. The strong oxidising properties

of these substances are explained by the high electronegativity of their atoms. Fluorine has the strongest oxidising properties, but in practice oxygen, chlorine, and bromine are used most frequently as oxidising agents.

The compounds used as oxidising agents also include acids. Hydrochloric, sulphuric, and nitric acids have the greatest practical significance. The oxidising element in hydrochloric acid is hydrogen, in nitric acid it is nitrogen, in dilute sulphuric acid—hydrogen, and in the concentrated acid—sulphur. Hence, the equation of reduction with hydrochloric and dilute sulphuric acids has the form

Nitric acid, depending on its concentration, temperature, and the nature of the reducing agent, can be reduced to different oxidation numbers of the oxygen. One of the usual products of its reduction is nitrogen monoxide NO:

Various products may also be formed in reduction with concentrated sulphuric acid. One of them is sulphur dioxide:

Other compounds of non-metals used as oxidising agents are hydrogen peroxide, the salts of acids in which the acid-forming element exhibits a high oxidation number—chlorates (KClO3), per-chlorates (KClO4).

Oxidation-Reduction Duality. Intramolecular Oxidation-Reduction

Compounds with the maximum oxidation number of a given element can only play the role of oxidising agents in redox reactions. The oxidation number can only lower in this case. Conversely, compounds with the minimum oxidation number can only be reducing' agents. Here the oxidation number of an element can only grow. If an element is in an intermediate oxidation state, however, its atoms can, depending on the conditions prevailing, either take on or give up electrons. In the first case, the oxidation number of the element will lower, and in the second one it will grow. Consequently, compounds containing elements in intermediate oxidation states have oxidation-reduction duality—they are capable of entering into-reactions with either oxidising or reducing agents.

For example, nitrogen forms compounds in which its oxidation number changes from —3 (ammonia and ammonium salts) to 4-5 (nitric acid and its salts). The nitrogen in ammonia can only be a reducing agent, and that in nitric acid—only an oxidising agent. nitrous acid HNO3 and its salts, in which the oxidation number of. nitrogen is +3  however, enter into reactions with both strong oxidising and strong reducing agents. In the first case, it is oxidised to nitric acid, and in the second, it is usually reduced to nitrogen monoxide NO. We can exemplify the oxidation-reduction duality of nitrous acid by the reactions:

In addition to nitrous acid, sulphur, iodine, hydrogen peroxide, and a number of other substances have oxidation-reduction duality.

Substances containing an element in an intermediate oxidation state often have another characteristic property. It consists in that in definite conditions such a substance experiences a process in the course of which one part of the element is oxidised and the other part is reduced. This process is known as autoxidation-autoreduction.For instance, when chlorine reacts with water, a mixture of hydrochloric and hypochlorous (HOC1) acids is produced:

Here the chlorine undergoes both oxidation and reduction:

Autoxidation-autoreduction is also called disproportionation.Some compounds in definite conditions (usually when heated) experience intramolecular oxidation-reduction.In this process, one constituent part of the substance is an oxidising agent, and the other is a reducing agent. Examples of intramolecular oxidation-reduction are many processes of thermal dissociation. For instance, when water vapour dissociates:

the oxygen is oxidised (its oxidation number grows from —2 to 0), and the hydrogen is reduced (its oxidation number diminishes from +1 to 0).

Another example is the decomposition of ammonium nitrite employed in the laboratory to obtain pure nitrogen:

NH4NO2 = N2 + 2H2O

Here the ion NH3 is oxidised and the ion NO2- is reduced to free nitrogen.

Chemical Sources of Electrical Energy

We know that in any redox reaction, electrons pass from The reducing agent to the oxidising one. For example, when a zinc plate is lowered into a copper sulphate solution, the following reaction occurs:

Here the reducing agent—zinc—loses electrons. This half-reaction is

The oxidising agent —the copper ion—gains electrons. The equation of this half-reaction is

In the above example, both half-reactions occur at the place of contact of the zinc with the solution so that the electrons pass directly from the zinc atoms to the copper ions. We can conduct this reaction in such a way, however, that the oxidation and reduction half-reactions will be separated in space, and the electrons will pass from the reducing agent to the oxidising one not directly, but via a conductor of an electric current—through an external circuit. This directed flow of electrons is an electric current. When a redox reaction is conducted in such a way, its energy will be converted into electrical energy that can be used by connecting a consumer of electrical energy (for example, an electrical heating appliance or an electrical lamp) to the external circuit.

Devices used for the direct conversion of the energy of a chemical  reaction into electrical energy are called galvanic (or voltaic) cells,  They are also known as chemical sources of electrical energy or  chemical sources of current.

It is customary practice in engineering to apply the name galvanic  cell only to a. chemical source of current in which virtually irreversible reactions proceed. Such current sources cannot usually be

recharged: they are intended for use only once (in one or more stages). Chemical sources of current in which virtually reversible reactions " occur are called accumulators: they can be recharged and used repeatedly.

The functioning of any galvanic cell is based on the proceeding of a redox reaction in it. The simplest galvanic cell consists of two plates or rods made from different metals and immersed in a solution of an electrolyte. This system makes possible the separation of the redox reaction in space: oxidation occurs on one metal, and reduction on the other. Thus, the electrons are transferred from the reducing agent to the oxidising one via the external circuit.

Let us consider as an example a copper-zinc galvanic cell operating at the expense of the energy of the reaction between zinc and copper sulphate described above (Fig. 82). This cell (a Jacobi-Daniell cell) consists of a copper plate immersed in a solution of copper sulphate solution (a copper electrode) and a zinc plate immersed in a zinc sulphate solution (a zinc electrode). The two solutions are in contact with each other, but to prevent mixing, they are separated by a partition made from a porous material.

In operation of the cell, i.e. when the circuit is closed, the zinc becomes oxidised: on the surface of its contact with the solution, the zinc atoms transform into ions and after becoming hydrated pass into the solution. The liberated electrons travel through the external circuit to the copper electrode. The entire collection of these processes is depicted schematically by a half-reaction equation, or an electro-chemical equation:

Zn = Zn2+ + 2e-

At the copper electrode, reduction of the copper ions takes place. The electrons arriving here from the zinc electrode combine with the copper ions becoming dehydrated and leaving the solution; copper atoms are produced that separate in the form of metal. The corresponding electrochemical equation has the form Cu2++2e-= Cu

The net equation of the reaction proceeding in the cell is obtained when the equations of the two half-reactions are summated. Thus, in the operation of a galvanic cell, the electrons from the reducing agent pass to the oxidising agent through the external circuit, electrochemical processes occur at the electrodes, and the directed motion of the ions is observed in the solution.

The direction of motion of the ions in the solution is due to the electrochemical processes occurring at the electrodes. We have already indicated that at the zinc electrode the cations emerge into the solution, creating an excess positive charge in it, while at the copper electrode the solution, conversely, constantly becomes leaner in cations so that here the solution is charged negatively. The result is the setting up of an electric field in which the cations in the solution (Cu2+ and Zn2+) move from the zinc electrode to the copper one, while the anions— S042-— move in the opposite direction. As a result, the liquid at both electrodes remains electrically neutral.

The motion of the electrons and ions in the operation of a copper zinc cell is shown schematically inFig. 83.

The electrode at which oxidation occurs is called the anode. The one at which reduction occurs is called the cathode. In the copper zinc cell, the zinc electrode is the anode, and the copper one—the cathode.

The redox reaction proceeding in a galvanic cell is a complicated process. It includes the electrochemical steps proper (the transformations of atoms, ions, or molecules at the electrodes), the transfer of electrons, and the transfer of ions. All these steps are interlinked and proceed at the same rate. The number of electrons given up by the zinc in unit time equals the number of electrons received during this time by the copper ions. Hence, the rate of a reaction proceeding in a galvanic cell is proportional to the quantity of electricity transferred through the circuit in unit time, i.e. to the current in the circuit.

The electric current flowing in the external circuit of a galvanic cell can do useful work. But the work that can be done at the expense of the energy of a chemical reaction depends on its rate— it is the greatest when the reaction is conducted infinitely slowly, reversibly (see Sec. 67). Consequently, the work that can be done at the expense of the reaction proceeding in a galvanic cell depends on the magnitude of the current taken from it. If we lover the current in the external circuit to an infinitely small value by increasing the resistance of the circuit, the rate of the reaction in the cell will also be infinitely small, and the work will be maximum. The heat evolved in the internal circuit of the cell, on the contrary, will be minimum.

The work of an electric current is expressed by the product of the quantity of electricity ^flowing through, the circuit and the voltageV.

Fig. 83. Motion and electrons in the operation of a copper-zinc

In a copper-zinc cell, when one equivalent of zinc is oxidised and simultaneously one equivalent of copper ions is reduced, a quantity of electricity equal to one faraday (1 F =eNa = 96 485 coulombs*) wheree is the elementary electric charge andna is the Avogadro constant) will pass through the circuit. Hence, the useful workWwhich a current can do will be


whereF = Faraday's constant, numerically equal to one faradayV == voltage between the poles of the cell.

But since this work depends on the current, the voltage between the poles of the cell also depends on the current(F is a constant). In the limiting case corresponding to the reversible proceeding of the reaction, the voltage will be maximum. The maximum value of the voltage of a galvanic cell corresponding to the reversible proceeding of the reaction is called the electromotive force (e.m.f.) of the given cell.

For this limiting case, the useful work done by the current in a copper-zinc cell when one equivalent of zinc reacts with one equivalent of copper ions is expressed by the equation

whereE = Vmax is the e.m.f. of the cell.

It is clear that when one mole of zinc atoms reacts with one mole of copper ions, the equation becomes

In the general case upon the dissolving (or separation) of one mole of a substance whose ions have a charge of z, the maximum useful work is related to the e.m.f. by the equation

At a constant temperature and pressure, the maximum useful work of a reaction equals the change in the Gibbs energy AG taken with the reverse sign (see Sec. 67). Hence:

If the concentrations (more exactly, the activities) of the substances participating in a reaction equal unity, i.e. if standard conditions are observed, the e.m.f. of a cell is called its standard electromotive force and is designated by the symbolE°. The last equation

now becomes:

* In calculations, we shall use the value of this quantity approximated to three significant digits (96500 C/mol).

With a view to the standard change in the Gibbs energy of a reaction being associated with its equilibrium constant (see Sec. 68) by the expression

we get an equation relating the standard e.m.f. to the equilibrium constant of a reaction proceeding in a galvanic cell-

E.m.f. can be measured with a high accuracy. These measurements are one of the most accurate ways of finding the standard Gibbs energies and, consequently, the equilibrium constants for oxidation-reduction reactions in solutions.

Redox reactions proceed in a galvanic cell notwithstanding the fact that the oxidising and reducing agents are not in direct contact with each other. To gain an understanding of how this occurs, of how an e.m.f. appears when the processes of oxidation and reduction are separated in space, let us consider in greater detail the phenomena occurring at the phase interfaces in a galvanic cell.

Direct experiments with the use of radioactive indicators show that if we bring a metal (M) into contact with a solution of its salt, then the metal ions (Mz+ pass from the metal phase into the solution and from the solution into the metal. Since the energy state of the ions in these phases is not the same, then at the first moment after contact has been established, the metal ions pass from the metal into the solution and in the reverse direction at different rates. If transition of the ions from the metal phase into the solution predominates, the solution acquires a positive charge, while the metal electrode is charged negatively. As these charges increase, the transition of the cations into the solution having a like charge is hindered so that the rate of this process diminishes. On the other hand, the rate of transition of the cations from the solution onto the negatively charged electrode grows. As a result, the rates of the two processes level out, and equilibrium sets in between the metal and the solution:

The metal electrode is charged negatively, and the solution positively. If, when contact is established between a metal and a solution, the rate of transition of the cations from the metal into the solution was lower than the rate of their transition in the reverse direction, equilibrium also sets in between the electrode and the solution. Here, however, the electrode is charged positively, and the solution negatively.

In the Jacobi-DanielI cell, the corresponding equilibria set in between the zinc electrode and the zinc sulphate solution:

Zn2+ (metal) == Zn2+ (solution of ZnS04)

and also between the copper electrode and the copper sulphate solution:

Cu2+ (metal) == Cu2+ (solution of CuS04)

This cell has two other phase interfaces: between the solutions of zinc and copper sulphates, and also between the copper and the zinc(seeFig. 82). The interface between the solutions does not appreciably affect either the magnitude of the e.m.f. or the proceeding of the reaction when the cell is operating. As regards the interface between the metals, electrons can pass through it instead of ions as at a metal-solution interface*. Here too, owing to the different energy state of the electrons in the copper and in the zinc, the initial rates of electron transition from one metal into the other and in the reverse direction are different. In this case, equilibrium also sets in rapidly, however, and the metals also acquire charges of the opposite sign:

e~ (copper) ==e~ (zinc)

Consequently, when the circuit is open, equilibria set in on the three phase interfaces in the Jacobi-Daniell cell, the phases becoming charged. As a result, the energy state of the electrons at the ends of the open circuit becomes different: on the copper conductor in contact with the zinc electrode the Gibbs energy of the electrons is higher, and on the end connected to the copper electrode it is lower. It is exactly the difference, between the Gibbs energies of the electrons at the ends of the circuit that determines the e.m.f. of a given cell.

When the external circuit is closed, the electrons move from the zinc electrode to the copper one. Therefore, equilibria at the phase interfaces are violated. A directed transition of the zinc ions takes place from the metal into the solution, of the copper ions—from the solution into the metal, and of the electrons—from the zinc to the copper. A redox reaction occurs.

In principle, any redox reaction can yield electrical energy. The number of reactions practically used in chemical sources of electrical energy, however, is not great. This is associated with the circumstance that not any redox reaction makes it possible to create a galvanic cell having technically valuable properties (a high and virtually constant e.m.f., the possibility of obtaining high currents, a long lifetime, etc.). In addition, many redox reactions require the use of costly substances.

Unlike the copper-zinc cell, all modern galvanic cells and accumulators use not two, but one electrolyte. Such current sources are considerably more convenient in operation. For example, in lead ac-


The diffusion of atoms and ions from a metal into a metal occurs much more slowly and does not virtually affect the establishment of equilibrium at the interface between the metals.


cumulators, a sulphuric acid solution istheelectrolyte.

In virtually all the galvanic cells produced at present, the anode is made from zinc, while oxides of less active metals are used as the substance for the cathode.

For a description of the most important galvanic cells and of accumulators see.

Chemical sources of electrical energy arc used in various branches of engineering. In means of communication (radio, telephone, telegraph) and in electrical measuring' apparatus, they are sources of electrical power, on motor vehicles and aeroplanes they are used for actuating starters and other equipment, on transport and in many other fields they are used in portable lanterns for illumination. With the growing shortage of petroleum products in many countries, greater and greater attention is being given to the development of electric vehicles powered by accumulators.

All conventional chemical current sources are not free of two shortcomings. First, the cost of the substances needed for their operation (for example, lead and cadmium) is high. Second, the ratio of the amount of energy that a cell can give up to its mass is low. The last few decades have seen the conducting of investigations aimed at creating cells whose operation would require the consumption of cheap substances having a low density, similar to liquid or gaseous fuel (natural gas, kerosene, hydrogen, etc.). Such galvanic cells are known as fuel cells. Much attention is being given to the fuel cell problem at present, and it can be assumed that in the nearest future fuel cells will find broad application.

Electrode Potentials

Every oxidation-reduction reaction consists of oxidation and reduction half-reactions. When a reaction proceeds in a galvanic cell or is conducted by electrolysis, each half-reaction occurs at the corresponding electrode. This is why half-reactions are also called electrode processes.

We showed earlier that the e.m.f.E of a galvanic cell corresponds to the redox reaction proceeding in this cell. The e.m.f. is related to the change in the Gibbs energy of the reaction AG by the equation

In accordance with the division of a redox reaction into two half-reactions, it is customary practice to represent the e.m.f.'s too in the form of the difference between two quantities, each of which corresponds to the given half-reaction. These quantities are known as the electrode potentials.

For a copper-zinc cell, the reaction proceeding during its operation:

Zn + Cu2+ == Zn2++  Cu

s divided into two half-reactions:

Cu2+ + 2e- = Cu

Zn = Zn2+ + 2e-

Accordingly, the e.m.f. of this element(E} can be represented as the difference between the electrode potentials ((j), one of which (j1) corresponds to the first, and the other (j2) to the second of the above half-reactions:

The change in the Gibbs energy AG1 that corresponds to the thermo-dynamically reversible reduction of one mole of copper ions is

while the change in the Gibbs energy AG2 corresponding to the ther-modynamically reversible oxidation of one mole of zinc atoms is

In the general case, the electrode potentialj and a change in the Gibbs energy AG equal to

correspond to any electrode process Ox +ze- = Red

where Ox and Red stand for the oxidised and reduced forms, respectively, of the substances participating in the electrode process.

In the following, when dealing with electrode processes, we shall write their equations in the direction of reduction (except, naturally, when we are speaking exactly about oxidation).

Investigation of the potentials of various electrode processes has showTi that their magnitudes depend on the following three factors:

(1) the nature of the substances participating in the electrode process;

(2) the ratio between the concentrations* of these substances, and

(3) the temperature of the system. This relationship is expressed by the equation:

* Strictly speaking, the magnitude of an electrode potential depends on the ratio of the activities and not of the concentrations of substances. In all the following equations, the activity should be substituted for the concentration. But at low concentrations of solutions,the error introduced by substituting the concentration for the activity is cot great.

Herej° is the standard electrode potential of a given process— a constant whose physical meaning is considered below; R is the molar gas constant;T is the absolute temperature; z is the number of electrons participating in the process;F is the Faraday constant;

[Ox] and [Red] are the products of the concentrations of the substances participating in the process in the oxidised (Ox) and reduced (Red) forms.

The physical meaning of the quantityj° becomes clear if we consider the case when the concentrations (activities) of all the substances participating in a given electrode process are unity. For this condition, the second addend in the right-hand side of the equation vanishes (log 1=0), and the equation becomes

Concentrations (activities) equal to unity are called standard concentrations (activities). Therefore, the potential corresponding to this case is also called the standard potential. Thus,the standard electrode potential is the potential of a given electrode process at concentrations (more accurately, at activities)of all the substances participating in it equal to unity.

Hence, the first addend in the equation of the electrode potential (j°) takes into account the influence of the nature of the substances on its magnitude, and the second addend 2.3RT/zF log[Ox]/[Red] the influence of their concentrations. In addition, both terms vary with the temperature.

For the customary standard temperature used in electrochemical measurements (25 °C == 298 K), when the values of the constant quantities are introduced [R •= 8.31 J/(mol-K),F == 96 500 C/mol], the equation becomes

To construct a numerical scale of the electrode potentials, it is necessary to assume that the potential of a definite electrode process is zero. The following electrode process has been taken as the standard for constructing such a scale:

2H+ +2e- == H2

The change in the Gibbs energy associated with the proceeding of this half-reaction in standard conditions is adopted equal to zero. Accordingly, the standard potential of this electrode process is also taken equal to zero. All electrode potentials indicated in this book, and also in the majority of other modern publications, are expressed according to this hydrogen scale.

The above electrode process is carried out on a hydrogen electrode.The latter is a platinum plate electrolytically coated with spongy platinum arid immersed in a solution of an acid through which

hydrogen is passed (Fig. 84). The hydrogen dissolves well in the platinum; the hydrogen molecules partly decompose into atoms (the platinum catalys.es this decomposition). Oxidation of the hydrogen atoms or reduction of the hydrogen ions can proceed on the surface of contact of the platinum with the acid solution. The platinum does not virtually participate in the electrode reactions and plays, as it were, the role of a sponge impregnated with atomic hydrogen.

The potential of the hydrogen electrode is reproduced with a very high accuracy. This is exactly why the hydrogen electrode has been taken as the standard in creating a scale of electrode potentials.

Let us establish the form of the general equation of the electrode potential for the hydrogen electrode. For this electrode, z == 2, [Ox] = [H+]2, [Red] = [H2]. The concentration of the hydrogen dissolved in the platinum is proportional to its partial pressureP (H2):

[H2] =kp(H2)

wherek is a quantity that is constant at the given temperature. Using the equation of an electrode process and introducing the constantk into the value ofj°, we get

The partial pressure of hydrogenP (H2) is usually kept equal to standard atmospheric pressure, which is conventionally taken as unity. In this case, the last term of the equation obtained vanishes (log 1 = 0). Hence:

Since the standard potential of the process being considered is assumed to equal zero, then

or since log [H+] == —pH, we finally get

To find the potential of an electrode process, it is necessary to form a galvanic cell from the electrode being tested and a standard hydrogen electrode and measure its e.m.f. Seeing that the potential of a standard hydrogen electrode is zero, the measured e.m.f. will be the potential of the given electrode process.

A standard hydrogen electrode is not generally used in practice as a reference electrode because this involves considerable complications. More convenient electrodes are employed whose potentials in comparison with the standard hydrogen electrode are known. The e.m.f. of the cell should be calculated by the equation

E= ljref-—jxl

HereE is the e.m.f. of the cell,jref is the known potential of the reference electrode, andjx is the potential of the electrode being tested.

The reference electrodes in greatest favour are the silver-silver chloride and the calomel electrodes. The former is a silver wire coated with a layer of AgCl and immersed in a solution of hydrochloric

acid or its salt. The following reaction proceeds in it when the circuit is closed:

AgCl +e- = Ag+ + Cl-

The calomel electrode consists of mercury coated with a suspension of calomel (Hg2Cl2) in a solution of KC1. The potentials of these electrodes are reproduced with a high accuracy.Figure 85 shows a circuit with a calomel electrode.

To find the value of the electrode potential, it is necessary to measureaot the voltage of the operating cell, but its e.m.f. When measuring the latter, the resistance of the external circuit (i.e. of the measuring device) is very high. Virtually no reaction proceeds in the cell. Thus, the electrode potentials correspond to the reversible occurring of processes or, which is the same, to the state of electrochemical equilibrium at the electrodes.

Consequently, electrode potentials are often called equilibrium electrode potentials or simply equilibrium potentials.

Let us now consider the form which the general equation of the electrode potential will acquire in the most important cases.

1. The electrode process is expressed by the equation

M+ + ze    = M

where M stands for the atoms of a metal, and M** are its z-charged ions.

This case includes both electrodes of a copper-zinc cell and in general any metal electrode in a solution of a salt of the same metal. Here the metal ions are the oxidised form of the metal, and its atoms are the reduced form. Hence, [Ox] = [Mz+], and [Red] == const because the concentration of the atoms in a metal is a constant quantity at a constant temperature. Including the value of this constant in the quantityj°, we get:

For instance, for the process Ag+ + e- = Ag, we have:

and for the process Zn2+ +2e~ = Zn, we get:

2. The electrode process is expressed by the equations

Mz2+ + (z2-z1)e- =Mz1+

Here both the oxidised (Mz2+) and reduced (Mz+) forms of the metal are in the solution, and their concentrations are variable quantities. Consequently,

In this and in the cases treated below, the electrode at which the electrode process occurs is made from an inert material. Most often, this material is platinum.

We have considered examples when only ions consisting of one element participated in the electrode processes. Quite often, however, the substance becoming oxidised or reduced consists not of one, but of two or more elements. The oxidising agent most frequently contains oxygen; here water and the products of its dissociation— hydrogen ions (in an acid solution) or hydroxide ions (in a basic solution) generally participate in the electrode process. Let us see what the equations of the electrode process potentials will be like in such cases.

3. The electrode process is expressed by the equation:

02 + 4H+ + 4e- = 2H20

This half-reaction (when it proceeds in the direction of reduction) plays a very great part in the corrosion of metals. Oxygen is the most widespread oxidising agent causing metals to corrode in aqueous solutions.

In the above electrode process, oxygen is reduced with the participation of hydrogen ions to form water. Hence, [Red] = [H2O]2, and [Ox] = [O2] [H+]4. The concentration of the water in dilute solutions may be considered constant. The concentration of oxygen in a solution is proportional to its partial pressure over the solution {[O2] =kp(O2)}. After performing the required transformations and designating the sum of the constant quantities byj°, we get

For the process being considered,j ==1.228 V; hence,

At a partial pressure of the oxygen equal to standard atmospheric pressure (which is conventionally taken equal to unity), logp (O2) = = 0, and the last equation becomes

4. For electrode processes represented by more complicated equations, the expressions for the potentials contain a greater number of  variable concentrations. Let us consider, for example, the electrode  process

MnO4- + 8H+ + 5e- = Mn2+ + 4H2O                            |

This half-reaction proceeds (in the direction of reduction) when potassium permanganate reacts with most reducing agents in an acid solution.

The concentrations of all the substances participating in this electrode process (except water) are variable quantities. For this-process,j0 == 1.507 V. The equation of the electrode potential has the form

Examples 3 and 4 show that for electrochemical processes occurring with the participation of water, the hydrogen ion concentration is in the numerator of the logarithmic term of the potential equation. For this reason, the electrode potentials of such processes depend on the pH of the solution and their values grow with an increasing acidity of the solution.

We have already mentioned that the dependence of the electrode potential on the nature of the substances participating in an electrode process is described by the quantityj°. In this connection, it is customary practice to arrange all electrode processes in a series according to the magnitude of their standard potentials. Table 18 gives the equations of the most important electrode processes and the relevant electrode potentials in the order of increasing values ofjo.

The position of an electrochemical system in tills series characterises its oxidation-reduction ability. By an electrochemical systemhere is meant the collection of all the substances participating in the given electrode process.

The oxidation-reduction ability characterises an electrochemical system, but the oxidation-reduction ability of a substance (or ion) is also often spoken about. It must be borne in mind, however, that many substances can be oxidised or reduced to different products. For instance, potassium permanganate (the ion MnO4-) can, depending on the conditions and first of all on the pH of the solution, be reduced either to the ion Mn2+, or to MnO2, or to the ion Mn042-.

The corresponding electrode processes are expressed by the equations:

MnO4- + 8H+ + 5e- == Mn2+ + 4H20

MnO4-  + 4H+ +3e- = MnO2 + 2H2O

MnO4- +e- = MnO42-

Since the standard potentials of these three electrode processes are different (see Table 18), the positions of these three systems in the series ofj0 are also different. Hence, the same oxidising agent (MnO4-) can occupy several positions in the series of standard potentials.

Elements exhibiting only one oxidation number in their compounds have simple oxidation-reduction characteristics and occupy a small number of positions in the series of standard potentials. They chiefly include the metals of the main subgroups of Groups I-III of the periodic table. On the other hand, many positions in the series ofj° are occupied by the elements that form compounds with different oxidation numbers—non-metals and many metals of the auxiliary subgroups of the periodic table.

The series of standard electrode potentials makes it possible to solve the problem about the direction of spontaneous redox reactions

as in the general case of any chemical reaction, the determining factor here is the sign of the change in the Gibbs energy of the reaction. If we form a galvanic cell from two electrochemical systems, then in its operation the electrons will pass spontaneously from the negative pole of the cell to the positive one, i.e. from an electrochemical system with a lower value of the electrode potential to a system with a higher value of it. But this signifies that the first system will play the role of a reducing agent, and the second that of an oxidising agent. Consequently, in a galvanic cell,an oxidation-reduction. reaction can proceed spontaneously in a direction in which the electrochemical system with a higher value of the electrode potential plays the role of an oxidising agent, i.e. becomes reduced. Upon the direct reaction of the substances, the possible direction of the reaction will naturally be the same as when it is conducted in a galvanic cell.

If the oxidising and reducing agents are far from each other in the series ofj°, the direction of the reaction is practically completely determined by their mutual position in the series. For example, zinc (j° = —0.763 V) will displace copper (j° = +0.337 V) from an aqueous solution of its salt at any practically achievable concentration of the solution. If, on the other hand, the values ofj° for the oxidising and reducing agents are close to each other, then when solving the problem of the direction in which the reaction will proceed spontaneously, it is also necessary to take account of how the concentrations of the relevant substances affect the electrode potentials. For example, the reaction

Hg22+ +  2Fe2+ == 2Hg + 2Fe3+

Table 18

Electrode Potentials in Aqueous Solutions at 25 °C and at a Partial Pressure of the Gases Equal to Standard Atmospheric Pressure

can proceed spontaneously either from the left to the right or from the right to the left. Its direction is determined by the concentrations of the iron and mercury ions. Two electrochemical systems participate in this reaction:

Hg22+ +2e-=2Hg                                          (1)

2Fe3++e-=Fe2+                                              (2)

The following potentials correspond to the relevant electrode


Let us calculate the values ofj1andj2 at [Hg22+] = [Fe2*] = 10-1and [Fe3+l = 10-4 mol/1000 g ofH2O:

Oxidation-Reduction Reactions. Electrochemistry

Thus, with the ratio of the concentrations we have taken,j1 >j2and the reaction proceeds from left to right.

Now let us calculatej1 andj2 for the reverse ratio of the concentrations. Assume that [Hg22+] - [Fe2+] = 10-4, and [Fe3+] = - 10-1 mol/1000 g of H2O:

Consequently, at these concentrations,j1 >j2, and the reaction proceeds from right to left.

If a redox reaction proceeds with the participation of water and hydrogen or hydroxide ions, then the pH of the solution must also be taken into consideration.

Table 18 includes 38 half-reactions. Combining them with oneanother, we can solve the problem of the directions of the spontaneous proceeding of (38 X 37)/2 = 703 reactions.

Example. Establish the direction in which the following reaction can proceed:

We shall write the equation of the reaction in the net ionic form.

In Table 18, we find the standard electrode potentials for the electrochemical

systems participating in the reaction:

The oxidising agent is always the electrochemical system with a higher value of the electrode potential. Since herej0 is considerably greater thanj??, then at virtually any concentrations of the reacting substances the bromide ion will be the reducing agent and will be oxidised by the lead dioxide: the reaction will proceed spontaneously from left to right.

The farther a system is in the series of standard potentials, i.e. the higher its standard potential, the stronger an oxidising agent is its oxidised form. And, conversely, the earlier a system appears in the series, i.e. the smaller the value of (p°, the stronger a reducing agent is its reduced form. Indeed, among the oxidised forms of systems at the bottom of the series, we find such strong oxidising agents as f2, H2O2, and Mn04-. The strongest reducing agents, on the other hand, are the reduced forms of the systems at the top of the series; the alkali and alkaline-earth metals.

When redox reactions proceed, the concentrations of the reactants diminish, and of the products grow. The result is a change in the

magnitudes of the potentials of both half-reactions: the electrode potential of the oxidising agent drops, and that of the reducing agent grows. When the potentials of both processes become equal to each other, the reaction terminates—a state of chemical equilibrium sets in.

Electromotive Series of Metals

If from the entire series of standard electrode potentials we separate only the electrode processes that correspond to the general equation

Mz++ ze- = M

we get an electromotive (or activity) series of metals. In addition to metals, hydrogen is included in this series, which permits us to see what metals are capable of displacing hydrogen from aqueous solutions of acids. The electromotive series for the most important metals is given in Table 19. The position of a metal in the series character

Table 19Electromotive Series of Metals


ises its ability to participate in oxidation-reduction reactions in aqueous solutions in standard conditions. Ions of the metals are oxidising agents, and the metals in the form of elementary substances are reducing agents. The farther a metal is in the electromotive series, the stronger an oxidising agent in an aqueous solution are its ions. Conversely, the nearer a metal is to the top of the series the stronger are the reducing properties exhibited by the elementary substance—metal. The potential of the electrode process

2H+ + 2e- = H2

in a neutral medium (pH = 7) is —0.059 x 7 = —0.41 V The active metals at the top of the series having a potential that is considerably more negative than —0.41 V displace hydrogen from water. Magnesium displaces hydrogen only from hot water. The metals between magnesium and cadmium do not usually displace hydrogen from water. The surfaces of these metals become covered with oxide films having a protective action*.

The metals between magnesium and hydrogen displace hydrogen from solutions of acids. The surfaces of some metals also become covered with protective films that inhibit the reaction. For example, the oxide film on aluminium makes this metal stable not only in water, but also in solutions of some acids. Lead does not dissolve in sulphuric acid when its concentration is below 80% because the salt PbS04 formed when lead reacts with sulphuric acid is insoluble and produces a protective film on the surface of the metal. The phenomenon of the deep inhibition of the oxidation of a metal due to the presence of protective oxide or salt films on its surface is called passivation, and the state of the metal is called the passive state.

Metals are capable of displacing one another from solutions of salts. The direction of the reaction is determined by their relative position in the electromotive series. When considering specific cases of such reactions, one must remember that active metals displace hydrogen not only from water, but also from any aqueous solution. Consequently, the mutual displacement of metals from solutions of their salts occurs in practice only with metals below magnesium in the series.

The displacement of metals from their compounds by other metals was first studied in detail by N. Beketov**. As a result of his investi-


* The potentials of metals in water naturally differ from their potentials in standard conditions, in the majority of cases they have a more negative value. But, as a rule, this does not affect the correctness of our conclusions on the ability of metals to displace hydrogen from water.

** Nikolai Nikolaevich Beketov (1826-1911) was an eminent Russian scientist— a physicochemist. Beketov's outstanding work is hisInvestigations of Phenom' ena of the Displacement of Elements by Other Elements published in 1865.


gations, he arranged the metals in a "displacement series" according to their chemical activity. This series was the prototype of the electromotive series of metals.

The relative position of some metals in the electromotive series and in the periodic table, at first sight, is not the same. For example, according to its position in the periodic table, the chemical activity of potassium should be greater than that of sodium, and of sodium greater than that of lithium. In the electromotive series, however, lithium is the most active, while potassium occupies a position between lithium and sodium. Zinc and copper as regards their position in the periodic table should have an approximately identical chemical activity, but in the electromotive series, zinc is considerably ahead of copper. The reason for such discrepancies is as follows.

In comparing metals occupying definite positions in the periodic table, the magnitude of the ionisation energy of free atoms is taken as a measure of their chemical activity—their reducing power. Indeed, as we pass, for example, down the main subgroup of group I of the periodic table, the ionisation energy of atoms diminishes, which is associated with an increase in their radii (i.e., with a greater distance to the outer electrons from the nucleus) and with an increasing screening of the positive charge of the nucleus by intermediate electron layersConsequently, potassium atoms display a greater chemical activity—have stronger reducing properties— than sodium atoms, and sodium atoms display a greater activity than lithium atoms.

When comparing metals in the electromotive series, the measure of their chemical activity is the work of transforming a metal in the solid state into hydrated ions in an aqueous solution. This work can be represented as the sum of three addends: the energy of atomisation—the transformation of a metal crystal into isolated atoms, the energy of ionisation of the free atoms of the metal, and the energy of hydration of the ions formed. The atomisation energy characterises the strength of the crystal lattice of a given metal. The ionisation energy of atoms—the energy needed to detach valence electrons from them—is directly determined by the position of the metal in the periodic table. The energy evolved in hydration depends on the electronic structure of an ion, its charge and radius. Ions of lithium and potassium, having an identical charge, but different radii, will set up different electric fields around themselves. The field set up near small lithium ions will be stronger than that near big potassium ions. It is thus clear that lithium ions will become hydrated with the liberation of more energy than potassium ions.

Thus, in the course of the transformation being considered, energy is spent for atomisation and ionisation, and energy is evolved in hydration. When the summary expenditure of energy is smaller, the some electrochemical processes. Such cases will be given special mention below.

In considering cathode processes occurring in the electrolysis of aqueous solutions, we shall limit ourselves to the most important case—cathode reduction leading to separation of the elements in the free state. Here account must be taken of the magnitude of the potential in the process of hydrogen ion reduction. This potential depends on the concentration of the hydrogen ions and for neutral solutions (pH =7) it isj = —0.059 X 7 == —0.41 V. Therefore, if the cation of an electrolyte is a metal whose electrode potential is considerably more positive than —0.41 V, the metal will separate from a neutral solution of such an electrolyte at the cathode. In the electromotive series, such metals are near hydrogen (beginning approximately with tin) and after it. Conversely, if a metal having a potential considerably more negative than —0.41 V is the cation of an electrolyte, the metal will not be reduced, but hydrogen will be liberated. Such metals include the ones at the top of the electromotive series, approximately up to titanium. Finally, if the potential of a metal is close to the value —0.41 V (metals in the middle part of the series—Zn, Cr, Fe, Cd, Ni), then depending on the concentration of the solution and the conditions of electrolysis both reduction of the metal and the liberation of hydrogen are possible. Often the joint deposition of the metal and liberation of hydrogen are observed.

The electrochemical liberation of hydrogen from acid solutions occurs owing to the discharge of the hydrogen ions. For neutral or alkaline media, it is the result of the electrochemical reduction of water:

2H20 +2e- = H2 + 20H-

Thus, the nature of the cathode process in the electrolysis of aqueous solutions is determined first, of all by the position of the relevant metal in the electromotive series. In a number of cases, the pH of the solution, the concentration of the metal ions, and other conditions of electrolysis are of great importance.

In considering anode processes, we must remember that the material of the anode may become oxidised in the course of electrolysis. In this connection, electrolysis with an inert anode and electrolysis with an active anode are distinguished. An anode is called inertif its material does not become oxidised in electrolysis. An anode is called active if its material may become oxidised in the course of electrolysis. Most frequently, graphite, carbon, and platinum are used as the materials for inert anodes.

* The most important conditions of electrolysis include the current density, the temperature, and the composition of the solution. The current density is defined as the ratio of the current to the working surface area of an electrode.The electrochemical oxidation of water with the liberation of oxygen occurs at an inert anode in the electrolysis of aqueous solutions of alkalies, oxyacids and their salts, and also hydrogen fluoride and fluorides. Depending on the pH of the solution, this process goes on differently and can be expressed by different equations. For an alkaline medium, the equation is

40H- == O2 + 2H2O + 4e-

and for an acid or a neutral one, it is

2H20 == O2 + 4H+ + 4e-

In the above cases, the electrochemical oxidation of water is the most advantageous process from the energy viewpoint. Oxygen-containing anions are either not capable of becoming oxidised, or their oxidation occurs at very high potentials. For instance, the standard potential of oxidation of the ion S042-

2SO42- == S2O82- + 2e-

is 2.010 V, which considerably exceeds the standard potential of oxidation of water (1.228 V). The standard potential of oxidation of the ion F- has a still higher value (2.87 V).

In the electrolysis of aqueous solutions of acids not containing oxygen and their salts (except for HF and fluorides), the anions are discharged at the anode. Particularly, in the electrolysis of solutions of HI, HBr, HCl and their salts, the corresponding halogen is liberated at the anode. We must note that the liberation of chlorine in the electrolysis of HCl and its salts contradicts the mutual position of the systems

in the series of standard electrode potentials. Thi5 anomaly is connected with the considerable overvoltage (see Sec. 104) of the second of these two electrode processes—the material of the anode has an inhibiting action on the process of oxygen liberation.

With an active anode, the number of competing oxidising processes grows to three: the electrochemical oxidation of water with the liberation of oxygen, discharge of the anion (i.e. its oxidation), and electrochemical oxidation of the anode metal (the anode dissolvingof the metal). Of these possible processes, the one will occur that is more profitable from the energy viewpoint. If the metal of the anode is above the other two electrochemical systems in the series of standard potentials, then anode dissolving of the metal will be observed. Otherwise, oxygen will be liberated or the anion will be discharged.

Let us consider some typical cases of electrolysis of aqueous solutions.

The Electrolysis of a CuCl2 Solution with an Inert Anode. Copper is below hydrogen in the electromotive series. Consequently, the discharge of the Cu2+ ions and the deposition of metallic copper willoccur at the cathode. The chloride ions will discharge at the anode.

The scheme of electrolysis* of a solution of copper(II) chloride is as follows:

Electrolysis of a K2SO4 Solution with an Inert Anode. Since potassium is much higher than hydrogen in the electromotive series, the liberation of hydrogen and the accumulation of OH~ ions will occur at the cathode. The liberation of oxygen and the accumulation of H+ ions will occur at the anode. At the same time, K+ ions will arrive at the cathode space, and S042- ions at the anode space. Thus, the solution will remain electrically neutral in all its parts. An alkali will accumulate in the cathode space, however, and an acid in the anode space.

The scheme of electrolysis of a potassium sulphate solution is:

Electrolysis of an NiS04 Solution with a Nickel Anode. The standard potential of nickel (—0.250 V) is somewhat higher than —0.41 V. Hence, in the electrolysis of a neutral solution of NiSO4, discharge of the Ni2+ ions and deposition of the metal will mainly occur at the cathode. The opposite process—oxidation of the metal—occurs at the anode because the potential of nickel is much lower than the potential of oxidation of water and, moreover, than the potential of oxidation of the S042- ion. Hence, in the given case, electrolysis consists in dissolving of the anode metal and its deposition at the cathode.

* In this and the following schemes, the boxes contain the formulas of the substances that are the products of electrolysis. The scheme of electrolysis of a nickel sulphate solution is

This process is used for the electrical purificationof nickel (electrolytic refining, see Sec. 103).

Laws of Electrolysis

Electrolysis was first studied quantitatively in the thirties of the 19th century by the outstanding British scientist M. Faraday* who established the following laws of electrolysis:

1.The mass of a substance formed in electrolysis is proportional to the amount of electricity that has passed through the solution.

This law follows from the essence of electrolysis. We have already mentioned that an electrochemical process—the reaction of the ions or molecules of the electrolyte with the metal electrons—occurs at the place of contact of the metal with the solution, so that the electrolytic formation of a substance is a result of this process. It is quite obvious that the amount of a substance obtained at an electrode will always be proportional to the number of electrons passing through the circuit, i.e. to the amount of electricity.

2.In the electrolysis of different chemical compounds, equal amounts of electricity result in the electrochemical transformation of equivalent amounts of substances.

Assume, for example, that an electric current consecutively passes through solutions of hydrochloric acid, silver nitrate, copper(II) chloride, and tin(IV) chloride (Fig. 86). After a certain time, the

* Michael Faraday (1791-1867) was one of the most eminent British physicists and chemists. The major part of his works relate to the field of electricity. He established the laws of electrolysis, and discovered the phenomenon of electromagnetic induction. He was the first to obtain a number of gases (chlorine, ammonia, etc.) in the liquid state. He discovered benzene and isobutylene.

amounts of the electrolysis products are determined. It is found that in the time during which one gram of hydrogen, i.e. one mole of atoms is liberated from a solution of hydrochloric acid, the masses of the metals indicated below will be deposited from the other solutions:

Comparing the deposited masses of the metals with their atomic masses, we find that one mole of silver atoms, one-half mole of copper atoms, and one-fourth mole of tin atoms are deposited. In other words, the amounts of substances formed at the cathode equal their equivalents. The same result is obtained by measuring the amounts of the substances liberated at the anode. Thus, in each of the first, third and fourth cells, 35.5 g of chlorine are liberated, and in the second cell 8 g of oxygen. It is not difficult to see that here too the substances are liberated in amounts equal to their equivalents.

In considering the second law of electrolysis from the viewpoint of the electron theory, it is not difficult to understand why substances are liberated in chemically equivalent amounts in electrolysis. Let us take as an example the electrolysis of copper(II) chloride. When copper is deposited from the solution, each of its ions receives two electrons from the cathode, and at the same time two chloride ions give up electrons to the anode, transforming into chlorine atoms. Consequently, the number of deposited copper atoms will always be half the number of liberated chlorine atoms, i.e. the ratio of the masses of the copper and chlorine will equal the ratio of their equivalent masses.

Measurements have established that the amount of electricity causing the electrochemical transformation of one equivalent of a substance is 96 485 (approximately 96 500) -coulombs. This amount of electricity is called the faraday or Faraday's constant and is designated by the symbolF.

The second law of electrolysis gives a direct way of determining the equivalents of various elements. Calculations associated withelectrochemical production processes are based on the same law.

The laws of electrolysis relate to the electrolysis of solutions, melts, and solid electrolytes with purely ionic conductivity.

Electrolysis in Industry

Electrolysis finds a very important application in the metallurgical and chemical industries and in electrodeposition.

In the metallurgical industry, electrolysis of molten compounds and aqueous solutions is used to produce metals, and also for electrolytic refining—the purification of metals from harmful impurities and the extraction of valuable components.

The electrolysis of melts is used to produce metals having strongly negative electrode potentials, and some of their alloys.

At a high temperature, the electrolyte and the electrolysis products may enter into a reaction with one another, with air, and also with the materials of the electrodes and the electrolyser. As a result, this scheme of electrolysis, which is simple in principle (for instance, the electrolysis of MgCl2 in the production of magnesium) becomes more intricate.

Usually, not individual molten compounds, but their mixtures are used as the electrolyte. A very important advantage of mixtures is their relatively low melting point allowing electrolysis at a lower temperature.

At present, electrolysis of melts is used to produce aluminium, magnesium, sodium, lithium, beryllium, and calcium. It is virtually not used to produce potassium, barium, rubidium, and cesium owing to the high chemical activity of these metals and their high solubility in the molten salts. In recent years, the electrolysis of molten media has been acquiring some significance for the production of certain refractory metals.

The electrolytic separation of a metal from a solution is calledelectrical extraction. The ore or concentrate is treated with definite reagents to transfer the metal into a solution. After purification, the solution is subjected to electrolysis. The metal is deposited at the cathode and in the majority of cases is highly pure. This method is used mainly to produce zinc, copper, and cadmium.

Metals are subjected to electrolytic refining for removing impurities from them and for transferring the components contained in them into products convenient to process. The metal to be purified is cast into plates, and they are placed as anodes in an electrolyser. The passage of a current through the circuit causes the anode metal to dissolve—it passes into solution in the form of cations. The metal cations then become discharged at the cathode and form a compact deposit of pure metal. The impurities in the anode either remain undissolved, settling in the form of anode mud, or pass into the electrolyte, whence they are periodically or continuously withdrawn.

Let us consider as an example the electrolytic refining of copper. The main component of the solution is copper sulphate—the most widespread and cheapest salt of this metal. But a solution of CuS04has a low electrical conductivity. To increase it, sulphuric acid is added to the electrolyte. Small amounts of additions facilitating the production of a compact deposit of the metal are also introduced into the solution.

The metallic impurities contained in unrefined ("blister")copper can be divided into two groups:

(1) Fe, Zn, Ni, Co. These metals have considerably more negative electrode potentials than copper. Therefore, they are dissolved fromthe anode together with the copper, but are not deposited at thecathode. They accumulate in the electrolyte, and in this connection the latter is periodically purified.

(2) Au, Ag, Pb, Sn. The noble metals (Au, Ag) do not undergo anode dissolving, and in the course of the process settle near the anode, forming together with other impurities an anode mud that it periodically extracted. Tin and lead, on the other hand, dissolve together with the copper, but form poorly soluble compounds in the electrolyte that precipitate and are also removed.

Copper, nickel, lead, tin, silver, and gold are subjected to electrolytic refining.

Electrodeposition includes electroplating and galvanoplastics.Electroplating processes involve the application of other metals on the surface of metal articles by electrolysis to protect the articlesfrom corrosion, to impart hardness to their surfaces, and also fordecorative purposes. Among the numerous electroplating processes used in engineering, the most important are chromium plating, zinc plating (galvanisation), and nickel plating.

The essence of electroplating is as follows. The well cleaned and degreased component to be protected is immersed in a solution containing a salt of the metal which it is to be coated with, and is connected as the cathode to a direct-current circuit. When a current flows through the circuit, a coat of the protecting metal is deposited onto the component. The best protection is provided by finely crystalline dense deposits.  Such deposits also have better mechanical properties.

Galvanoplastics (or galvanoplasty) is the name given to processes of obtaining accurate metal copies of relief (embossed) objects by the electrodeposition of metal. Galvanoplastics is used to produce moulds and dies for pressing various articles (phonograph records, buttons, etc.), matrices for stamping leather and paper, printed radio engineering circuits, and stereotype plates. Galvanoplastics was discovered by the Russian academician B. Jacobi (1801-1874) in the thirties of the 19th century.

There are also other kinds of electrochemical treatment of metal surfaces such as the electrical polishing of steel and the oxidation of aluminium and magnesium. The latter consists in anode treatment of the metal in the course of which the structure of the oxide film on its surface is changed in a definite way. The result is improvement of the corrosion resistance of the metal. In addition, the metal acquires an attractive appearance.

In the chemical industry, electrolysis is used to prepare various products. Among them are fluorine, chlorine, sodium hydroxide, many oxidising agents, in particular highly pure hydrogen, and hydrogen peroxide.

Electrochemical Polarisation. Overvoltage

When an electrode is at a potential equal to the equilibrium one, electrochemical equilibrium sets in on it:

Ox + ze- = Red

When the potential of the electrode is displaced in the positive or negative direction, oxidation or reduction processes begin to proceed on it. The deviation of the potential of an electrode from its equilibrium value is called electrochemical polarisation or simply polarisation.

An electrode can be polarised by connecting it to a direct-current circuit. For this purpose, an electrolytic cell must be formed from an electrolyte and two electrodes—the one being studied and an auxiliary one. By connecting it to a direct-current circuit, we can make the electrode being studied the cathode or (with reverse connection of the cell) the anode. This method of polarisation is called polarisation from an external source of electrical energy.

Let us consider a simple example of polarisation. Assume that a copper electrode is in a 0.1m solution of CuS04 containing no impurities and no dissolved oxygen. As long as the circuit is not closed, the electrode potential at 25 °C will have an equilibrium value equal to

and electrochemical equilibrium will set in at the metal-solution interface:

Let us connect the electrode to the negative pole of the current source—we shall make it the cathode. The surplus of electrons that now appears at the electrode will shift the potential of the electrode in the negative direction and simultaneously violate equilibrium. The electrons will attract copper cations from the solution—a reduction process will occur:

If we connect the electrode to the positive pole of the current source instead of to the negative one, i.e. make it the anode, then owing to the withdrawal of part of the electrons, the potential of the electrode will shift in the positive direction and equilibrium will also be violated. But now an oxidation process will occur on the


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