As you already know, chemical processes can be accompanied by various phenomena - absorption and release of heat, light, sound, etc. In particular, they can lead to or be caused by an electric current. Such processes are called electrochemical, and their discovery played a significant role in both chemistry and physics.

Take two identical glasses. In one we pour a solution of copper chloride and lower a copper plate into it, in the other - a solution of zinc chloride and lower a zinc plate into it. Outwardly, nothing happens in both glasses. However, if we connect the metal plates with a conductor with a galvanometer and an ammeter built into it, then we will see that the galvanometer needle deviates, indicating the presence of a potential difference. In this case, the ammeter needle will remain at zero, which indicates the absence of current between the plates. What is going on?

Although we did not see anything when we lowered the copper plate into the copper salt solution, something did happen. In a very thin (practically monomolecular) solution layer adjacent to the metal, polar water molecules began to pull out its ions from the crystal lattice of copper:

Cu (tv) "Cu 2+ +2e -

This process can be considered as an ordinary chemical reaction, but with the participation of an unusual reagent - electrons, which remain in the metal as a result of the reaction, giving it a negative charge, the solution layer adjacent to the metal acquires a positive charge due to an excess of positive ions. A potential difference arises, which tends to return

copper ions back into the metal, and equilibrium is established. It turns out that in as a result of a chemical process, an electrical device appeared - a capacitor (though having molecular dimensions). It is called a double electric layer, and the entire created system (metal - a solution of its salt) - half element In contrast to the usual chemical equilibrium, the one obtained by us is characterized not only by the ratio of the concentrations of reactants and products, but also by the potential difference in the double electric layer. This difference is called electrode potential metal and characterizes the redox ability of a solid metal. (We note right away that such an ability for a gaseous metal is characterized by a completely different value - ionization potential, which is equal to the energy required to detach an electron from an isolated atom).

It is practically impossible to directly measure the electrode potential - after all, it exists between objects separated by one layer of molecules. However, if we take two half-elements formed by different metals (as in our experiment), then the potentials on the metal plates will be different, which we noticed. The resulting system of two half-cells is called a galvanic cell.

: If we connect the glasses in our experiment with a tube with a solution of some salt (salt bridge), then the ammeter will show the presence of current. At the same time, since the electrode potential of zinc is lower than that of copper, the electrons from the zinc plate will go to the copper one. According to Le Chatelier's principle, the equilibrium in the double electric layer will shift in both half-elements (after all, electrons participate in the reaction!) This will lead to the fact that copper from the solution will be deposited on the copper plate, and zinc will leave the zinc plate into the solution. Through the salt bridge, the excess of positive ions from the glass of zinc chloride will pass into the copper chloride solution, restoring the electrostatic balance. This process will continue until either the zinc is completely dissolved or the copper chloride runs out. If we abstract from electrical processes and consider only chemical ones, then we get the reaction: Cl 2 + Zn \u003d Cu + ZnCl 2

But it can be carried out without a galvanic cell! However, only his participation explains why the reaction goes exactly in this direction, and, say, not vice versa. Thus, knowledge of the values ​​of electrode potentials makes it possible to predict the possibility

ness and direction of redox reactions. How can you recognize them?

If you use the same half-element (reference electrode) in combination with various others, you can get a set of values ​​that will differ from the electrode potentials of the compared metals by the same amount - by the potential of the reference electrode. In practice, these quantities can be used in the same way as the electrode potentials themselves.

In fact, as a reference electrode is used hydrogen electrode. It is a specially prepared platinum plate immersed in a sulfuric acid solution with a hydrogen ion concentration of 1 mol/l and washed by a continuous stream of hydrogen under a pressure of 100,000 Pa at a temperature of 25°C. In this case, the following processes occur on the platinum surface.

H "H + + e - (2)

Reaction (2), as can be seen, is very similar to that which occurs in a metal half-element. A potential appears on the platinum plate, which is conditionally taken as zero.

If a plate of metal immersed in a solution of its salt with a concentration of 1 mol / l is connected to a galvanic cell with a hydrogen electrode at a temperature of 25 ° C, then the resulting potential difference is called the standard electrode potential of the metal and is denoted as E °.

Metals, arranged in ascending order of their standard electrode potentials, form the so-called electrochemical series of metal voltages

Li, Rb, K, Ba, Sr, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Cd, Co, Ni, Sn, Pb, H, Sb, Bi, Cu, Hg, Ag, Pd, Pt, Au

If we recall what happened in our galvanic cell, it is easy to understand why the arrangement of metals in this row predicts their properties:

1) Each metal can displace (restore) from solutions of their salts those metals that are in the series of voltages after it.

2) All metals that have a negative electrode potential (that is, standing in a series of voltages up to hydrogen) can displace (restore) it from acid solutions.

As you might guess, the concept of the standard electrode potential is applicable not only to the metal/metal ion system, but also to any reaction involving electrons. These reactions are well known to you: you wrote them, making up an electron-ion balance to equalize redox reactions, for example:

Cr 2 O 2- 7 + 14H + + be - ®2Cr 3+ + 7H 2 O

We will not dwell on how the standard electrode potentials of such half-reactions are measured - this is beyond the scope of this course, but such methods exist, and with their help the standard redox potentials of a huge number of reactions have been determined. They are summarized in tables, where standard reaction potentials are given in the form:

| oxidized form | + ne - ® | restored form |

and, accordingly, show the oxidative capacity of the oxidized form. In order to understand whether a redox reaction is possible, it is necessary to find the difference in the standard potentials of the corresponding half-reactions. For example, we will find out whether it is possible to obtain free halogens by oxidizing bromides and chlorides using an acidic dichromate solution. We find in table 12 the half-reaction for the oxidizing agent

In the case of bromide, the potential difference is 0.28 V > 0 and the reaction K 2 Cr 2 O 7 + KBr + H 2 SO 4 ®Cr 2 (SO 4) 3 + K 2 SO 4 + H 2 O + Br 2

will go. In the case of chloride, the difference is -0.01 V<0 и аналогичная реакция происходить не будет. Напротив, будет идти обратная реакция, то есть окисление трехвалентного хрома в кислом растворе хлором. Однако нужно помнить, что выяснять направление реакции с помощью стандартных потенциалов можно только при условии, что реакция проходит при 25°С, а Концентрации всех реагентов - 1 моль/л. Так, на самом деле реакция окисления хлорида калия бихроматом калия будет идти, так как при 25°С невозможно создать в растворе концентрацию хлора 1 моль/л.

Standard (normal) hydrogen electrode. Standard electrode potential. Tables of standard redox potentials

In electrochemistry, the standard electrode potential, denoted Eo, E0, or EO, ​​is a measure of the individual potential of a reversible electrode (in equilibrium) in the standard state, which occurs in solutions at an effective concentration of 1 mol/kg and in gases at a pressure of 1 atmosphere or 100 kPa (kilopascals). Volumes are most often taken at 25 °C. The basis for an electrochemical cell such as a galvanic cell is always a redox reaction, which can be broken down into two half-reactions: oxidation at the anode (losing an electron) and reduction at the cathode (gaining an electron). Electricity is generated due to the difference in the electrostatic potential of the two electrodes. This potential difference is created as a result of differences in the individual potentials of the two metals of the electrodes in relation to the electrolyte. Calculation of standard electrode potentials

The electrode potential cannot be obtained empirically. The potential of the galvanic cell results from the "pair" of electrodes. Thus, it is not possible to determine the value for each electrode in a pair using the empirically derived potential of a galvanic cell. For this, a standard hydrogen electrode is installed, for which this potential is precisely determined and equal to 0.00 V, and any electrode for which the electronic potential is not yet known can be correlated with a standard hydrogen electrode to form a galvanic cell - and in this case, the potential of the galvanic cell gives the potential of the unknown electrode.

Since electrode potentials are traditionally defined as reduction potentials, the sign of an oxidizing metal electrode must be reversed when calculating the total cell potential. You also need to keep in mind that the potentials do not depend on the number of electrons transferred in half-reactions (even if it is different), since they are calculated per 1 mole of electrons transferred. Hence, when calculating any electrode potential based on the other two, care should be taken.

For example:

(equation 1) Fe3+ + 3e? --> Fe(tv) -0.036 V

(equation 2) Fe2+ + 2e? --> Fe(tv) -0.44V

To get the third equation:

(equation 3) Fe3+ + e? --> Fe2+ (+0.77 V)

multiply the potential of the first equation by 3, reverse equation 2 (change the sign) and multiply its potential by 2. Adding these two potentials will give the standard potential of equation 3.

Table of standard electrode potentials

Main article: Table of standard electrode potentials

The higher the standard reduction potentials, the easier they can be reduced, in other words, the stronger oxidizing agents they are. And vice versa: a large negative potential means that this form is a strong reducing agent. For example, F2 has 2.87 V, and Li+ has -3.05 V, fluorine is an oxidizing agent, lithium is a reducing agent. Thus, Zn2+, whose standard reduction potential is -0.76V, can be oxidized by any other electrode whose standard potential is greater than -0.76V. (e.g. H+(0V), Cu2+(0.16V) , F2 (2.87 V)) and can be restored by any electrode whose standard potential is less than -0.76 V (for example, H? (-2.23 V), Na + (-2.71 V), Li + ( -3.05 V)). In a galvanic cell, where a spontaneous redox reaction causes the cell to produce an electrical potential, the Gibbs energy DGo must be negative, according to the following equation:

DGoyach = -nFEoyach

where n is the number of moles of electrons per mole of products and F is Faraday's constant, ~96485 C/mol. Therefore, the following rules apply:

if Eocell > 0, then the process is spontaneous (galvanic cell)

if Eocell< 0, тогда процесс несамопроизвольный (электролитическая ячейка)

Non-standard conditions

Standard electrode potentials are given under standard conditions. However, real cells can operate under non-standard conditions. Given a standard potential, the potential at non-standard effective concentrations can be calculated using the Nernst equation:

The E0 values ​​are temperature dependent (except for the standard hydrogen electrode) and generally refer to a standard hydrogen electrode at that temperature. For condensed phases, the magnitudes of the potentials also depend on pressure.

Potential. From the course of physics it is known that the electric potential is the work of moving a single positive charge from a given point in space to infinity. Each electrode has some electrical potential. The absolute value of the electrode potential cannot be determined. You can only compare the potentials of different electrodes with each other. To do this, two electrodes must be combined into an electrochemical circuit. To do this, the metal parts are connected by a conductor, and the electrolyte solutions in which they are immersed, by a glass tube filled with an electrolyte solution (usually potassium chloride). This tube is called an electrolytic key or salt bridge. It provides ionic conductivity between solutions. Thus, a closed circuit or a galvanic cell is formed, which is shown in fig. 3.

The difference in electric potentials of two electrodes in such a circuit is called the electromotive force of the EMF circuit (Fig. 4. Electrochemical circuit with a standard hydrogen electrode: - standard hydrogen electrode, 2 - electrode under study, 3 - electrolytic key). The EMF value can be measured, making it possible to compare the potentials of the electrodes with each other. Usually, a standard hydrogen electrode is used as an electrode, relative to which the potentials of all systems are determined. Its potential is conditionally taken equal to zero.

Thus, the electrode potential is called the EMF of an electrochemical circuit - a galvanic cell, composed of the electrode under study and a standard hydrogen electrode. Such a circuit is shown in Fig. 4. The electrode potential is usually denoted by the letter E.

The electrode against which the potential is measured is called the reference electrode. In addition to hydrogen, silver chloride, calomel and some others are used as reference electrodes. In all cases, the potential of the reference electrode is assumed to be zero. You can go from one scale of potentials to another. For example, the standard potential of a zinc electrode on a hydrogen scale is - 0.76 V, and the potential of a silver chloride electrode is + 0.22 V (on the same scale). Therefore, the potential of the zinc electrode on the scale of the silver chloride electrode will be: - 0.76 - 0.22 = 0.98 V. Measurement of electrode potentials.

It is rather difficult to accurately measure the electrode potential, since it is necessary that the balance on the electrodes is not disturbed during the measurement process. For this reason, it is impossible to obtain an exact value of E using a conventional voltmeter: if we close the circuit using a voltmeter instead of a conductor, then a rather large current will begin to flow in it, which will upset the balance on the electrodes. For measurement, you can use special voltmeters with high input resistance (more than 1012 ohms). When included in the circuit of such a device, the current flowing is too small to have a significant effect on the electrode balance.

The standard electrode potential is the potential of the electrode under standard conditions, it is denoted by the symbol E °. These potentials have been determined for many redox systems and are usually given in chemical reference books. If the electrodes (for example, metal electrodes of the 1st kind) are arranged in order of increasing potential, then we will get a table called a series of standard electrode potentials. This series is often referred to as the stress series, but this term is obsolete and should not be used.

With the help of a number of standard electrode potentials, it is possible to characterize some of the chemical properties of metals. For example, it is used to determine the sequence in which metal ions are reduced during electrolysis, as well as to describe other properties of metals.

The smaller the algebraic value of the potential, the higher the reducing ability of this metal and the lower the oxidizing ability of its ions. As follows from this series, lithium metal is the strongest reducing agent, while gold is the weakest. Conversely, the gold ion Au3+ is the strongest oxidizing agent, and the lithium ion Li+ is the weakest.

Each metal in the series of standard electrode potentials has the ability to displace all metals following it from solutions of their salts. However, this does not mean that repression necessarily occurs in all cases. For example, aluminum displaces copper from a solution of copper (II) chloride CuCl2, but practically does not displace it from a solution of copper (II) sulfate CuSO4. This is due to the fact that the chloride ion Cl- quickly destroys the protective surface film on aluminum, and the sulfate ion SO4 2- practically does not destroy it.

All metals having negative values ​​of standard electrode potentials, i.e. standing in a row up to hydrogen, displace hydrogen from dilute acids, the anions of which do not show oxidizing properties (for example, from HCl or dilute H2SO4) and dissolve in them. However, there are exceptions. For example, lead is practically insoluble in sulfuric acid. This is due to the formation of a protective film of sparingly soluble lead sulfate PbSO4 on the metal surface, which makes it difficult for the metal to contact with the acid solution. Therefore, we can conclude that a number of standard electrode potentials should be used, taking into account all the features of the processes under consideration.

Standard redox potentials. The possibility of any redox reaction occurring under real conditions is due to a number of reasons: temperature, the nature of the oxidizing agent and reducing agent, the acidity of the medium, the concentration of substances involved in the reaction, etc. It can be difficult to take into account all these factors, but, remembering that any redox reaction proceeds with the transfer of electrons from the reducing agent to the oxidizing agent, it is possible to establish a criterion for the possibility of such a reaction.

The quantitative characteristics of redox processes are the normal redox potentials of oxidizing agents and reducing agents (or standard electrode potentials).

To understand the physicochemical meaning of such potentials, it is necessary to analyze the so-called electrochemical processes.

Chemical processes accompanied by the appearance of an electric current or caused by it are called electrochemical.

To understand the nature of electrochemical processes, we turn to the consideration of several fairly simple situations. Imagine a metal plate immersed in water. Under the action of polar water molecules, metal ions are detached from the surface of the plate and hydrated, they pass into the liquid phase. In this case, the latter becomes positively charged, and an excess of electrons appears on the metal plate. The further the process proceeds, the greater the charge becomes, both of the plate and of the liquid phase.

Due to the electrostatic attraction of solution cations and excess metal electrons, a so-called electric double layer appears at the phase boundary, which inhibits the further transition of metal ions into the liquid phase. Finally, there comes a moment when an equilibrium is established between the solution and the metal plate, which can be expressed by the equation:

or taking into account the hydration of ions in solution:

The state of this equilibrium depends on the nature of the metal, the concentration of its ions in solution, on temperature and pressure.

When a metal is immersed not in water, but in a solution of a salt of this metal, the equilibrium shifts to the left in accordance with Le Chatelier's principle and the more, the higher the concentration of metal ions in the solution. Active metals, whose ions have a good ability to go into solution, will in this case be negatively charged, although to a lesser extent than in pure water.

The equilibrium can be shifted to the right if electrons are removed from the metal in one way or another. This will dissolve the metal plate. On the contrary, if electrons are brought to a metal plate from the outside, then ions will precipitate from the solution on it.

When a metal is immersed in a solution, a double electric layer is formed at the phase boundary. The potential difference that occurs between the metal and the surrounding liquid phase is called the electrode potential. This potential is a characteristic of the redox ability of the metal in the form of a solid phase.

In an isolated metal atom (the state of a monatomic vapor that occurs at high temperatures and high degrees of rarefaction), the redox properties are characterized by a different quantity called the ionization potential. The ionization potential is the energy required to detach an electron from an isolated atom.

The absolute value of the electrode potential cannot be measured directly. At the same time, it is not difficult to measure the electrode potential difference that occurs in a system consisting of two metal-solution pairs. Such pairs are called half elements. We agreed to determine the electrode potentials of metals with respect to the so-called standard hydrogen electrode, the potential of which is arbitrarily taken as zero. A standard hydrogen electrode consists of a specially prepared platinum plate immersed in an acid solution with a hydrogen ion concentration of 1 mol/l and washed by a hydrogen gas jet at a pressure of 105 Pa at a temperature of 25 °C.

A number of standard electrode potentials. If a metal plate immersed in a solution of its salt with a concentration of metal ions equal to 1 mol / l is connected to a standard hydrogen electrode, then a galvanic cell will be obtained. The electromotive force of this element (EMF), measured at 25 ° C, characterizes the standard electrode potential of the metal, usually denoted as E °.

The standard potentials of electrodes that act as reducing agents with respect to hydrogen have the “-” sign, and the “+” sign has the standard potentials of electrodes that are oxidizing agents.

Metals, arranged in ascending order of their standard electrode potentials, form the so-called electrochemical series of metal voltages: Li, Rb, K, Ba, Sr, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Cd, Co, Ni , Sn, Pb, H, Sb, Bi, Cu, Hg, Ag, Pd, Pt, Au.

A number of stresses characterize the chemical properties of metals:

1. The more negative the electrode potential of the metal, the greater its reducing ability.

2. Each metal is able to displace (restore) from salt solutions those metals that are in the electrochemical series of metal voltages after it.

3. All metals that have a negative standard electrode potential, that is, those that are in the electrochemical series of metal voltages to the left of hydrogen, are able to displace it from acid solutions.

As in the case of determining the E° value of metals, the E° values ​​of non-metals are measured at a temperature of 25 ° C and at a concentration of all atomic and molecular particles involved in the equilibrium equal to 1 mol/l.

The algebraic value of the standard redox potential characterizes the oxidative activity of the corresponding oxidized form. Therefore, the comparison of values

standard redox potentials allows you to answer the question: does this or that redox reaction take place?

A quantitative criterion for assessing the possibility of a particular redox reaction is the positive value of the difference between the standard redox potentials of the oxidation and reduction half-reactions.

HYDROGEN electrode in electrochemistry - usually a platinum plate immersed in an acid solution with a certain concentration of H + ions and washed by hydrogen gas. At a hydrogen pressure of 0.1 MPa and the thermodynamic activity of its ions equal to unity, the potential of the hydrogen electrode is conditionally taken to be zero. Such a hydrogen electrode is called a standard electrode; it serves as a reference electrode, from which the potentials of other electrodes are measured.

32 Thermodynamics of electrode processes. Spontaneous occurrence of redox reactions. Relationship between the EMF of a galvanic cell and the Gibbs energy. Connection of EMF with the equilibrium constant

Any chemical reactions are associated with the movement of electrons, so they can be used to generate an electric current. In this case, the source of electrical energy is the energy released during a chemical reaction. Such a transformation of the energy of a chemical reaction into electrical energy is possible only with the help of a special device called a galvanic cell. It allows you to direct the flow of electrons through metal conductors.

Simple combustion of hydrogen is accompanied by a large release of heat. If it is carried out at a constant volume, for example, in a calorimetric bomb, then DU = -284.5 kJ / mol. If the same reaction is carried out in a galvanic cell by an electrochemical method, then part of this loss of internal energy can be used to generate an electric current. A diagram of such a galvanic cell is shown in Fig. IX.1. Two platinum electrodes are immersed in an aqueous solution (for example, NaOH). The left electrode is bathed in hydrogen bubbles, and the right electrode is bathed in oxygen. The hydrogen on the left side of this cell is dissolved in platinum and ionized. Due to the high affinity for water molecules, a certain amount of protons passes into the solution layer immediately adjacent to the electrode. In this case, hydroxonium ions H3O+ are formed - they are indicated by pluses in the right part of Fig. IX. 1, and the electrons (minuses) remain on the surface of the platinum electrode. Due to the electrostatic attraction between electrons and hydronium ions, the latter remain near the electrode and do not go into the bulk of the solution. Due to this, a so-called double electric layer arises at the metal-solution interface, similar to two plates of a capacitor. On the surface of the right electrode, the reaction of the formation of hydroxyl ions occurs:

3/2O2g + H2Ol + 2e = 2OH-

which removes two electrons from the metal. The surface of the metal is therefore positively charged and a double electric layer is also formed on it, but of the opposite sign. If you connect the left and right electrodes with a metal conductor, then an electric current will flow through it. The arrow in fig. IX.1 indicates the direction of electron flow. The difference in electrical potentials at the electrodes of an open galvanic cell is called its electromotive force (emf).

Obviously, the flow of electrons that occurs in the element can be used to produce work, for example, to rotate an electric motor. The flow of current leads to a decrease in the charges of double electric layers. Therefore, H3O+ and OH- ions get the opportunity to move away from the electrodes and form neutral water molecules in the solution. At the same time, due to reactions on the electrodes, double layers are restored again. The changes occurring on the electrodes and in the solution are reflected by the following equations:

H2g = 2H+ + 2e;

3/2 O2g + H2Ol + 2e = 2OH-;

2H+ + 2OH- = 2H2Ol,

the sum of which is the reaction of water formation:

H2g + 1/2O2g = H2Ozh,

Thus, the same reaction of formation of water from elements can be carried out in two different ways. Which of these methods is more profitable in terms of converting the energy of a chemical reaction into work? In the first method, when hydrogen is burned in a calorimetric bomb (V = const) at 298 K, the decrease in internal energy is equal to the amount of heat released -DU = 284.5 kJ / mol, and the work is zero.

In the second case, part of this change in internal energy (DG) can be converted into electrical work. If the reaction in the galvanic cell is carried out reversibly, then the accompanying loss of Gibbs energy goes entirely to the production of electrical work.

In the case under consideration, DG0 = -237.2 kJ/mol, and, consequently, only ?47 kJ/mol is converted into heat. This example shows that, in general, the energy released during the combustion of natural fuels is more advantageous to directly convert into electrical energy, since the efficiency of heat engines and thermal power plants is low. The described hydrogen-oxygen cell is an example of so-called fuel cells.

Work on the creation of such elements has recently been widely developed in connection with new technical problems. In these cells, the fuel and oxidizer must be stored separately and fed to the electrodes on which the electrochemical reactions take place. In this case, the element can operate continuously if reagents are supplied to it and reaction products are removed, which is especially convenient when using liquid and gaseous substances. Instead of burning coal, it is possible to use the reaction St + O2g = CO2g to produce electric current.

Obviously, in real conditions, galvanic cells work irreversibly, so only a part of the change in the Gibbs energy of the reaction occurring in the cell is converted into work. We repeat that a galvanic cell can operate under the condition that a spontaneous chemical reaction or some other spontaneous process occurs in it, accompanied by a decrease in the Gibbs energy.

If a sufficiently large potential difference is applied to the considered galvanic cell from the outside, exceeding its e. d.s. and having the opposite direction, then the decomposition of water will occur with the release of hydrogen and oxygen. Thus, the processes of obtaining electric current in galvanic cells and electrolysis are mutually opposite.

A feature of the electrochemical process in a galvanic cell is the theoretically important possibility of its implementation under conditions very close to reversibility. This is achieved thanks to the potentiometric method, in which e. d.s. of the studied galvanic cell is almost completely compensated with the help of an oppositely directed em. with. external source. This technique allows you to measure emf. in the absence of current in the circuit, i.e. when the element does not work, and its emf. maximum. Control over the absence of current is carried out by galvanometers (null-tools) of high sensitivity. They give a deviation in the passage of current with a force of 10-8 - 10-9 A. Such a weak current, when passing through the electrolyte, even for many years, could not release any noticeable amounts of substance.

Rice. IX.2. E.m.f. measurement scheme compensation method.

Schematic diagram of the measurement e. d.s. galvanic cell by the compensation method is shown in fig. IX.2. Direct current from the auxiliary battery WB is supplied to the ends of the reochord AB - a wire with a constant cross section. Therefore, the voltage drop along the reochord is proportional to the length of the corresponding segment on the straight line AB. With the help of moving contact C, it is possible to select an arbitrary part of the voltage drop between points A and B. From fig. IX.2 it can be seen that the stress removed from any section of the reochord, for example AC, is directed towards e. d.s. element x.

By moving contact C along the rheochord, a position is found at which the zero-galvanometer G indicates the absence of current in the AHGS circuit. This means that the potential drop from the WB on the AC segment completely compensates for e. d.s. element x.

If e. d.s. auxiliary battery WB is equal to EB, then e. d.s. element X EX is determined from the proportion:

EX/EB = AC/AB, whence EX = (AC/AB) EB.

In order to calibrate the auxiliary battery before the EX measurements, instead of the X element, another one is switched on, e. d.s. which is exactly known, for example, the standard Weston element. The device of this element will be described below.

Let us repeat that e. d.s. maximum, since during the measurement there is no potential drop both outside and inside the element. The work done by an element with a negligible current during a reversible process would be maximum.

Of theoretical and practical interest are galvanic cells with metal electrodes. Consider, for example, the reaction Znt + CuSO4aq. rr. = ZnSO4 aq. solution + Cut or Znt + Cu2+ = Zn+2 + +Cut, which can be done in two ways. One of them is completely irreversible. A zinc plate is placed in an aqueous solution of copper sulphate, and metallic copper is released and zinc dissolves. The electrons go from zinc directly to copper, and the reaction proceeds without work being done, but is accompanied only by the release of heat. In the case of a hydrogen-oxygen element, conditions can be created in which electrons will move along a metal conductor and do work. This is achieved in a galvanic cell, where the zinc electrode is immersed in a ZnSO4 solution, and the copper electrode in a CuSO4 solution.

The solutions are separated from each other by a porous (ceramic) partition, which prevents their mixing, but ensures the passage of an electric current due to the diffusion of ions through the pores. Such an element, on the electrodes of which double electrical layers are formed, was designed by the Russian electrochemist B.S. Jacobi.

The magnitude and sign of electric charges in double layers are determined by the work of removing an electron from the metal and the hydration energy of its ions. Those metals will easily pass into the solution, which have a lower work function of electrons and a higher energy of hydration of ions, i.e. less noble metals. Since zinc is less noble than copper, it will be charged more negatively than copper. If you connect both electrodes with a metal conductor, then the electrons will move from zinc to copper. As a result, Zn2+ zinc ions are not retained in the double layer by the attraction of electrons, they pass into the bulk of the solution, and the electrons that have passed to the copper electrode discharge the Cu2+ ions, transferring them to the metallic state.

Therefore, during the operation of the element, the zinc electrode is dissolved and copper is deposited on the copper electrode. For the element to work, the circuit must be closed, i.e. there must be electrical contact between the solutions. The current transfer inside the element is carried out by ions. In the element, the transition of electrons from zinc to copper occurs not under conditions of direct contact of these metals, but with the help of a conductor. The overall reaction in the cell consists of two spatially separated electrode processes.

Reactions occurring in galvanic cells are redox. In the case under consideration, zinc is oxidized, which loses electrons, and copper is reduced, gaining electrons. In general, any redox reaction can be used to generate electric current using a galvanic cell. As mentioned, such a reaction can be the combustion of any type of fuel.

When schematically recording galvanic cells, the boundaries between phases are marked with vertical lines. Provided that there is no potential difference at the boundary of two liquids (in this case, solutions of ZnSO4 and CuSO4), it is indicated by two vertical lines. The scheme of the considered element has the following form:

Zn? ZnSO4? CuSO4? Cu.

It is customary to write such circuits in such a way that the left electrode is negative (electrons flow along a metal conductor from left to right and positive electricity inside the element is carried by ions in the same direction). Such a record corresponds to the course of the reaction, accompanied by a decrease in the Gibbs energy and a positive value of e. d.s.

Galvanic cells can be built not only with the use of aqueous solutions of electrolytes, but also with the use of melts. An example of such an element is the chain Ag? AgBr? Br2, in which the left electrode is silver, and the right one is graphite bathed in gaseous bromine, and the electrolyte is molten AgBr. Silver dissolves on the left electrode: Agt > Ag+ + e, and on the right electrode, bromine adsorbed by graphite: 1/2Br2g + e = Br-. Thus, the reaction occurs in the element: Agt + 1/2Br2g = AgBrzh.

Recently, galvanic cells with solid electrolytes having oxygen conductivity have gained great importance (see Chapter VIII), for example,

The left electrode is a mixture of iron and its oxide. Here, the oxidation reaction of iron occurs with O2- ions coming through the solid electrolyte. In this case, electrons are released, and the electrode receives a negative charge. On the right electrode, consisting of a mixture of Mo and MoO3, the oxide is reduced. This is accompanied by the absorption of electrons in such a way that the electrode becomes positively charged and the released O2 ions can migrate through the electrolyte to the left electrode. The reaction at the electrode is represented by the following equation 3Fet + 3O2- = 3FeOt + 6e; on the right electrode: MoO3m + 6e = Mot + 3O2-.

Note that the sum of these two reactions 3Fet + MoOt = 3FeOt + MoOt is the process of reducing molybdenum oxide with iron, the spontaneous occurrence of which is the source of electrical energy produced by the element.

It can be seen from the considered examples that the reaction taking place in a galvanic cell can be represented as two separate electrode reactions.

It can be assumed that e. d.s. galvanic cell should depend on the nature of the reactants, their concentrations and temperature. To find expressions for these dependencies, it is necessary to consider the thermodynamic relationships that characterize the operation of a galvanic cell.

Let the following reaction take place in a galvanic cell: M + Nn+ = Mn+. The work done by an element at a consumption of 1 mole M is determined by the product of the amount of electricity nF by the value e. d.s. E, i.e. W = nFE, where n is the number of moles of electrons flowing through the circuit; F - Faraday number, equal to 96493 C. For example, for the reaction Zn + Cu2+ = Zn2+ + Cu, n = 2. If the element works reversibly at constant pressure and temperature, then the work done by it is equal to the loss of the Gibbs energy, i.e. DG = W:

DG = -nFE = -96493E. (IX.1)

If the element works irreversibly, then nFE< -ДG, т.е. э.д.с. меньше, чем при обратимом проведении реакции. Выражая E в В, получаем величину ДG в Дж.

Thus, if the stoichiometric equation of the reaction occurring in the galvanic cell and the tabular data on the change in the Gibbs energy are known, e can be calculated. d.s.

So, for the hydrogen-oxygen element considered above, which operates due to the energy released during the reaction H2g + 1/2O2g = H2Ol, for which DG 0

298 = -237200 J, n = 2, рH2 = рO2 = 1.

/n 96493 = -(-237200/2) 96493 ?? 1.2 V.

Equation IX.1 implies that the measurement e. d.s. galvanic cell allows you to find the change in the Gibbs energy of the reaction occurring in it. Therefore, the method e. d.s. widely used to determine the thermodynamic properties of substances.

In the example above, this method makes it possible to find the DG of the reduction reaction of MoO3 with iron. Knowing the standard change in the Gibbs energy during the formation of FeO(DG 0 f FeO) from the found value of DG, we can find the Gibbs energy of MoO3 formation from the equation:

e. d.s. from temperature. Since the Gibbs energy is a function of temperature, then e. d.s. galvanic cell should also depend on temperature.

To find this dependence, we use the Gibbs-Helmholtz equation: ДG = ДH + T(?ДG/?T)p substituting the expression ДG into it through e. d.s. In this case, we get -nEF = ДH - TnF(dE/dT) or

DH = nF, (IX.2)

DH = W - TnF(dE/dT). (IX.3)

Let us first imagine that the galvanic cell placed in the calorimeter is short-circuited. In this case, the electrical energy produced by it will completely turn into heat, the amount of which is equal to the enthalpy of the reaction DH, and, therefore, the work will be equal to zero.

Now let the reaction in the element be carried out reversibly, for example, the wires from the electrodes are taken out of the calorimeter, connected to the motor, and the electric current does work. Then part of the energy released during the reaction will turn into electrical work W, and the other part Q will remain in the form of heat and will be measured in the calorimeter. According to the first law of thermodynamics

DH = W - Q (IX.4)

Comparison of equations (IX.3) and (IX.4) shows that

Q = TnF(dE/dT). (IX.5)

Obviously, the closer the course of reactions in a galvanic cell to the conditions of reversibility, the greater part of DG is converted into work. The quantity Q, which characterizes the bound energy, determines the amount of heat inevitably released (or absorbed) when the element works reversibly. Since (?DG /?T)p \u003d -DS and (?DG /?T)p \u003d -pF (dE / dT), then

DS = nF(dE/dT), (IX.6)

and, consequently, measurements of the temperature dependence of e. d.s. allow you to calculate the change in entropy during a reaction occurring in a galvanic cell. It should be emphasized that a galvanic cell can work both with the release and absorption of heat. In the latter case, it converts the heat of the environment into work. This is not in conflict with the second law of thermodynamics, since the processes in galvanic cells are not continuous and stop when the electrode material is used up.

The sign and magnitude of Q determine the temperature dependence of e. d.s. If heat is generated during the operation of the element, i.e. Q< 0, то температурный коэффициент э. д. с. dE/dT < 0. Это наиболее часто встречающийся случай, так как большинство элементов работает с выделением тепла. Наоборот, при Q >0 e. d.s. grows with temperature.

For galvanic cells serving as standards, during electrical measurements, such reactions are selected in which Q is very small and dE / dT is close to zero. So, dependence d.s. on the temperature of the commonly used standard Weston element is expressed by the equation:

E = 1.0183 - 0.0000406 (t - 20) V.

It is composed according to the scheme: Cd ? CdSO4? ? Hg2SO4? Hg, and the reaction Cdt + 2Hg+ = Cd2+ + 2Hgl1 takes place in it.

As an example of the application of equations (IX.4) and (IX.5), we calculate the value of dE / dT for an element in which the reaction Znt + 2AgCl = ZnCl2 + 2Agt takes place

DH = 217760 J and E = 1.015 V at 0°C. Hence

Q \u003d -DH \u003d 217760 - 2 96493 1.015 \u003d 21880 J.

dE/dT = -218807(273 2 96493) ?? - 4 10-4 W/C.

An example of an element with a positive temperature coefficient is the Hg ? Hg2Cl2, KCl? KOH? Hg2O? Hg, in which the reaction Hg2Cl2 + 2KOH = 2KCl + Hg2O + H2O takes place.

The left electrode of this element, called the calomel electrode, is often used in electrochemical measurements. It consists of liquid mercury in contact with solid calomel Hg2Cl2 and an aqueous solution of some strong electrolyte, such as KS1. The reaction taking place in the element under consideration is endothermic, DH = 13720 J, and W = 31570 J. Thus, Q = 13720 + 31570 = 45240 J, i.e. the element absorbs heat equal to 45240 J from the environment. Part of this heat, equal to 31570 J, goes to the production of work.

e. d.s. on the concentrations of electrolytes involved in the reaction can be: found using the equation of the isotherm of a chemical reaction.

Let the reaction A + B = 2D proceed in a galvanic cell, while DG = RTlnK + RTln (c 2 D/cAcB). Substituting the value - nEF instead of ДG and dividing both parts of the equation by -пF, we get E = RTln(K/nF) - . or, denoting the value of RTlnK/nF, which depends only on temperature, through E0, we will have:

E = E0 - (RT/nF. (IX.7.)

The value of E0 is called the standard e. d.s. element. It characterizes an element in which the concentrations of all substances participating in the reaction are equal to unity, and the change in the Gibbs energy is equal to the standard DG0. Replacing the natural logarithm with the decimal logarithm in equation (IX.7), we obtain for a temperature of 25 °C.

Obviously, for electrolytes, one cannot simply use the analytical concentrations of the corresponding substances, but it is necessary to take into account the dissociation and interaction of ions. In this regard, the problem arises of determining the activity of electrolytes.

9.1. REDOX POTENTIAL (ORP): THERMODYNAMIC MEANING, PROPERTIES. STANDARD AND FORMAL ORP

9.1.1. thermodynamic meaning. Standard ORP Table

ORP is defined as the electrical work of electron transfer in the course of redox interaction. At equilibrium, it is equal to the maximum chemical work performed during this interaction, which, in turn, is equal to the change in the Gibbs energy D

ORP can be measured in various ways - potentiometric, colorimetric, voltammetric, polarographic. The most common potentiometric method, which will be described below, in accordance with the focus of this book. Within the framework of this method, we obtain an expression for the quantitative assessment of the ORP.

To do this, it is necessary to construct a galvanic cell, on the electrodes of which an equilibrium is established between the oxenvironment of half-reactions with the participation of substances whose ability to be oxidized and reduced we want to characterize. We proceed in the same way as in Chap. 3 when considering the interaction oxmedia of two systems, 1 and 2 according to reaction (3.2),

passing on the electrodes of a galvanic cell (moreover, in Ch. 3 Ox, = A, Red, = B, Red, = L, Ox, = M).

(Galvani potentials at phase boundaries are shown).

EMF of the cell in which the reaction (9.1) takes place:

Let system (2) be a standard hydrogen electrode (SHE). Then cell (9-1) will be written

i.e. Oh 2 \u003d H +, Rcd 2 \u003d H 2.

The reaction taking place in the cell is as follows:

At R, 7'= const the change in Gibbs energy caused by the reversible passage of this reaction is

because "n+ and R n= 1. For one run of the reaction, where nF Cl electricity, according to (3.1) e.m.f. cells (9-11)

where

The same formula was obtained in Chap. 3 for the potential of the oxend electrode of the zeroth kind:

This is the quantitative measure of what we call redox potential or ORP (oxidizing reduction potential).

Before continuing the presentation of the material, it is necessary to say a few words about the term itself. Most of the old textbooks used the term "oxidation-reduction potential" or redox, or oxred potential, and until 1953 there were 2 scales of such potentials - the American one, in which the stronger restorative capacity The red shape of the system, the more positive the potential, and the stronger oxidation capacity Ox forms, the more negative it is (Li / Li + +3 V; 2СГ / СЬ -1.36 V, etc., see Latimer's well-known monograph-reference book). In the European system - on the contrary. Since 1953, the European system has been adopted everywhere.

School of B.P. Nikolsky promoted the term "oxidizing potential" instead of "redox", on the grounds that, under the European system of signs, the higher oxidation capacity the system in solution, the higher (positive) the potential (See textbook BPN, books NPYa and ShPP).

It can be seen from the IUPAC recommendations that the reference system H "/H 2 was introduced into the determination of the characteristics of the redox ability of the system, and its ability restore another system. Therefore, the potential should rather be called "restorative" which is done in the well-known textbook of general and inorganic chemistry by A. B. Nikolsky and A. V. Suvorov, according to which students of St. Petersburg State University study. At the same time, the term "redox" continues to be widely used, because, according to many authors, which we share, this term reflects both sides of the interaction.

For electrodes of other kinds, the action of which is ultimately reduced to an oxidative reaction, regardless of whether the Ox and Red forms are in the same or different phases, formula (9.6) is even simplified. If in such electrodes the Red- or Ox- form is one-component solid, liquid or gaseous phases, the activity of ions in them is taken equal to 1, and the formula for ORP takes the form (9.6 a, b).

Table and on 9.1

Standard electrode potentials for some redox half-reactions

in an aqueous medium at 25 °C and a pressure of 1 atm

(http://en.wikipedia.org/wiki/Standard_electrode_potential_(data_page))

	Half reaction			Half reaction
	Sr+ e~ - Sr			Cu 2 0 + H 2 0 + 2e= 2Cu + 20H
	3N2+ 2e~+ 2H + \u003d 2HN 3			Ag+ e= Ag + I
	La (OH) 3 (s) + Ze \u003d La (s) + ZONE
	A1 (OH) 3 + Ze \u003d A1 + ZON			H" + e= 1/2H,
	A1F 6 3 - + Ze^= A1 + 6F"			T1 2 0 3 + 3H.0 + 4e-= 2TG + 60 H
				AgBr(s) + e = Ag(s) + Br
	Zr0 2 (s) + 41-G + 4e=Zr(s) + 2H,0			AgCl+ e~= Ag + Cl
	Zn0 2 "+ 2Н,0 + 2nd= Zn + 40N			3- + e- = 4 "
	Zn(OH) 3 "+ 2e= Zn(s) + 4OH			0 2 + 2H,0 + 4e= 40H
	Fe(C 5 H 5) 2 + e" = Fe(C 5 H 5) 2			Cu + + e~= Cu
	2H20+ 2e~\u003d H 2 + 20 H "			I 3 - + 2e~= ZG
	Cr 3+ + e = Cr +			l 2 (s) + 2e~= 21
	Eu 3+ + e= EU 2+			PtCl 2 '+ 2e-= Pt+4CI

Table 9.1 (continuation)

	Half reaction			Half reaction
	Fe' + + e~ = Fe2+			HC10 2 (aq) + 2H " + 2e \u003d HClO (aq) + H 2 0
	AgF+ e~= Ag + F~			MNO; + 4H + + 3e" = MnO,(s) + 2H_,0
	MnO“ + H + +" = HMnO"			Ce4+ +e~= Ce 3+
	Mn0 2 (s) + 4H + + e = Mn3+ + 2H.0			PbO, + SO 2 "+4H + + 2e= PbS0 4 + 2H.0
	Cu 2+ + 2CN" + e =			SOUTH + 2e+ 6H + = Bi 2+ + 3H.0
	I0 3 - + 5 H+ + 4e~= HIO(aq) + 2H,0			H 2 0 2 (aq) + 2 H + + 2e= 2 H 2 0
	ClOj + 2H‘+e“ = C10,(g) + H 2 0			Co 3+ + 2e" = Co +
	0,+4H++ 4e~= 2H30
	MnO,(s) + 4 H + + 2 e = Mn 2+ + 2H,0			S.0 2 "+ 2e= 2S02"
	Tl 3+ +2e= Tl +			0,(g) + 2H* + 2e= 0 2 (g) + H 2 0
	Pb0 2 (s)+ 4H + 2e= Pb 2 - + 2H,0			HMn0 4 + 3H + + 2e-= Mn0 2 (s) + 2H.0
	Mn0 4 + 8H + + 5e~= Mn 2+ + 4H 2 0			F2+2H+ +2e~= 2HF
	HO", + H + + e H,0,(aq)			XeF+e=Xe+F"
	2HC10(aq) + 2H + + 2e = Cl 2 (g) + 2H.0

At a Ok = 1 AND % ed = 1 ^Ox/Red = ^Ox/Red

ORP. The standard ORP values ​​are given in Table 9.1, which we will refer to again and again. A similar table has already appeared in Chap. 3, Tab. 3.2. The position of the system in the table characterizes its redox ability. Half-reactions in Table. 9.1 are written according to the principle Ox + Red. Positive values?ox/Red mean that this reaction (reduction) proceeds spontaneously from left to right under standard conditions, negative ones mean vice versa. the stronger the Ox-form as an oxidizing agent.

Tab. 9.1 contains mainly inorganic OM systems in which there is a change in the degree of oxidation of certain elements that make up the oxidizing agent or reducing agent. These systems can be variously classified: homogeneous and heterogeneous liquid/gas or liquid/solid type, containing and not containing H,0 and non-complex ion ions in one or both forms; oxyanions in one or both forms, etc. Under the condition that electrode reactions are reversible, homogeneous systems of one kind or another can form electrodes of the zeroth kind (for example, Fe 3+ /Fe 2+. CN) ^ +; heterogeneous - electrodes of the 1st, 2nd and 3rd kind (for example, Me + / Me (s); SG, AgCl (s) / Ag (s); Ca 2+, CaC 2 0 4 (s), PbC 2 0 4 /Pb).The electrode potential of the last three systems obeys a formula like (9.6a), since Red = 1. But all components of the system contribute to the standard ORP:

And the last. To emphasize the linkage of the ORP to the r.e. scale, the notation is often used in the literature for it Eh or E n. In what follows, we will denote

The standard EMF of some pairs of half-elements can be calculated, without resorting to potentiometric measurements, through the defining EMF equation (9.12) using the Gibbs energies of the formation of reaction participants in the cell, if they are known:

In addition, there is a way of calculating which in many cases is simpler, more direct, and sometimes more accurate. For this, the standard electrode potentials of reduction reactions in an aqueous medium, published in tables of physicochemical quantities, serve.

The standard electrode potential of the reduction reaction is the standard EMF of an element composed of a given electrode and a hydrogen electrode, and the half-reaction on the hydrogen electrode is considered as the oxidation of hydrogen. That is, in the corresponding cell diagram, the hydrogen electrode is on the left anyway, so that the standard electrode potential refers to the hydrogen reduction reaction. Is he marked? e, like the standard EMF. It should not be understood as the electric potential of a terminal, electrode, or any other part in the construction of an element, although the term is often used as if it were.

For example, the standard emf of a Harned cell

considered in the previous sections is the standard electrode potential of the reaction:

In the tables, its value is indicated for the half-reaction AgCl (t) + + e " = A? (t) + SG (a), which should be understood as a conditional record of the complete reduction reaction of silver (+1) with hydrogen H 2.

Like any standard thermodynamic function, the standard electrode potential depends only on temperature and the choice of standard states.

The standard electrode potential of the hydrogen electrode is the standard EMF of the RDT element)|H 2 (g)|N + (th)|H 2 (g)|RDt). It is zero at any temperature.

Since the values ​​of the standard EMF are related to the standard Gibbs energy of the reaction by equation (9.20), they have an additivity property similar to this property for the values ​​of DS e. This can be seen by example. Let's talk about a galvanic cell

The total reaction of this element is:

The standard electrode potential of the left half-cell in (9.21) is equal to the standard EMF of the element

with reaction

The standard electrode potential of the right half-cell in (9.21) is equal to the standard EMF of the element

with reaction

It can be seen that reaction (1) is the difference between reactions (3) and (2). Therefore, in accordance with Hess's law, it is true

Therefore:

In reactions (1), (2), and (3), the stoichiometric numbers of electrons y e (y 1? y 2 and y 3) are equal to 2. Therefore, they cancel out, as does the Faraday constant. Then it turns out: = ?^ - ?This ratio is valid for any element. It is a consequence of Hess's law and can serve as a general rule according to which the standard EMF of any electrochemical element is equal to the difference between the standard electrode potentials of half-reactions occurring on the right and left electrodes.

Using this relation, one can calculate the standard emf of any element from the standard electrode potentials of the corresponding half-reactions, if they are known. To do this, it is not at all necessary to imagine this electrode paired with hydrogen. It is easier to follow another rule: both half-reactions of the element should be written (or mentally represented) as reduction half-reactions with electrons on the left side, find these half-reactions in the table of standard electrode potentials and calculate using (9.22). For example, according to this recipe for the element (9.21), two half-reactions have the form:

In the table of standard electrode potentials, you can find for them the values ​​\u200b\u200bof -0.403 and 0.222 V, respectively. Then according to the formula (9.22) it turns out:

It should be noted that standard EMF and standard electrode potentials do not depend on the nature of ions that do not directly participate in electrode reactions. This follows from the fact that the standard state of ions of a given sort in solution is a hypothetical solution with the properties of an ideally dilute solution. At ideal dilution, the properties of a given kind of ion are independent of the other ions present. Therefore, the derivation of equation (9.22), given above, will not change if, instead of element (9.21), we consider an element with transfer:

with any anions in the solution of the left half-cell and with any cations in the solution of the right half-cell. In the same way, the standard electrode potentials of the reactions in the tables do not depend on which ions of the opposite sign are conjugated with the ions indicated in these reactions.

electrode in electrochemistry called the interface between an electric current conductor with electronic conductivity and an electric current conductor with ionic conductivity, or, in other words , the place where the electronic mechanism of electric charge transfer changes to ionic (and vice versa). In a narrower sense, an electrode is often called a conductor of electric current with electronic conductivity.

Rice. 7.1.Schematic representation of a galvanic cell

Let's carry out the interaction reaction of Sn 2+ and Fe 3+ so that the processes of oxidation and reduction are spatially separated (Fig. 7.1). In a vessel containing Sn 2+ and Sn 4+ , ​​the following processes will take place. The Sn 2+ ions will donate electrons to the platinum wire and turn into Sn 4+ . In parallel, the reverse process will also take place. After some time, equilibrium will be established in the system:

Sn 4+ + Sn 2+

Rice. 7.2.Occurrence of electrode potential

Due to the establishment of this equilibrium, the surface of the platinum wire and the solution near it will have a different charge, the formation of the so-called "double electric layer" will occur (Fig. 7.2). At the interface "metal - solution" there will be a potential difference called electrode potential.

Similar processes will also occur in a system containing Fe 2+ and Fe 3+ . However, since Fe 2+ ions have a lower ability to donate electrons than Sn 2+, and Fe 3+ ions, respectively, a greater ability to accept electrons than Sn 4+ , ​​the surface of a platinum wire dipped into a solution containing Fe 2+ and Fe 3+ will be less negatively charged than the Sn 2+ and Sn 4+ dipped into the solution.

We connect the platinum plates dipped into the solutions with a metal conductor. To close the circuit, we connect both solutions with a salt bridge - a tube containing a KCl solution. In the resulting system, called galvanic element, electric current will begin to flow. If you include a potentiometer or a high-resistance voltmeter in this circuit, then you can measure its EMF, which will characterize the ability of Fe 3+ ions to receive electrons from Sn 2+.

The absolute value of the electrode potential of an individual electrode cannot be determined. It is possible to determine only the potential difference of two electrodes. In principle, this can be done for each specific reaction. However, it is much more convenient to choose one standard electrode, relative to which all measurements of electrode potentials will then be carried out. A standard hydrogen electrode is used as such a reference electrode.

Rice. 7.3 Standard hydrogen electrode

The standard hydrogen electrode is a platinum plate saturated with hydrogen, which is in a solution of H 2 SO 4 or HCl (Fig. 7.3). To increase the adsorption capacity, platinum is covered with a layer of spongy platinum. To saturate the platinum surface with hydrogen, gaseous H 2 is passed through the solution (p = 1 atm). An equilibrium is established between hydrogen dissolved in platinum and hydrated hydrogen cations in solution:

2H + +  H 2 (Pt)

The potential of a standard hydrogen electrode is assumed to be zero at any temperature.

Standard half-reaction electrode potential(E 0 , 0) - this is the EMF of a galvanic cell, consisting of an electrode located under standard conditions, on which this half-reaction occurs, and a standard hydrogen electrode.

The hydrogen electrode is inconvenient in operation, therefore, in practice, secondary standard electrodes are used as standard electrodes, the potential of which relative to the SHE is determined with high accuracy. One such electrode is the silver chloride electrode,

The sign of the standard half-reaction potential depends on the chosen direction of the half-reaction. When the direction changes, the sign changes to the opposite. For example, for the half-reaction (A) E 0 \u003d +0.771 V, therefore, for the inverse half-reaction (B) E 0 \u003d - 0.771 V.

(A) Fe 3+ +  Fe 2+ (B) Fe 2+ -  Fe 3+

The potential characterizing the recovery process, for example, such as (A), is called restorative, and the potential characterizing the oxidation process, for example, such as (B) - oxidative. At present, the value of the electrode potential of the half-reaction is usually referred to as the process of reducing the oxidized form

The greater the value of the electrode potential, the stronger the oxidizing properties of the oxidized form of the substance and the weaker reducing properties of its reduced form. For example, the permanganate ion under standard conditions in an acidic environment is a stronger oxidizing agent than the dichromate ion.

Cr 2 O 7 2- + 14H + +  2Cr 3+ + 7H 2 O E 0 = +1.33 V

MnO 4 - + 8H + +  Mn 2+ + 4H 2 O E 0 = +1.51 V

If for the half-reaction of interest to us, the value of E 0 in the reference literature, for one reason or another, is not given, then it can be calculated using the potentials of other half-reactions.

Example 7.1.Calculate the value of E 0 for redox pairFe 3+ / Fe if it is known that

Fe 2+ + 2Fe( \u003d -0.473V) Fe 3+ +Fe 2+ ( = +0.771V)

When adding the first and second equations, we get the equation of the half-reaction of interest to us:

Fe 3+ + 3Fe

The value of the standard electrode potential of this half-reaction will not be equal to the sum of and, i.e. 0.298V. The value of E 0 does not depend on the amount of substance (potential is an intensive, not an extensive quantity), therefore Potentials cannot be added.

Unlike the electrode potential, G depends on the amount of substance, therefore G 3 =G 1 +G 2. Hence

The difference between the electrode potentials of the oxidizing agent involved in the direct reaction and the oxidized form of the reducing agent formed during the reaction is calledEMF of the reaction ( E).

By the magnitude of the EMF, one can judge whether or not the spontaneous occurrence of this reaction is possible or not.

Example 7.2.Determine whether the reaction of oxidation of iodide ions by ions can spontaneously proceed under standard conditionsFe 3+ .

2Fe 3+ + 2I -  2Fe 2+ + I 2

=
-
= 0.771 - 0.536 = 0.235V

This reaction can proceed spontaneously in the forward direction.