A number of standard electrode potentials quantitatively characterizes the reducing ability of metal atoms and the oxidizing ability of their ions.

A number of standard electrode potentials make it possible to resolve the issue of the direction of spontaneous occurrence of redox reactions. As in the general case, any chemical reaction, the determining factor here is the sign of the change in the isobaric potential of the reaction. But this means that the first of these systems will act as a reducing agent, and the second as an oxidizing agent. With direct interaction of substances, the possible direction of the reaction will, of course, be the same as when it is carried out in a galvanic cell.

A number of standard electrode potentials make it possible to resolve the issue of the direction of spontaneous occurrence of redox reactions. As in the general case of any chemical reaction, the determining factor here is the sign of the change in the Gibbs energy of the reaction. But this means that the first of these systems will act as a reducing agent, and the second as an oxidizing agent. With direct interaction of substances, the possible direction of the reaction will, of course, be the same as when it is carried out in a galvanic cell.

A number of standard electrode potentials characterize the chemical properties of metals.

A number of standard electrode potentials characterize the chemical properties of metals. It is used when considering the sequence of ion discharge during electrolysis, as well as when describing the general properties of metals.

A number of standard electrode potentials make it possible to resolve the issue of the direction of spontaneous occurrence of oxidizing and non-reducing reactions. As in the general case of any chemical reaction, the determining factor here is the change in the isobaric potential of the reaction. But this means that the first of these systems will act as a reducing agent, and the second - as an oxidizing agent. With direct interaction of substances, the possible direction of the reaction will, of course, be the same as when carried out in a galvanic cell.

A number of standard electrode potentials characterize the chemical properties of metals. It is used to determine the discharge sequence of ions during electrolysis, as well as to describe the general properties of metals. In this case, the values ​​of standard electrode potentials quantitatively characterize the reducing ability of metals and the oxidizing ability of their ions.

Part II. INORGANIC CHEMISTRY

Section 12. GENERAL PROPERTIES OF METALS

§ 12.5. A range of standard electrode potentials

IN high school you study the electrochemical voltage series of metals. Its more precise name is a series of standard electrode potentials of metals. For some metals it is given in table. 12.1. How does such a series consist? Why, for example, is sodium in it after calcium? How to use this nearby?

The answer to the first question can be given on the basis of the material already studied. When any metal is immersed in an electrolyte solution, a potential difference arises at the metal/solution interface, which is called the electrode potential or electrode potential. The potential of each electrode depends on the nature of the metal, the concentration of its ions in the solution and temperature.

It is not possible to directly measure the potential of an individual electrode. Therefore, electrode potentials are measured relative to a standard hydrogen electrode, the potential of which is conventionally taken to be zero at all temperatures. The hydrogen electrode consists of platinum

Rice. 12.3. Standard hydrogen electrode

platinum plate coated with platinum black (electrolytically deposited platinum), which is immersed in a solution of sulfuric acid with a hydrogen ion concentration of 1 mol/l and is washed by a stream of gaseous hydrogen under a pressure of 101.325 kPa at 25°C (Fig. 12.3).

Molecular hydrogen, passing through the solution, dissolves and approaches the surface of the platinum. On the surface of platinum, hydrogen molecules are split into atoms and adsorbed (fixed on the surface). Adsorbed hydrogen atoms H ads are ionized:

H ads - e - -> H + ,

and hydrogen ions, adding electrons, pass into the adsorbed state:

N + + e - -> N adc.

More complete equilibrium in the hydrogen electrode is expressed by the diagram:

2Н + + 2е - ⇆ 2 H a ds (Pt) = Н 2.

The middle part of this equilibrium is of course omitted, although it should be borne in mind what a large role platinum plays in establishing such an equilibrium state.

If now a plate of any metal, immersed in a solution of its salt with a concentration of metal ions of 1 mol/l, is connected to a standard hydrogen electrode, as shown in Fig. 12.4, then a galvanic element (electrochemical circuit) is formed, the electromotive force (abbreviated EMF) of which is easy to measure. This EMF is called the standard electrode potential of a given electrode (usually denoted E°). So, the electrode potential is the EMF of a galvanic cell (electrochemical circuit), which consists of

Rice. 12.4. Galvanic circuit for measuring the standard electrode potential of a metal:

1 - measured electrode;

2 - potentiometer;

3 - standard hydrogen electrode;

4 - potassium chloride solution

the electrode under study and the standard hydrogen electrode.

Such a circuit is shown in Fig. 12.4. Electrode potential is also called redox potential.

When designating electrode potentials E and standard electrode potentials E°, it is customary to put an index on the signs, which corresponds to the system to which this potential belongs. Thus, the standard electrode potential of the system is 2H + + 2e -⇆ H 2 denotes E°2H+/H 2, systems Li + + e - ⇆ Li - E ° Li +/ Li, and the systems M nO - 4 + 8H + + 5e - ⇆ Mn 2+ + 4H 2 O write E°M n O - 4 + 8 H+/M n 2 ++ 4H 2 O .

By placing metals in increasing order of the algebraic value of their standard electrode potentials, we obtain the series presented in Table. 12.1. It can also include other redox systems (including non-metallic ones), respectively, up to E° values, for example E ° C l 2 / C l - = 1.36 V, E ° F 2 / 2 F - = 2.87 V, E ° S / S 2- = -0.51 V, etc. The series presented in table. 12.1 can be considered only as a fragment from a number of standard electrode potentials of redox systems in aqueous solutions at 25°C, composed of the most important metals 1. Historically, this series was preceded by the “Vitisky series” of G. M. Beketov.

Symbol H g applied to a mercury electrode immersed in a solution of hydrargium (I) salt, the ion of which is usually depicted as a dimer:

Н g 2+ 2 + 2е - = 2Н g.

Most standard electrode potentials can be determined experimentally. However, for alkali and alkaline earth metals the E° value is calculated only theoretically, since these metals interact with water.

A number of standard electrode potentials characterize the chemical properties of metals. It is used to find out in what sequence ions are reduced during electrolysis (§ 7.7), as well as to describe other properties of metals (§ 10.9 and 12.5).

1 In the USA, opposite signs of electrode potentials are accepted: the most positive (+3.04 V) electrode Li+/Li and the most negative (-2.87 V) system F 2 /2 F - . This order of reference can also be seen in American educational literature translated into Russian.

The lower 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, and gold is the weakest. Conversely, gold ion Au is the strongest oxidizing agent, and lithium ion Li+ - the weakest (in Table 12.1 the growth of these properties is indicated by arrows).

Each metal in a series of standard electrode potentials has the property of displacing all subsequent metals from solutions of their salts. However, this does not mean that repression will necessarily occur in all cases. Thus, aluminum displaces copper from a solution of cuprum(II) chloride CuC l 2 , but practically does not displace it from cuprum(II) sulfate solution CuSO4 . This is explained by the fact that chloride ions C l - destroy the protective surface film on aluminum much faster compared to sulfate ions SO2-.

Very often, based on a number of standard electrode potentials, equations are written for the reactions of the displacement of metals from solutions of their salts by more active alkali and meadow-earth metals and, naturally, they are mistaken. In this case, the displacement of metals does not occur, because alkali and alkaline earth metals themselves react with water.

All metals that have negative values ​​of standard electrode potentials, that is, standing in the series before hydrogen, displace hydrogen from dilute acids (such as HC l or H 2 SO 4 ) and at the same time dissolve in them. However, lead is practically insoluble in dilute solutions of sulfuric acid. This happens because a protective layer of low-grade lead sulfate salt immediately forms on the surface of the lead. PbSO4 , which disrupts the contact of the solution with the metal. Metals in the row after hydrogen do not displace it from acids.

From the examples given, we can conclude that a number of standard electrode potentials should be used, taking into account the characteristics of the processes under consideration. The most important thing is to keep in mind that a number of standard electrode potentials can only be used for aqueous solutions and that it characterizes the chemical activity of metals only in redox reactions that occur in an aqueous environment.

In the series of standard electrode potentials, sodium is located after calcium Ca: it has a higher algebraic value of the standard electrode potential.

The EMF of any galvanic cell can be calculated from the difference in standard electrode potentials E°. It should be borne in mind that EMF is always a positive quantity. Therefore, from the potential of the electrode, which has a greater algebraic value, it is necessary to subtract the potential of the electrode, whose algebraic value is smaller. For example, the emf of a copper-zinc cell under standard conditions will be 0.34 - (-0.76) = 1.1 V.

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

then we get a series of metal stresses. In addition to metals, this series will always include hydrogen, which allows you to see which metals are capable of displacing hydrogen from aqueous solutions of acids.

Table 19. Series of metal stresses

A number of stresses for the most important metals are given in table. 19. The position of a particular metal in the stress series characterizes its ability to undergo redox interactions in aqueous solutions under standard conditions. Metal ions are oxidizing agents, and metals in the form of simple substances are reducing agents. Moreover, the further a metal is located in the voltage series, the stronger the oxidizing agent in an aqueous solution are its ions, and vice versa, the closer the metal is to the beginning of the series, the stronger the reducing properties of a simple substance - the metal.

Electrode process potential

in a neutral environment it is equal to B (see page 273). Active metals at the beginning of the series, having a potential significantly more negative than -0.41 V, displace hydrogen from water. Magnesium displaces hydrogen only from hot water. Metals located between magnesium and cadmium generally do not displace hydrogen from water. On the surface of these metals, oxide films are formed that have protective effect.

Metals located between magnesium and hydrogen displace hydrogen from acid solutions. At the same time, protective films are also formed on the surface of some metals, inhibiting the reaction. Thus, the oxide film on aluminum makes this metal stable not only in water, but also in solutions of certain acids. Lead does not dissolve in sulfuric acid at its concentration below, since the salt formed when lead reacts with sulfuric acid is insoluble and creates a protective film on the metal surface. The phenomenon of deep inhibition of metal oxidation, due to the presence of protective oxide or salt films on its surface, is called passivity, and the state of the metal in this case is called a passive state.

Metals are capable of displacing each other from salt solutions. The direction of the reaction is determined by their relative position in the series of stresses. When considering specific cases of such reactions, it should be remembered that active metals displace hydrogen not only from water, but also from any aqueous solution. Therefore, the mutual displacement of metals from solutions of their salts practically occurs only in the case of metals located in the series after magnesium.

Beketov was the first to study in detail the displacement of metals from their compounds by other metals. As a result of his work, he arranged metals according to their chemical activity into a displacement series, which is the prototype of a series of metal stresses.

The relative position of some metals in the stress series and in the periodic table at first glance does not correspond to each other. For example, according to the position in the periodic table, the chemical activity of potassium should be greater than sodium, and sodium - greater than lithium. In the series of voltages, lithium is the most active, and potassium occupies a middle position between lithium and sodium. Zinc and copper, according to their position in the periodic table, should have approximately equal chemical activity, but in the voltage series, zinc is located much earlier than copper. The reason for this kind of inconsistency is as follows.

When comparing metals occupying one or another position in the periodic table, the ionization energy of free atoms is taken as a measure of their chemical activity - reducing ability. Indeed, when moving, for example, from top to bottom along the main subgroup of group I of the periodic system, the ionization energy of atoms decreases, which is associated with an increase in their radii (i.e., with a greater distance of outer electrons from the nucleus) and with increasing screening of the positive charge of the nucleus by intermediate electronic layers (see § 31). Therefore, potassium atoms exhibit greater chemical activity - they have stronger reducing properties - than sodium atoms, and sodium atoms exhibit greater activity than lithium atoms.

When comparing metals in a series of voltages, the work of converting a metal in a solid state into hydrated ions in an aqueous solution is taken as a measure of chemical activity. This work can be represented as the sum of three terms: the atomization energy - the transformation of a metal crystal into isolated atoms, the ionization energy of free metal atoms and the hydration energy of the resulting ions. Atomization energy characterizes the strength of the crystal lattice of a given metal. The energy of ionization of atoms - the removal of valence electrons from them - is directly determined by the position of the metal in the periodic table. The energy released during hydration depends on the electronic structure of the ion, its charge and radius.

Lithium and potassium ions, having the same charge but different radii, will create unequal electric fields around themselves. The field generated near small lithium ions will be stronger than the field near large potassium ions. From this it is clear that lithium ions will hydrate with the release of more energy than potassium ions.

Thus, during the transformation under consideration, energy is expended on atomization and ionization and energy is released during hydration. The lower the total energy consumption, the easier the entire process will be and the closer to the beginning of the stress series the given metal will be located. But of the three terms of the general energy balance, only one - the ionization energy - is directly determined by the position of the metal in the periodic table. Consequently, there is no reason to expect that the relative position of certain metals in the stress series will always correspond to their position in the periodic table. Thus, for lithium, the total energy consumption turns out to be less than for potassium, according to which lithium comes before potassium in the voltage series.

For copper and zinc, the energy expenditure for the ionization of free atoms and the energy gain during ion hydration are close. But metallic copper forms a stronger crystal lattice, than zinc, as can be seen from a comparison of the melting temperatures of these metals: zinc melts at , and copper only at . Therefore, the energy spent on the atomization of these metals is significantly different, as a result of which the total energy costs for the entire process in the case of copper are much greater than in the case of zinc, which explains the relative position of these metals in the stress series.

When passing from water to non-aqueous solvents, the relative positions of metals in the voltage series may change. The reason for this is that the solvation energy of ions various metals changes differently when moving from one solvent to another.

In particular, the copper ion is solvated quite vigorously in some organic solvents; This leads to the fact that in such solvents copper is located in the voltage series before hydrogen and displaces it from acid solutions.

Thus, in contrast to the periodic system of elements, a series of metal stresses is not a reflection of a general pattern, on the basis of which it is possible to give versatile Characteristics chemical properties metals A series of voltages characterizes only the redox ability Electrochemical system“metal - metal ion” under strictly defined conditions: the values ​​​​given in it refer to aqueous solution, temperature and unit concentration (activity) of metal ions.

Electrochemical corrosion of metal. Cathodic protection. Anodic protection. Passive protection. Electrode potentials - table.

In the vast majority of cases, metal corrosion refers to the oxidation of a material. In practice, the greatest harm is caused by the so-called. electrochemical corrosion accompanied by active transfer of matter. Metal surfaces are susceptible to electrochemical destruction (corrosion) when they come into contact with electrolytes (corrosion agents). Such agents can be atmospheric gases, such as sea, city or industrial air (i.e. sulfur dioxide, hydrogen chloride and sulfite, etc.) or active liquids - brines, alkalis, sea ​​water etc. (for example, sweaty handprints).

If a galvanic couple is formed as a result of the contact of a corrosion agent on metal surfaces, then the transfer of a substance from one electrode of the couple to another is intensified many times over. The corrosion rate is determined by the difference in the electrode potentials of the pair. This process is usually meant when talking about electrochemical corrosion.

Having a tendency to give up electrons, due to the negative electrode potential, most metals oxidize during the corrosion process. If a certain additional positive potential is applied to the protected object = a certain negative potential of the order of a tenth of a volt is maintained on it, then the probability of an oxidation reaction drops almost to zero. This method protection is usually meant when talking about cathodic protection.

If a certain amount of a substance with a lower electrode potential (for example, zinc or magnesium to protect iron) is placed at the point of probable corrosion, then the oxidation reaction will take place on it. Good electrical contact must be ensured between this additional protective anode(sacrificial anode) and protected metal. Have you guessed why pipes are galvanized? What about iron sheets for roofing? Naturally, when the protective anode dissolves entirely, everything will go as usual.

Under passive protection understand the coating of the protected sample with a dielectric to prevent the occurrence of a galvanic circuit. For example, you can paint a metal structure oil paint etc.

Table. Standard electrode potentials of some substances:

Electrochemical systems

general characteristics

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 cause 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. Job 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 1M 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 will 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).

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 that are in the voltage series before 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/dm 3), 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

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, therefore, low specific capacity, hydrogen evolution during charging, 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

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% . Nickel-cadmium batteries are of greatest practical importance. Cadmium-nickel battery diagram

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

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.