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 engineering 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 with higher speed how ions move from the metal into 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 ​​are determined experimentally electrode potentials relative to the value of the other 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).

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

By arranging the metals in ascending order of their standard electrode potentials, a series of Nikolai Nikolaevich Beketov (1827-1911) voltages, 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 oxidative 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 metals differ in the values ​​of electrode potentials, the more higher value e.m.f. 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), natural logarithm decimal and substituting the values ​​of R, T and F, we get

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

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 due to 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, large number 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

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 nickel-cadmium 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.

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 occur 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 completely, 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:

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.

A number of standard electrode potentials quantitatively characterize 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 of 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. In direct interaction of substances possible direction 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 chemical properties 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.  

Li, K, Ca, Na, Mg, Al, Zn, Cr, Fe, Pb, H 2 , Cu, Ag, Hg, Au

The further to the left a metal is in the series of standard electrode potentials, the stronger the reducing agent it is; the strongest reducing agent is lithium metal, gold is the weakest, and, conversely, gold (III) ion is the strongest oxidizing agent, lithium (I) is the weakest .

Each metal is capable of reducing from salts in solution those metals that are in the series of stresses after it; for example, iron can displace copper from solutions of its salts. However, remember that alkali and alkaline earth metals will react directly with water.

Metals located in the voltage series to the left of hydrogen are capable of displacing it from solutions of dilute acids, while dissolving in them.

The reduction activity of a metal does not always correspond to its position in the periodic table, because when determining a metal’s place in a series, not only its ability to donate electrons is taken into account, but also the energy expended on the destruction of the metal’s crystal lattice, as well as the energy expended on the hydration of ions.

Interaction with simple substances

WITH oxygen Most metals form oxides - amphoteric and basic:

4Li + O 2 = 2Li 2 O,

4Al + 3O 2 = 2Al 2 O 3.

Alkali metals, with the exception of lithium, form peroxides:

2Na + O 2 = Na 2 O 2.

WITH halogens metals form salts of hydrohalic acids, for example,

Cu + Cl 2 = CuCl 2.

WITH hydrogen the most active metals form ionic hydrides - salt-like substances in which hydrogen has an oxidation state of -1.

2Na + H2 = 2NaH.

WITH gray metals form sulfides - salts of hydrogen sulfide acid:

WITH nitrogen Some metals form nitrides; the reaction almost always occurs when heated:

3Mg + N2 = Mg3N2.

WITH carbon carbides are formed:

4Al + 3C = Al 3 C 4.

WITH phosphorus – phosphides:

3Ca + 2P = Ca 3 P 2 .

Metals can interact with each other, forming intermetallic compounds :

2Na + Sb = Na 2 Sb,

3Cu + Au = Cu 3 Au.

Metals can dissolve into each other at high temperatures without reacting, forming alloys.

Alloys

Alloys are called systems consisting of two or more metals, as well as metals and non-metals, which have characteristic properties inherent only in the metallic state.

The properties of alloys are very diverse and differ from the properties of their components, for example, in order for gold to become harder and more suitable for making jewelry, silver is added to it, and an alloy containing 40% cadmium and 60% bismuth has a melting point of 144 °C, i.e. much lower than the melting point of its components (Cd 321 °C, Bi 271 °C).

The following types of alloys are possible:

Molten metals are mixed with each other in any ratio, dissolving in each other indefinitely, for example, Ag-Au, Ag-Cu, Cu-Ni and others. These alloys are homogeneous in composition, have high chemical resistance, and conduct electric current;

The straightened metals are mixed with each other in any ratio, but when cooled they separate, and a mass is obtained consisting of individual crystals of components, for example, Pb-Sn, Bi-Cd, Ag-Pb and others.