Electrochemical processes. Electrode potential is a series of standard electrode potentials for metals. Nernst equation. A range of standard electrode potentials

26.09.2019

Metals include s-elements of groups 1 and 2, all d- and f-elements, as well as a number of p-elements of the main subgroups: 3 (except boron), 4 (Ge, Sn, Pb), 5 (Sb, Bi) and Ro. The most typical metal elements are located at the beginning of periods. Earlier we talked about the fact that highly delocalized bonds occur in metals. This is due to the fact that, due to the screening effect, the valence electrons in metal atoms are weaker attracted to the nucleus and the first ionization energies for them are relatively low. At our usual temperature (about 300 K), which is quite far from absolute zero, the energy of thermal motion is sufficient for the free movement of electrons throughout the metal.

Since the bond in metals is highly delocalized and extends over the entire crystal, metals have high plasticity, electrical and thermal conductivity. Silver and copper have the highest electrical and thermal conductivity, and mercury has the lowest. The latter is also the most fusible metal (-38.9 C). The most refractory metal is tungsten (3390 C). Such a large difference in the melting and boiling points is explained by the presence in metals, in addition to the metallic bond, of a certain proportion covalent bonds, especially for transition elements with a large number of valence electrons.

Let's consider the electronic configurations of mercury and tungsten.

Hg – 5d 10 6s 2; W – 5d 4 6s 2 . The intermolecular interaction between mercury atoms is very small, so small that in general, at high density, due to the gravity of the atoms, it is the most fusible metal. Since all sublevels in the mercury atom are filled, the formation of covalent bonds is generally impossible, and the metallic bond is quite weak, weaker than in alkali metals, which are generally the most fusible among all metals. On the contrary, the formation of four valence bonds at once is possible in the W atom. In addition, metallic bonding is the strongest of all 5d elements, and the atoms themselves are heavier than their electronic counterparts: Mo and Cr. The combination of these factors leads to the greatest refractoriness of tungsten.

The electronic configuration of osmium (5d 6 6s 2) is such that it lacks 4 electrons before completing the 5d sublevel, so it is most strongly capable of attracting electrons from neighboring atoms, which causes a shortening of the metal-metal bond. Therefore, osmium has the highest density (22.4 g/cm3).

IN pure form metals are relatively rare. Basically, these are chemically inert metals (gold, as well as platinum group metals - platinum, rhodium, iridium, osmium, etc.). Silver, copper, mercury, and tin can be found both in the native state and in the form of compounds. The remaining metals occur in the form of compounds called ores.

Metals are obtained from their compounds by reducing them from oxides. C, CO, active metals, hydrogen, and methane are used as reducing agents. If the ore is metal sulfide (ZnS, FeS 2), then it is first converted into oxide. The reduction of metals from their compounds by other metals is called metallothermy. Some metals are extracted from solutions of their salts by electrolysis, for example, aluminum or sodium. All methods for obtaining metals from their compounds are based on redox processes.

The process of electron transfer in a redox half-reaction can be represented by the following general equation:

The process of electron transition corresponds to a change in the Gibbs energy equal to ∆G = –nFE, where F (Faraday’s constant, corresponds to the amount of electricity required for the reduction or oxidation of one mole of a substance) = 96500 C/mol, n is the number of electrons, E is the electrode potential, B is the voltage difference between the oxidizing agent and the reducing agent. On the other hand, ∆G = –RTlnK = –nFE; RTlnK = nFE. Hence E = RTlnK/nF. Since K = /, and 2.3lnK = logK, the dependence of the electrode potential on the concentrations of substances - participants in the electrode process - and on temperature is expressed by the following equation:

E = E 0 + log/ – Nernst equation.

At standard temperature (298 K) the equation takes the form:

E = E 0 + 0.059lg/

The concentration of the oxidizing agent is always given in the numerator, and the potential is always given for the reduction half-reaction: Ox + ne = Red.

At equilibrium concentrations of the oxidizing agent and the reducing agent equal to unity, E = E 0 is the standard electrode potential: this is the potential of a given electrode process at unit concentrations of all substances. Since the absolute value of standard electrode potentials cannot be determined, then the half-reaction potential is taken as the starting point: 2Н + + 2е = Н 2 . The potential of this electrode process is assumed to be 0 at unit concentrations of the hydrogen cation. The hydrogen electrode consists of a platinum plate, which is immersed in a solution of sulfuric acid with [H + ] = 1 mol/l and washed by a stream of H 2 under a pressure of 101325 Pa at 298 K.

The electrode potential is the EMF of a galvanic cell, which consists of the electrode under study and a standard hydrogen electrode. By arranging metals in increasing order of the magnitude of their electrode potentials, we obtain a number of standard electrode potentials of metals. It characterizes the chemical properties of metals. Each metal in the series displaces all subsequent metals from the solution of their salts. Metals in the row to the left of hydrogen displace it from acid solutions.

The potential of any redox reaction can be calculated based on the values of the half-reaction potentials.

Let's consider a simple example: Zn + 2HCl = ZnCl 2 + H 2. For this process, two half-reactions take place:

Zn 2+ + 2e = Zn 0 E 0 (Zn 2+ /Zn) = –0.76 B

2H + + 2e = H 2 0 E 0 (2H + /H 2) = 0.00 B

Since the potential of the second half-reaction is higher than the first, the second half-reaction will proceed from left to right, that is, towards the formation of hydrogen molecules. The first half-reaction will proceed from right to left, that is, towards the formation of zinc cations.

When considering the production of metals, we talked about the fact that a number of metals are reduced from their oxides by other, more active metals. For example, magnesium can reduce copper from copper(II) oxide. Let's compare two half-reactions:

Cu 2+ + 2e = Cu E 0 = +0.34 V

Mg 2+ + 2e = Mg E 0 = –2.36 V

The potential of the first half-reaction is higher than the second and it is the one that will proceed from left to right, and the second - from right to left.

Thus, to determine the direction of redox reactions, it is necessary to write down two half-reactions from the oxidized form to the reduced form and compare their potentials. A reaction with a higher potential will proceed from left to right, and one with a lower potential will proceed from right to left.

Almost all reactions of metals are redox processes and to determine their direction it is necessary, first of all, to take into account the potentials of each of the half-reactions in the redox process. But, besides, there are exceptions. For example, lead is insoluble in sulfuric acid, despite the fact that the potential of the Pb 2+ /Pb pair is –0.15 V. The fact is that lead sulfate is insoluble and its formation prevents further oxidation of lead.

Lecture 15.

Electrolysis.

In solutions and melts of electrolytes there are oppositely charged ions (cations and anions), which are in constant movement. If inert (graphite) electrodes are immersed in this kind of liquid, for example, in a melt of sodium chloride (melts at 801 0 C) and a constant electric current is passed, then the ions under the influence of an external electric field will move towards the electrodes, cations - towards the cathode, and anions - towards anode. Sodium cations, having reached the cathode, accept electrons from it and are reduced to metallic sodium:

Chloride ions are oxidized at the anode:

2Сl – – 2e = Cl 2 0

As a result, metallic sodium is released at the cathode, and molecular chlorine at the anode. The overall equation for the electrolysis of molten sodium chloride is as follows.

K: Na + + e = Na 0 2

A: 2Сl – – 2e = Cl 2 0 1

2Na + + 2Сl – electrolysis ® 2Na 0 + Cl 2 0

2NaСl = 2Na + Cl 2

This reaction is redox: an oxidation process occurs at the anode, and a reduction process occurs at the cathode.

The redox process occurring on the electrodes during the passage electric current through a melt or electrolyte solution is called electrolysis.

The essence of electrolysis is the implementation of chemical reactions using electrical energy. In this case, the cathode gives electrons to cations, and the anode accepts electrons from anions. The action of direct electric current is much stronger than the action of chemical reducing agents and oxidizing agents. It was through electrolysis that fluorine gas was first obtained.

Electrolysis was carried out in a solution of potassium fluoride in hydrofluoric acid. IN in this case Fluorine is released at the anode, and hydrogen is released at the cathode. Electrolysis is carried out in an electrolytic bath.

It is necessary to distinguish between the electrolysis of molten electrolytes and their solutions. IN the latter case water molecules may participate in the processes. For example, during the electrolysis of an aqueous solution of sodium chloride on inert (graphite) electrodes, water molecules are reduced at the cathode instead of sodium cations.

2H 2 O + 2e = H 2 + 2OH –

and chloride ions are oxidized at the anode:

2Сl – – 2e = Cl 2 0

As a result, hydrogen is released at the cathode, chlorine is released at the anode, and sodium hydroxide molecules accumulate in the solution. General equation electrolysis of an aqueous solution of sodium chloride has the form:

K: 2H 2 O + 2e = H 2 + 2OH –

A: 2Сl – – 2e = Cl 2 0

2H 2 O + 2Сl – = H 2 + Cl 2 + 2OH –

By the way, this is how industry produces hydroxides of all alkali and some alkaline earth metals, as well as aluminum.

What is the difference between the electrolysis of melts and aqueous solutions of electrolytes? Reduction processes at the cathode of aqueous solutions of electrolytes depend on the value of the standard electrode potentials of metals, namely, they most often act as cations that are reduced at the cathode. There are three possible options here:

1. Metal cations that have a standard electrode potential are higher than that of hydrogen, that is, more than zero during electrolysis are completely reduced at the cathode (copper, silver, gold and others).

2. Metal cations having very small value standard electrode potential (from lithium to aluminum inclusive), are not reduced at the cathode, but water molecules are reduced.

3. Metal cations, whose standard electrode potential is less than that of hydrogen, but greater than that of aluminum, are reduced during electrolysis at the cathode along with water molecules.

If several metal cations are simultaneously present in an aqueous solution, then during electrolysis their release at the cathode proceeds in the order of decreasing algebraic value of the standard electrode potential of the corresponding metal. For example, when analyzing bronze type BrAZh or BrAZhMts (copper, aluminum, iron and manganese), you can select a certain current value, separate the copper onto an inert electrode (for example, platinum), pull out the electrode, weigh it and determine the copper content. Then separate the aluminum and determine its content. This method is good for separating metals with a positive standard electrode potential.

All electrodes are divided into insoluble (inert) - carbon, graphite, platinum, iridium. Soluble - copper, silver, zinc, cadmium, nickel and others. The concept of a soluble electrode is important for the anode, since it is the one that is capable of dissolving during electrolysis. At the insoluble anode, during electrolysis, oxidation of anions or water molecules occurs. In this case, the anions of oxygen-free acids are quite easily oxidized. If anions of oxygen-containing acids are present in the solution, then water molecules are oxidized at the anode, releasing oxygen according to the reaction:

2H 2 O – 4e = O 2 + 4H +

During electrolysis, the soluble anode itself oxidizes, donating electrons to the external electrical circuit and passing into solution:

A: Me Û Me n+ + ne –

Let's look at examples of electrolysis of melts and electrolyte solutions.

Electrochemistry - a branch of chemistry that studies the processes of the occurrence of potential differences and the conversion of chemical energy into electrical energy (galvanic cells), as well as the implementation of chemical reactions due to the expenditure of electrical energy (electrolysis). These two processes, which have a common nature, are widely used in modern technology.

Galvanic cells are used as autonomous and small-sized energy sources for machines, radio devices and control devices. Using electrolysis, various substances are obtained, surfaces are treated, and products of the desired shape are created.

Electrochemical processes do not always benefit humans, and sometimes cause great harm, causing increased corrosion and destruction of metal structures. In order to skillfully use electrochemical processes and combat undesirable phenomena, they must be studied and be able to regulate.

The reason for the occurrence of electrochemical phenomena is the transfer of electrons or a change in the oxidation state of atoms of substances participating in electrochemical processes, that is, redox reactions occurring in heterogeneous systems. In redox reactions, electrons are directly transferred from the reducing agent to the oxidizing agent. If the processes of oxidation and reduction are spatially separated, and electrons are directed along a metal conductor, then such a system will represent a galvanic cell. The reason for the occurrence and flow of electric current in a galvanic cell is the potential difference.

Electrode potential. Measuring electrode potentials

If you take a plate of any metal and lower it into water, then the ions of the surface layer, under the influence of polar water molecules, come off and hydrate into the liquid. As a result of this transition, the liquid is charged positively and the metal negatively, since an excess of electrons appears on it. The accumulation of metal ions in the liquid begins to inhibit the dissolution of the metal. A mobile equilibrium is established

Me 0 + mH 2 O = Me n + × m H 2 O + ne -

The state of equilibrium depends both on the activity of the metal and on the concentration of its ions in solution. In the case of active metals in the voltage series up to hydrogen, interaction with polar water molecules ends with the separation of positive metal ions from the surface and the transition of hydrated ions into solution (Fig. b). The metal becomes negatively charged. The process is oxidation. As the concentration of ions near the surface increases, the reverse process becomes possible - the reduction of ions. The electrostatic attraction between cations in solution and excess electrons on the surface forms an electrical double layer. This leads to the appearance of a certain potential difference, or potential jump, at the interface between the metal and the liquid. The potential difference that arises between a metal and its surrounding aqueous environment is called electrode potential. When a metal is immersed in a solution of a salt of that metal, the equilibrium shifts. Increasing the concentration of ions of a given metal in solution facilitates the process of transition of ions from solution to metal. Metals whose ions have a significant ability to pass into solution will be positively charged in such a solution, but to a lesser extent than in pure water.

For inactive metals, the equilibrium concentration of metal ions in solution is very small. If such a metal is immersed in a solution of a salt of this metal, then positively charged ions are released on the metal at a faster rate than the transition of ions from the metal to the solution. The metal surface will receive a positive charge, and the solution will receive a negative charge due to the excess salt anions. And in this case, an electric double layer appears at the metal-solution interface, hence a certain potential difference (Fig. c). In the case considered, the electrode potential is positive.

Rice. The process of transition of an ion from a metal to a solution:

a – balance; b – dissolution; c – deposition

The potential of each electrode depends on the nature of the metal, the concentration of its ions in the solution and temperature. If a metal is immersed in a solution of its salt containing one mole metal ion per 1 dm 3 (the activity of which is 1), then the electrode potential will be a constant value at a temperature of 25 o C and a pressure of 1 atm. This potential is called standard electrode potential (E o).

Metal ions having a positive charge, penetrating into the solution and moving in the potential field of the metal-solution interface, expend energy. This energy is compensated by the work of isothermal expansion from a higher concentration of ions on the surface to a lower one in the solution. Positive ions accumulate in the surface layer to a concentration With O, and then go into solution, where the concentration of free ions With. The work of the electric field EnF is equal to the isothermal work of expansion RTln(с o /с). By equating both expressions of work, we can derive the magnitude of the potential

En F = RTln(s o /s), -E = RTln(s/s o)/nF,

where E is the metal potential, V; R – universal gas constant, J/mol K; T – temperature, K; n – ion charge; F – Faraday number; с – concentration of free ions;

с о – concentration of ions in the surface layer.

It is not possible to directly measure the potential value, since it is impossible to experimentally determine the value of the potential. The values of the electrode potentials are determined empirically relative to the value of another electrode, the potential of which is conventionally assumed to be zero. Such a standard or reference electrode is normal hydrogen electrode (n.v.e.) . The structure of the hydrogen electrode is shown in the figure. It consists of a platinum plate coated with electrolytically deposited platinum. The electrode is immersed in a 1 M solution of sulfuric acid (the activity of hydrogen ions is 1 mol/dm3) and is washed by a stream of hydrogen gas under a pressure of 101 kPa and T = 298 K. When platinum is saturated with hydrogen, equilibrium is established on the metal surface, the overall process is expressed by the equation

2Н + +2е ↔ Н 2 .

If a plate of metal immersed in a 1M solution of a salt of this metal is connected by an external conductor to a standard hydrogen electrode, and the solutions are connected by an electrolytic key, then we obtain a galvanic cell (Fig. 32). The electromotive force of this galvanic cell will be the quantity standard electrode potential of a given metal (E O ).

Scheme for measuring standard electrode potential

relative to the hydrogen electrode

Taking zinc in a 1 M solution of zinc sulfate as an electrode and connecting it with a hydrogen electrode, we obtain a galvanic cell, the circuit of which can be written as follows:

(-) Zn/Zn 2+ // 2H + /H 2, Pt (+).

In the diagram, one line indicates the interface between the electrode and the solution, two lines indicate the interface between solutions. The anode is written on the left, the cathode on the right. In such an element, the reaction Zn o + 2H + = Zn 2+ + H 2 takes place, and electrons pass through the external circuit from the zinc to the hydrogen electrode. Standard electrode potential for zinc electrode (-0.76 V).

Taking a copper plate as an electrode, under the specified conditions in combination with a standard hydrogen electrode, we obtain a galvanic cell

(-) Pt, H 2 /2H + //Cu 2+ /Cu (+).

In this case, the reaction occurs: Cu 2+ + H 2 = Cu o + 2H +. Electrons move through the external circuit from the hydrogen electrode to the copper electrode. Standard electrode potential of copper electrode (+0.34 V).

Electrochemical systems

General characteristics

Electrode potential. Measuring electrode potentials

Me 0 + mH 2 O = Me n + × m H 2 O + ne -

Rice. The process of transition of an ion from a metal to a solution:

a – balance; b – dissolution; c – deposition

En F = RTln(s o /s), -E = RTln(s/s o)/nF,

where E is the metal potential, V; R – universal gas constant, J/mol K; T – temperature, K; n – ion charge; F – Faraday number; с – concentration of free ions;

с о – concentration of ions in the surface layer.

2Н + +2е ↔ Н 2 .

Scheme for measuring standard electrode potential

relative to the hydrogen electrode

Taking zinc in a 1 M solution of zinc sulfate as an electrode and connecting it with a hydrogen electrode, we obtain a galvanic cell, the circuit of which can be written as follows:

(-) Zn/Zn 2+ // 2H + /H 2, Pt (+).

Taking a copper plate as an electrode, under the specified conditions in combination with a standard hydrogen electrode, we obtain a galvanic cell

(-) Pt, H 2 /2H + //Cu 2+ /Cu (+).

A number of standard electrode potentials (voltages). Nernst equation

By arranging metals in increasing order of their standard electrode potentials, a series of voltages of Nikolai Nikolaevich Beketov (1827-1911), or a series of standard electrode potentials, is obtained. Numerical values of standard electrode potentials for a number of technically important metals are given in the table.

Metal stress range

A number of stresses characterize some properties of metals:

1. The lower the electrode potential of a metal, the more chemically active it is, the easier it is to oxidize and the more difficult it is to recover from its ions. Active metals in nature exist only in the form of compounds Na, K, ..., are found in nature both in the form of compounds and in the free state of Cu, Ag, Hg; Au, Pt - only in a free state;

2. Metals that have a more negative electrode potential than magnesium displace hydrogen from water;

3. Metals in the voltage series up to hydrogen displace hydrogen from solutions of dilute acids (the anions of which do not exhibit oxidizing properties);

4. Each metal in the series that does not decompose water displaces metals that have more positive values of electrode potentials from solutions of their salts;

5. The more the metals differ in the values of the electrode potentials, the greater the emf value. will have a galvanic cell constructed from them.

The dependence of the electrode potential (E) on the nature of the metal, the activity of its ions in solution and temperature is expressed by the Nernst equation

E Me = E o Me + RTln(a Me n +)/nF,

where E o Me is the standard electrode potential of the metal, and Men + is the activity of metal ions in solution. At a standard temperature of 25 o C, for dilute solutions, replacing activity (a) with concentration (c), the natural logarithm with a decimal one and substituting the values of R, T and F, we obtain

E Me = E o Me + (0.059/n)logс.

For example, for a zinc electrode placed in a solution of its salt, the concentration of hydrated ions Zn 2+ × mH 2 O Let us abbreviate it as Zn 2+ , then

E Zn = E o Zn + (0.059/n) log[ Zn 2+ ].

If = 1 mol/dm 3, then E Zn = E o Zn.

Galvanic cells, their electromotive force

Two metals immersed in solutions of their salts, connected by a conductor, form a galvanic cell. The first galvanic cell was invented by Alexander Volt in 1800. The cell consisted of copper and zinc plates separated by cloth soaked in a solution of sulfuric acid. When a large number of plates are connected in series, the Volta element has a significant electromotive force (emf).

The occurrence of an electric current in a galvanic cell is caused by the difference in the electrode potentials of the metals taken and is accompanied by chemical transformations occurring at the electrodes. Let's consider the operation of a galvanic cell using the example of a copper-zinc cell (J. Daniel - B.S. Jacobi).

Diagram of a copper-zinc Daniel-Jacobi galvanic cell

On a zinc electrode immersed in a solution of zinc sulfate (c = 1 mol/dm 3), zinc oxidation (zinc dissolution) occurs Zn o - 2e = Zn 2+. Electrons enter the external circuit. Zn is a source of electrons. The source of electrons is considered to be the negative electrode - the anode. On a copper electrode immersed in a copper sulfate solution (c = 1 mol/dm3), metal ions are reduced. Copper atoms are deposited on the electrode Cu 2+ + 2e = Cu o. The copper electrode is positive. It is the cathode. At the same time, some SO 4 2- ions pass through the salt bridge into a vessel with a ZnSO 4 solution . Adding up the equations of the processes occurring at the anode and cathode, we obtain the total equation

Boris Semenovich Jacobi (Moritz Hermann) (1801-1874)

or in molecular form

This is a common redox reaction occurring at the metal-solution interface. The electrical energy of a galvanic cell is obtained through a chemical reaction. The considered galvanic cell can be written in the form of a brief electrochemical circuit

(-) Zn/Zn 2+ //Cu 2+ /Cu (+).

A necessary condition for the operation of a galvanic cell is the potential difference, it is called electromotive force of a galvanic cell (emf) . E.m.f. any working galvanic element has a positive value. To calculate the emf. galvanic cell, it is necessary to subtract the value of the less positive potential from the value of the more positive potential. So e.m.f. copper-zinc galvanic cell under standard conditions (t = 25 o C, c = 1 mol/dm 3, P = 1 atm) is equal to the difference between the standard electrode potentials of copper (cathode) and zinc (anode), that is

e.m.f. = E o C u 2+ / Cu - E o Zn 2+ / Zn = +0.34 V – (-0.76 V) = +1.10 V.

When paired with zinc, the Cu 2+ ion is reduced.

The difference in electrode potentials required for operation can be created using the same solution of different concentrations and the same electrodes. Such a galvanic cell is called concentration , and it works by equalizing the concentrations of the solution. An example would be a cell composed of two hydrogen electrodes

Pt, H 2 / H 2 SO 4 (s`) // H 2 SO 4 (s``) / H 2, Pt,

where c` = `; c`` = ``.

If p = 101 kPa, s`< с``, то его э.д.с. при 25 о С определяется уравнением

E = 0.059lg(s``/s`).

At с` = 1 mol-ion/dm 3 emf. element is determined by the concentration of hydrogen ions in the second solution, that is, E = 0.059lgс`` = -0.059 pH.

Determination of the concentration of hydrogen ions and, consequently, the pH of the medium by measuring the emf. the corresponding galvanic element is called potentiometry.

Batteries

Batteries are called galvanic cells of reusable and reversible action. They are capable of converting accumulated chemical energy into electrical energy during discharge, and electrical energy into chemical energy, creating a reserve during charging. Since the e.m.f. batteries are small; during operation they are usually connected into batteries.

Lead acid battery . A lead-acid battery consists of two perforated lead plates, one of which (negative) after charging contains a filler - spongy active lead, and the other (positive) - lead dioxide. Both plates are immersed in a 25 - 30% sulfuric acid solution (Fig. 35). Battery circuit

(-) Pb/ p -p H 2 SO 4 / PbO 2 / Pb(+) .

Before charging, a paste containing, in addition to the organic binder, lead oxide PbO, is smeared into the pores of the lead electrodes. As a result of the interaction of lead oxide with sulfuric acid, lead sulfate is formed in the pores of the electrode plates

PbO + H 2 SO 4 = PbSO 4 + H 2 O .

Batteries are charged by passing electric current

Discharging process

In total, the processes that occur when charging and discharging a battery can be represented as follows:

When charging a battery, the density of the electrolyte (sulfuric acid) increases, and when discharging it decreases. The density of the electrolyte determines the degree of discharge of the battery. E.m.f. lead battery 2.1 V.

Advantages lead-acid battery - high electrical capacity, stable operation, a large number of cycles (discharge-charge). Flaws - large mass and, consequently, low specific capacity, hydrogen evolution during charging, and non-tightness in the presence of a concentrated sulfuric acid solution. Alkaline batteries are better in this regard.

Alkaline batteries. These include T. Edison cadmium-nickel and iron-nickel batteries.

Edison battery and lead battery circuits

Thomas Edison(1847-1931)

They are similar to each other. The difference lies in the material of the negative electrode plates. In the first case they are cadmium, in the second they are iron. The electrolyte is a KOH solution ω = 20% . Greatest practical significance have cadmium-nickel batteries. Cadmium-nickel battery diagram

(-) Cd / KOH solution / Ni 2 O 3 / Ni (+).

The operation of a cadmium-nickel battery is based on a redox reaction involving Ni 3+

E.m.f. of a charged nickel-cadmium battery is 1.4 V.

The table shows the characteristics of the Edison battery and the lead battery.

Rice. 128. Device for measuring the normal potential of a metal

There are several theories explaining the occurrence of current in galvanic cells. The simplest of them was put forward by Nernst (1888) and later developed in detail by Academician L.V. Pisarzhevsky based on ideas about the structure of metals from positively charged ions and free electrons.

Lev Vladimirovich Pisarzhevsky was born in 1874. Chisinau. After graduating from the Faculty of Natural Sciences of Novorossiysk University (Odessa), Pisarzhevsky was left with him to prepare for the title of professor. In 1902 he defended his master's thesis, and in 1913 he was elected professor at the Ekaterinoslav Mining Institute (Dnepropetrovsk). Since 1930, Pisarzhevsky was a full member of the USSR Academy of Sciences.

A prominent scientist and brilliant teacher, Pisarzhevsky boldly used the achievements of physics to study and explain chemical processes. His most important works were devoted to the study of peroxides and peracids, the development of the theory of solutions, the application of electronic theory to chemistry, and the development of the theory of the occurrence of current in galvanic cells.

The generation of current in a galvanic cell occurs as follows. If you immerse any metal in water, its ions begin to go into solution under the influence of the attraction they experience from polar water molecules. As a result, the metal in whichexcess electrons remain and become negatively charged, while the solution becomes positively charged. However, the number of ions that the metal sends into the solution, as experience shows, is very small. The negative charge that appears on the metal as the ions leave begins to attract back the ions that have left the metal, so that soon a state of equilibrium occurs, in which per unit time as many ions leave the metal as return to it:

metal⇄metal ions

(in solution)

The ions that have passed into the solution are not distributed evenly throughout the entire mass of the solution, but due to attraction to the negatively charged metal, they are located near its surface, forming the so-called electric double layer (Fig. 127). As a result, a certain difference potentials.

Lev Vladimirovich Pisarzhevsky (1874-1938)

Let us now suppose that we add to the water in which the metal is immersed a certain amount of salt of the same metal. Due to an increase in the concentration of metal ions in the solution, the equilibrium between them and the metal will be disrupted and some of the ions will go back into the metal. Therefore, into a solution of your salt

the metal should send fewer ions than in clean water, and the less, the higher the concentration of ions in the solution. If the salt concentration is high enough, the ions may not move from the metal into the solution at all, so neither the metal nor the solution will be charged.

Finally, if the concentration of metal ions in the solution is sufficiently high and the activity of the metal is relatively low, the metal not only does not send ions into the solution, but, on the contrary, some of the ions pass from the solution to the metal. In this case, a potential difference also arises between the metal and the solution, but now the solution is charged negatively due to the excess of negative ions of the salt, and the metal is charged positively. In practice, the situation is that some (more active) are always charged negatively in solutions of their salts, while others (less active) are positively charged.

It should be noted that in all cases, when a metal is immersed in a solution of its salt, the amount of ions passing into the solution or released from the solution is so small that it cannot be detected chemically. However, their charge is large enough to create a measurable potential difference.

The theory outlined above very simply explains the mechanism of action of galvanic cells. Consider, for example, a copper-zinc element. In this element, a certain negative charge appears on a zinc plate immersed in a ZnSO 4 solution, and a positive charge appears on copper immersed in a CuSO 4 solution. If they are not connected to each other by a conductor, the appearance of these charges, as we saw above, should immediately stop both the further transition of zinc ions into the solution and the release of copper ions from the solution. But if you connect both plates with a wire, then the electrons accumulated on the zinc will constantly flow to the copper plate, where they are missing. Thus, it becomes possible to send more and more quantities of Zn ions into the solution, while at the copper plate the Cu ions are discharged and released in the form of metallic copper. This process continues until all the copper salt is dissolved or used up.

Rice. 127. Electric double layer

In galvanic cells, the electrode that is destroyed during operation of the cell, sending ions into the solution, is called the anode, and the electrode at which positive ions are discharged is called the cathode.

A galvanic cell can be constructed from any two metals immersed in solutions of their salts. In this case, it is not at all necessary that one metal be charged negatively and the other positively. The only condition for the flow of electrons from one charged body to another is the existence of a potential difference between them. But the latter must arise, no matter what we taken, since the ability to detach electrons and transform into ions is different for all metals. If, for example, a galvanic cell is composed of zinc and iron immersed in normal solutions of their salts, then, although both metals are negatively charged in solutions, some potential difference will still arise between them. When metals are connected by a conductor, electrons will flow from zinc, as a more active metal, to iron; will dissolve and - be released from solution. The reaction occurring in the element is expressed by the equation

Zn + Fe = Fe + Zn

The potential difference that arises between a metal and a solution of its salt is called the electrode potential of the metal and can serve as a measure of its ability to donate electrons or, what is the same, a measure of its chemical activity during reactions in solutions. Therefore, by measuring the potentials of all metals at the same concentrations of their ions, we could quantitatively characterize the activity of the metals.

Unfortunately, direct measurement of these quantities is very difficult and does not give accurate results. This is already clear from the fact that it is impossible, for example, to connect a voltmeter to a solution without immersing a metal conductor in the solution. But then a potential difference arises between the conductor and the solution, so that the voltage shown by the voltmeter will depend on two potential differences: the potential difference between the metal of interest to us and the solution of its salt, and the potential difference between the metal conductor and the same solution.

It is much easier to measure the potential difference (the difference in electron voltage) between two different metal electrodes immersed in solutions of the corresponding salts, that is, to find out how much the potential of one metal is greater or less than the potential of another metal. If we measure the relative potentials of all metals in this way, comparing their potentials with the potential of any one of them, then the resulting numbers will characterize the activity of the metals just as accurately as the absolute values of their potentials.

The so-called normal hydrogen electrode is adopted as a standard electrode, with the potential of which the potentials of other metals are compared. The latter consists of a platinum plate coated with a loose layer of platinum and immersed in a two-normal solution of sulfuric acid. Pressure is continuously passed through the solution in 1 at current pure hydrogen, which, coming into contact with platinum, is quite large quantities is absorbed by it. A platinum plate saturated with hydrogen behaves as if it were made of hydrogen. When it comes into contact with a solution of sulfuric acid, a certain potential difference arises (potential of the hydrogen electrode), which is conventionally accepted as zero when measuring relative potentials.

The potential difference between a metal immersed in a solution of its salt containing 1 gram of metal ion per liter and a normal hydrogen electrode is called the normal potential of the metal.

To measure normal potentials, instruments similar to those shown in Fig. are usually used. 128. Essentially, such a device is a galvanic cell, one of the electrodes of which is the metal being tested, and the other is a hydrogen electrode. Since the potential of a hydrogen electrode is taken to be zero, by measuring the potential difference at the poles of such an element or its electromotive force, we directly find the normal potential of the metal.

In table 27 indicates the normal potentials of the most important metals. They are taken with a minus sign when the potential of the metal is lower than the potential of the hydrogen electrode, and with a plus sign when the potential of the metal is higher than it.

If we arrange metals, including and, according to the decreasing voltage of their electrodes, i.e., according to decreasing negative normal potentials (and increasing positive ones), then the same series of voltages will be obtained.

Table 27

Normal potentials of metals

Metal	Ion	Potential in volts	Metal	Ion	Potential in volts
TO	TO	- 2,92	Ni	Ni	- 0,23
Sa	Sa	- 2,84	Sn	Sn	- 0,14
Na	Na	- 2,713	Pb	Pb	- 0,126
Mg	Mg	- 2,38	n 2	H	0,000
Al	Al	- 1,66	Cu	Cu	+ 0,34
Mn	Mn	- 1,05	Hg	Hg 2	+ 0,798
Zn	Zn	- 0,763	Ag	Ag	+ 0,799
Fe	Fe	- 0,44	Au	Au	+ 1,42

Knowing the normal potentials of metals, it is easy to determine the electromotive force of any element consisting of two metals immersed in solutions of their salts. To do this, you only need to find the difference in the normal potentials of the taken metals.

In order for the electromotive force to have a positive value, the smaller one is always subtracted from the larger potential. For example, the electromotive force of a copper-zinc element is:

e. d.s. = 0.34 - (-0.763) = 1.103

It is clear that it will have such a value if the concentrations of Zn and Cu ions in the corresponding solutions are equal to 1 gram per 1 liter. For other concentrations, metal potentials, and therefore electromotive forces, can be calculated using the formula derived by Nernst:

All electrochemical processes can be divided into two opposing groups: electrolysis processes, in which chemical reactions occur under the influence of an external source of electricity, and processes of the emergence of electromotive force and electric current as a result of certain chemical reactions.

In the first group of processes, electrical energy is converted into chemical energy, in the second, on the contrary, chemical energy is converted into electrical energy.

Examples of both types of processes include processes occurring in batteries. Thus, when a lead battery of an electrical energy generator operates, the following reaction occurs:

Pb + PbO 2 + 4H + + 2SO 4 2- → PbSO 4 + 2H 2 O.

As a result of this reaction, energy is released, which turns into electricity. When the battery is discharged, it is charged by passing electric current through it in the opposite direction.

The chemical reaction also occurs in the opposite direction:

2PbSO 4 + 2H 2 O → Pb + PbO 2 + 4H + + 2SO 4 2-.

In this case, electrical energy turned into chemical energy. The battery now has energy reserves and can be discharged again.

All electrochemical reactions occur when an electric current flows in a circuit. This circle necessarily consists of metal conductors connected in series and an electrolyte solution (or melt). In metal conductors, as we know, current is carried by electrons, in a solution of electrolytes - by ions. The continuity of current flow in the circuit is ensured only when processes occur on the electrodes, i.e. at the metal - electrolyte boundary. On one electrode the process of receiving electrons occurs - reduction, on the second electrode - the process of releasing electrons, i.e. oxidation.

Feature electrochemical processes, unlike conventional chemical ones, is the spatial separation of oxidation and reduction processes. Of these processes, which cannot occur without each other, and consists as a whole chemical process in an electrochemical system.

If you immerse a metal plate (electrode) in an electrolyte solution, a potential difference arises between the plate and the solution, which is called the electrode potential.

Let's consider the reasons for its occurrence. The nodes of the metal crystal lattice contain only positively charged ions. Due to their interaction with polar solvent molecules, they break away from the crystal and go into solution. As a result of this transition, an excess of electrons remains in the metal plate, causing it to acquire a negative charge. Positively charged ions that enter the solution due to electrostatic attraction remain directly at the surface of the metal electrode. An electrical double layer is formed. A potential jump occurs between the electrode and the solution, which is called the electrode potential.

Along with the transition of ions from the metal to the solution, the reverse process also occurs. The rate of transition of ions from metal to solution V 1 may be greater than the rate of reverse transition of ions from solution to metal V 2 (V 2 ˃ V 1).

This difference in speed will result in a decrease in the number of positive ions in the metal and an increase in them in the solution. The metal electrode acquires a negative charge, and the solution acquires a positive charge.

The greater the difference V 1 ‒V 2, the more negative the charge of the metal electrode will be. In turn, the value of V 2 depends on the content of metal ions in the solution; their large concentrations correspond to high speed V 2. Consequently, with increasing concentration of ions in the solution, the negative charge of the metal electrode decreases.

If, on the contrary, the rate of transition of metal ions into solution is less than the rate of the reverse process (V 1< V 2), то на металлическом электроде будет избыток положительных ионов, а в растворе ‒ их нехватка. В таком случае электрод вступит положительный заряд, а раствор ‒ негативного.

In both cases, the potential difference, which arises as a result of the uneven distribution of charges, accelerates the slow process and slows down faster. As a result, a moment will come when the rates of both processes become equal. There will be an equilibrium that will be dynamic. The transition of ions from the metal to the solution and back will occur all the time and in a state of equilibrium. The rates of these processes in equilibrium will be the same (V 1p = V 2p). The amount of electrode potential that is kept in equilibrium is called the equilibrium electrode potential.

The potential that arises between a metal and a solution if the metal is immersed in a solution in which the concentration of ions of this metal is equal to one gram ion is called the normal or standard electrode potential.

If we place the normal potentials of electrode reactions for various metals so that their algebraic values consistently increase, then we obtain the well-known general course chemistry series of stresses. In this row, all elements are placed depending on their electrochemical properties, which are directly related to chemical properties. Thus, all metals located in copper (i.e., with more negative potentials) are relatively easily oxidized, and all metals located after copper are oxidized with rather great difficulty.

K, Na, Ca, Mg, A1, Mn, Zn, Fe,

Ni, Sn, Pb, H2, Cu, Hg, Ag, Au.

Each member of the series, being more active, can displace from the connections any member of the series standing to the right of it in the series of stresses.

Let us consider the mechanism of action of a galvanic cell, the diagram of which is shown in Fig. The element consists of a zinc plate immersed in a zinc sulfate solution and a copper plate immersed in a copper sulfate solution.

Rice. Diagram of a copper-zinc galvanic cell

Both are vessels with solutions, called half-cells, connected to each other by an electrolytic switch to form a galvanic cell. This key (a glass tube filled with electrolyte) allows ions to move from one vessel (half cell) to another. Together, solutions of zinc sulfate and copper sulfate do not mix.

If the electrical circuit is open, then no changes occur in the metal plates or in the solution, but when the circle is closed, current will flow through the circle. Electrons from a place where the negative charge density is higher (i.e. the zinc plate) move to places with a lower negative charge density or to a place with a positive charge (i.e. the copper plate). Due to the movement of electrons, the equilibrium at the metal-solution interface will be disrupted. The excess of negative charges in the zinc plate will decrease, the attractive forces will correspondingly decrease, and some of the zinc ions from the electrical double layer will move into the total volume of the solution. This will lead to a decrease in the rate of transition of Zn 2+ ions from solution to metal. The difference V 1 ‒V 2 (which is zero in the equilibrium state) will increase, and a new amount of zinc ions will move from the metal into the solution. This will cause an excess of electrons to appear in the zinc plate, which will immediately move to the copper plate, and again everything will be repeated continuously. As a result, the zinc dissolves, and an electric current continuously flows in the circle.

It is clear that the continuous movement of electrons from the zinc plate to the copper plate is possible only when they are assimilated on the copper plate. The appearance of excess electrons in the copper plate will lead to rearrangement of the double layer. Negative SO 4 2- ions will repel, and positive copper ions that are in the solution will enter the electrical double layer due to electrostatic attraction caused by the appearance of electrons. The rate of transition of ions to metal V 2 will increase. Cu 2+ ions penetrate into the crystal lattice of the copper plate, adding electrons. It is this process of electron assimilation on the copper plate that will ensure the continuity of the process as a whole.

The magnitude of the EMF E is equal to the difference between the electrode potentials E 1 and E 2 on the electrodes: E = E 1 – E 2.

The processes that occur on the electrodes can be represented by a diagram: on the face there is a zinc plate - electrolyte Zn - 2e - = Zn 2+, on the face there is a copper plate electrolyte Cu 2+ + 2e - = Cu.

As we can see, the processes of zinc oxidation and copper reduction are separated in space; they occur on different electrodes. Generally chemical reaction, which occurs in a copper-zinc element, can be written in ionic form as follows:

Zn + Cu 2+ = Zn 2+ + Cu.

The same picture will be observed in the case when both plates are negatively charged relative to the solution. Let's immerse two copper plates in dilute solutions of copper sulfate. The concentration of copper ions in these solutions is C 1 and C 2 (C 2 > C 1). Let us assume that both plates are negatively charged relative to the solutions. But plate A in a vessel with solution concentration C 1 will be charged more negatively due to the fact that the concentration of copper ions in this vessel is less than in the second vessel, and accordingly the rate of penetration of Cu 2+ ions into the crystal lattice will be less. If you close the circle, then the electrons will move from plate A, where their density is greater, to plate B. On the edge of plate A with the electrolyte, the process Cu° ‒ 2е - = Cu 2+ occurs, on the edge of plate B with the electrolyte Cu 2+ + 2е - + Cu°.

Both plates, as already noted, are negatively charged relative to the solution. But plate A is negatively charged relative to plate B and therefore acts as a negative electrode in a galvanic cell, and plate B acts as a positive electrode.

The magnitude of the EMF, equal to the difference in electrode potentials, will be greater, the greater the difference in ion concentrations in solutions.

Nernst equation- an equation connecting the redox potential of the system with the activities of the substances included in the electrochemical equation and the standard electrode potentials of redox pairs.

Electrode potential, - standard electrode potential, measured in volts;