Standard electrode potentials of metals table. A range of standard electrode potentials

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.

Grosse E., Weissmantel H.

Chemistry for the curious. Basics of chemistry and entertaining experiments.

SMALL COURSE IN ELECTROCHEMISTRY OF METALS

METALS STRESS SERIES

LET'S LOOK BEHIND THE SCENES

Me 1 + Me 2 2+ = Me 1 2+ + Me 2

Moreover, for the first experiment Me 1 = Fe, Me 2 = Cu.
So, the process consists of the exchange of charges (electrons) between atoms and ions of both metals. If we separately consider (as intermediate reactions) the dissolution of iron or the precipitation of copper, we obtain:

Fe = Fe 2+ + 2 e --

Cu 2+ + 2 e-- = Cu

Now consider the case when the metal is immersed in water or in a salt solution, with the cation of which exchange is impossible due to its position in the stress series. Despite this, the metal tends to go into solution in the form of an ion. In this case, the metal atom gives up two electrons (if the metal is divalent), the surface of the metal immersed in the solution becomes negatively charged relative to the solution, and a double electric layer is formed at the interface. This potential difference prevents further dissolution of the metal, so that the process soon stops.
If two different metals are immersed in a solution, they will both charge, but the less active one will be somewhat weaker, due to the fact that its atoms are less prone to losing electrons.
Let's connect both metals with a conductor. Due to the potential difference, a flow of electrons will flow from the more active metal to the less active one, which forms the positive pole of the element. A process occurs in which the more active metal goes into solution, and cations from the solution are released on the more noble metal. Let us now illustrate the somewhat abstract reasoning above (which, moreover, represents a gross simplification) with several experiments.
First, fill a 250 ml beaker to the middle with a 10% solution of sulfuric acid and immerse not too small pieces of zinc and copper in it. We solder or rivet copper wire to both electrodes, the ends of which should not touch the solution.
As long as the ends of the wire are not connected to each other, we will observe the dissolution of zinc, which is accompanied by the release of hydrogen. Zinc, as follows from the voltage series, is more active than hydrogen, so the metal can displace hydrogen from the ionic state. An electrical double layer is formed on both metals. The easiest way to detect the potential difference between the electrodes is with a voltmeter. Immediately after connecting the device to the circuit, the arrow will indicate approximately 1 V, but then the voltage will quickly drop. If you connect a small light bulb that consumes 1 V to the element, it will light up - at first quite strongly, and then the glow will become weak.
Based on the polarity of the device terminals, we can conclude that the copper electrode is the positive pole. This can be proven without a device by considering the electrochemistry of the process. Let's prepare a saturated solution of table salt in a small beaker or test tube, add about 0.5 ml of an alcohol solution of the phenolphthalein indicator and immerse both electrodes closed with wire into the solution. A faint reddish color will be observed near the negative pole, which is caused by the formation of sodium hydroxide at the cathode.
In other experiments, one can place various pairs of metals in a cell and determine the resulting voltage. For example, magnesium and silver will give a particularly large potential difference due to the significant distance between them and a series of voltages, while zinc and iron, on the contrary, will give a very small one, less than a tenth of a volt. By using aluminum, we will not receive practically any current due to passivation.
All these elements, or, as electrochemists say, circuits, have the disadvantage that when measuring current, the voltage across them drops very quickly. Therefore, electrochemists always measure the true magnitude of the voltage in the de-energized state using the voltage compensation method, that is, comparing it with the voltage of another current source.
Let us consider the processes in the copper-zinc element in a little more detail. At the cathode, zinc goes into solution according to the following equation:

Zn = Zn 2+ + 2 e --

Hydrogen ions of sulfuric acid are discharged at the copper anode. They attach electrons coming through the wire from the zinc cathode and as a result, hydrogen bubbles are formed:

2H + + 2 e-- = N 2

After a short period of time, the copper will be covered with a thin layer of hydrogen bubbles. In this case, the copper electrode will turn into a hydrogen one, and the potential difference will decrease. This process is called electrode polarization. The polarization of the copper electrode can be eliminated by adding a little potassium dichromate solution to the cell after the voltage drop. After this, the voltage will increase again, as potassium dichromate will oxidize hydrogen to water. Potassium dichromate acts in this case as a depolarizer.
In practice, galvanic circuits are used whose electrodes are not polarized, or circuits whose polarization can be eliminated by adding depolarizers.
As an example of a non-polarizable element, consider the Daniel element, which was often used in the past as a current source. This is also a copper-zinc element, but both metals are immersed in different solutions. The zinc electrode is placed in a porous clay cell filled with dilute (about 20%) sulfuric acid. The clay cell is suspended in a large glass containing a concentrated solution of copper sulfate, and at the bottom there is a layer of copper sulfate crystals. The second electrode in this vessel is a cylinder made of copper sheet.
This element can be made from a glass jar, a commercially available clay cell (in extreme cases, we use a flower pot, closing the hole in the bottom) and two electrodes of suitable size.
During operation of the element, zinc dissolves to form zinc sulfate, and copper ions are released at the copper electrode. But at the same time, the copper electrode is not polarized and the element produces a voltage of about 1 V. Actually, theoretically, the voltage at the terminals is 1.10 V, but when collecting current we measure a slightly lower value due to the electrical resistance of the cell.
If we do not remove the current from the element, we need to remove the zinc electrode from the sulfuric acid solution, because otherwise it will dissolve to form hydrogen.
A diagram of a simple cell that does not require a porous partition is shown in the figure. The zinc electrode is located at the top of the glass jar, and the copper electrode is located near the bottom. The entire cell is filled with a saturated solution of table salt. Place a handful of copper sulfate crystals at the bottom of the jar. The resulting concentrated copper sulfate solution will mix with the table salt solution very slowly. Therefore, when the cell operates, copper will be released on the copper electrode, and zinc will dissolve in the form of sulfate or chloride in the upper part of the cell.
Nowadays batteries use almost exclusively dry cells, which are more convenient to use. Their ancestor is the Leclanche element. The electrodes are a zinc cylinder and a carbon rod. The electrolyte is a paste that mainly consists of ammonium chloride. Zinc dissolves in the paste, and hydrogen is released on the coal. To avoid polarization, the carbon rod is dipped into a linen bag containing a mixture of coal powder and pyrolusite. The carbon powder increases the electrode surface, and the pyrolusite acts as a depolarizer, slowly oxidizing the hydrogen.
True, the depolarizing ability of pyrolusite is weaker than that of the previously mentioned potassium dichromate. Therefore, when current is received in dry cells, the voltage drops quickly, they " get tired"due to polarization. Only after some time does the oxidation of hydrogen occur with pyrolusite. Thus, the elements " resting", if you do not pass current for some time. Let's check this on a flashlight battery, to which we connect a light bulb. In parallel with the lamp, that is, directly to the terminals, we connect a voltmeter.
At first, the voltage will be about 4.5 V. (Most often, such batteries have three cells connected in series, each with a theoretical voltage of 1.48 V.) After some time, the voltage will drop and the glow of the light bulb will weaken. Based on the voltmeter readings, we can judge how long the battery needs to rest.
A special place is occupied by regenerating elements known as batteries. They undergo reversible reactions and can be recharged after the cell has been discharged by connecting to an external DC source.
Currently, lead-acid batteries are the most common; The electrolyte in them is dilute sulfuric acid, into which two lead plates are immersed. The positive electrode is coated with lead dioxide PbO 2, the negative is metallic lead. The voltage at the terminals is approximately 2.1 V. When discharging, lead sulfate is formed on both plates, which again turns into metallic lead and lead peroxide when charging.

APPLICATION OF GALVANIC COATINGS

Metal is deposited by current

Having bent the ends of two thin sheet copper plates, we hang them on opposite walls of a beaker or, better yet, a small glass aquarium. We attach the wires to the plates with terminals.
Electrolyte Let's prepare according to the following recipe: 125 g of crystalline copper sulfate, 50 g of concentrated sulfuric acid and 50 g of alcohol (denatured alcohol), the rest is water up to 1 liter. To do this, first dissolve copper sulfate in 500 ml of water, then carefully add sulfuric acid in small portions ( Heating! Liquid may splash!), then add alcohol and add water to a volume of 1 liter.
Fill the coulometer with the prepared solution and connect a variable resistance, an ammeter and a lead battery to the circuit. Using resistance, we adjust the current so that its density is 0.02-0.01 A/cm 2 of the electrode surface. If the copper plate has an area of ​​50 cm2, then the current strength should be in the range of 0.5-1 A.
After some time, light red metallic copper will begin to precipitate at the cathode (negative electrode), and copper will go into solution at the anode (positive electrode). To clean the copper plates, we will pass current through the coulometer for about half an hour. Then we take out the cathode, carefully dry it with filter paper and weigh it accurately. Let's install an electrode in the cell, close the circuit using a rheostat and maintain a constant current, for example 1 A. After an hour, open the circuit and weigh the dried cathode again. At a current of 1 A, its mass will increase by 1.18 g per hour of operation.
Therefore, an amount of electricity equal to 1 ampere hour passing through a solution can release 1.18 g of copper. Or in general: the amount of substance released is directly proportional to the amount of electricity passing through the solution.
To isolate 1 equivalent of an ion, it is necessary to pass through the solution an amount of electricity equal to the product of the electrode charge e and Avogadro's number N A:
e*N A = 1.6021 * 10 -19 * 6.0225 * 10 23 = 9.65 * 10 4 A * s * mol -1 This value is indicated by the symbol F and is named after the discoverer of the quantitative laws of electrolysis Faraday number(exact value F- 96,498 A*s*mol -1). Therefore, to isolate a given number of equivalents from a solution n e an amount of electricity should be passed through the solution equal to F*n e A*s*mol -1 . In other words,
I*t =F*n uh Here I- current, t- time of passage of current through the solution. In the section " Titration Basics"It has already been shown that the number of equivalents of a substance n e is equal to the product of the number of moles and the equivalent number:
n e = n*Z Hence:

I*t = F*n*Z

In this case Z- ion charge (for Ag + Z= 1, for Cu 2+ Z= 2, for Al 3+ Z= 3, etc.). If we express the number of moles as the ratio of mass to molar mass ( n = m/M), then we get a formula that allows us to calculate all the processes occurring during electrolysis:

I*t =F*m*Z/M

Using this formula you can calculate the current:

I = F*m*Z/(t*M)= 9.65*10 4 *1.18*2 / (3600*63.54) A*s*g*mol/(s*mol*g) = 0.996 A

If we introduce the relation for electrical work W el

W el = U*I*t And W email/ U = I*t

Then, knowing the tension U, you can calculate:

W el = F*m*Z*U/M

It is also possible to calculate how long it takes for a certain amount of a substance to be electrolytically released, or how much of a substance will be released in a certain time. During the experiment, the current density must be maintained within specified limits. If it is less than 0.01 A/cm2, then too little metal will be released, since copper(I) ions will be partially formed. If the current density is too high, the adhesion of the coating to the electrode will be weak and when the electrode is removed from the solution, it may crumble.
In practice, galvanic coatings on metals are used primarily to protect against corrosion and to obtain a mirror-like shine.
In addition, metals, especially copper and lead, are purified by anodic dissolution and subsequent separation at the cathode (electrolytic refining).
To plate iron with copper or nickel, you must first thoroughly clean the surface of the object. To do this, polish it with washed chalk and successively degrease it with a diluted solution of caustic soda, water and alcohol. If the item is covered with rust, you need to pickle it in advance in a 10-15% solution of sulfuric acid.
We hang the cleaned product in an electrolytic bath (a small aquarium or a beaker), where it will serve as a cathode.
The solution for applying copper coating contains 250 g of copper sulfate and 80-100 g of concentrated sulfuric acid in 1 liter of water (Caution!). In this case, the copper plate will serve as the anode. The surface of the anode should be approximately equal to the surface of the object being coated. Therefore, you must always ensure that the copper anode hangs in the bath at the same depth as the cathode.
The process will be carried out at a voltage of 3-4 V (two batteries) and a current density of 0.02-0.4 A/cm 2. The temperature of the solution in the bath should be 18-25 °C.
Let us pay attention to the fact that the anode plane and the surface to be coated are parallel to each other. It is better not to use objects with complex shapes. By varying the duration of electrolysis, it is possible to obtain copper coatings of different thicknesses.
Often they resort to preliminary copper plating in order to apply a durable coating of another metal to this layer. This is especially often used for chrome plating of iron, nickel plating of zinc casting and in other cases. True, very poisonous cyanide electrolytes are used for this purpose.
To prepare an electrolyte for nickel plating, dissolve 25 g of crystalline nickel sulfate, 10 g of boric acid or 10 g of sodium citrate in 450 ml of water. You can prepare sodium citrate yourself by neutralizing a solution of 10 g of citric acid with a dilute solution of sodium hydroxide or soda solution. Let the anode be a nickel plate of the largest possible area, and take the battery as a voltage source.
Using a variable resistance, we will maintain the current density equal to 0.005 A/cm 2 . For example, with an object surface of 20 cm2, you need to work at a current strength of 0.1 A. After half an hour of work, the object will already be nickel-plated. Let's take it out of the bath and wipe it with a cloth. However, it is better not to interrupt the nickel plating process, since then the nickel layer may become passivated and the subsequent nickel coating will not adhere well.
To achieve a mirror shine without mechanical polishing, we introduce a so-called shine-forming additive into the galvanic bath. Such additives include, for example, glue, gelatin, sugar. You can add, for example, a few grams of sugar to a nickel bath and study its effect.
To prepare an electrolyte for chromium plating of iron (after preliminary copper plating), dissolve 40 g of chromic acid anhydride CrO 3 (Caution! Poison!) and exactly 0.5 g of sulfuric acid (in no case more!) in 100 ml of water. The process occurs at a current density of about 0.1 A/cm 2, and a lead plate is used as the anode, the area of ​​which should be slightly less than the area of ​​the chrome-plated surface.
Nickel and chrome baths are best heated slightly (to about 35 ° C). Please note that electrolytes for chrome plating, especially during a long process and high current, emit vapors containing chromic acid, which are very harmful to health. Therefore, chrome plating should be carried out under traction or in the open air, for example on a balcony.
When chrome plating (and to a lesser extent, nickel plating), not all of the current is used for metal deposition. At the same time, hydrogen is released. Based on a number of voltages, it would be expected that metals in front of hydrogen should not be released from aqueous solutions at all, but on the contrary, less active hydrogen should be released. However, here, as with the anodic dissolution of metals, the cathodic evolution of hydrogen is often inhibited and is observed only at high voltage. This phenomenon is called hydrogen overvoltage, and it is especially large, for example, on lead. Thanks to this circumstance, a lead-acid battery can function. When charging a battery, instead of PbO 2, hydrogen should appear at the cathode, but, due to overvoltage, the evolution of hydrogen begins when the battery is almost fully charged.

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. Oxide films are formed on the surface of these metals, which have a 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 different metal ions 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 a comprehensive Characteristic of the chemical properties of metals. A series of voltages characterizes only the redox ability of the Electrochemical system “metal - metal ion” under strictly defined conditions: the values ​​​​given in it refer to an aqueous solution, temperature and unit concentration (activity) of metal ions.

In an electrochemical cell (galvanic cell), the electrons remaining after the formation of ions are removed through a metal wire and recombine with ions of another type. That is, the charge in the external circuit is transferred by electrons, and inside the cell, through the electrolyte in which the metal electrodes are immersed, by ions. This creates a closed electrical circuit.

The potential difference measured in an electrochemical cell is o is explained by the difference in the ability of each metal to donate electrons. Each electrode has its own potential, each electrode-electrolyte system is a half-cell, and any two half-cells form an electrochemical cell. The potential of one electrode is called the half-cell potential, and it determines the ability of the electrode to donate electrons. It is obvious that the potential of each half-element does not depend on the presence of another half-element and its potential. The half-cell potential is determined by the concentration of ions in the electrolyte and temperature.

Hydrogen was chosen as the “zero” half-element, i.e. it is believed that no work is done for it when an electron is added or removed to form an ion. The “zero” potential value is necessary to understand the relative ability of each of the two half-cells of the cell to give and accept electrons.

Half-cell potentials measured relative to a hydrogen electrode are called the hydrogen scale. If the thermodynamic tendency to donate electrons in one half of the electrochemical cell is higher than in the other, then the potential of the first half-cell is higher than the potential of the second. Under the influence of the potential difference, electron flow will occur. When two metals are combined, it is possible to determine the potential difference that arises between them and the direction of electron flow.

An electropositive metal has a higher ability to accept electrons, so it will be cathodic or noble. On the other side are electronegative metals, which are capable of spontaneously donating electrons. These metals are reactive and therefore anodic:

- → 0 → +

Al Mn Zn Fe Sn Pb H 2 Cu Ag Au

For example Cu gives up electrons more easily Ag, but worse than Fe . In the presence of a copper electrode, silver nonions will begin to combine with electrons, resulting in the formation of copper ions and the precipitation of metallic silver:

2 Ag + + Cu → Cu 2+ + 2 Ag

However, the same copper is less reactive than iron. When metallic iron comes into contact with copper nonates, it will precipitate and the iron will go into solution:

Fe + Cu 2+ → Fe 2+ + Cu.

We can say that copper is a cathode metal relative to iron and an anodic metal relative to silver.

The standard electrode potential is considered to be the potential of a half-cell of fully annealed pure metal as an electrode in contact with ions at 25 0 C. In these measurements, the hydrogen electrode acts as a reference electrode. In the case of a divalent metal, we can write down the reaction occurring in the corresponding electrochemical cell:

M + 2H +→ M 2+ + H 2.

If we arrange metals in descending order of their standard electrode potentials, we obtain the so-called electrochemical series of metal voltages (Table 1).

Table 1. Electrochemical series of metal voltages

For example, in a copper-zinc galvanic cell, there is a flow of electrons from zinc to copper. The copper electrode is the positive pole in this circuit, and the zinc electrode is the negative pole. The more reactive zinc loses electrons:

Zn → Zn 2+ + 2е - ; E °=+0.763 V.

Copper is less reactive and accepts electrons from zinc:

Cu 2+ + 2e - → Cu; E °=+0.337 V.

The voltage on the metal wire connecting the electrodes will be:

0.763 V + 0.337 V = 1.1 V.

Table 2. Stationary potentials of some metals and alloys in sea water in relation to a normal hydrogen electrode (GOST 9.005-72).

Note . The indicated numerical values ​​of potentials and the order of metals in a series can vary to varying degrees depending on the purity of the metals, the composition of sea water, the degree of aeration and the state of the surface of the metals.

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. So, when the 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 is converted 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:

2РbSO 4 + 2Н 2 O → Рb + РbO 2 + 4Н + + 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.

A feature of electrochemical processes, in contrast to conventional chemical ones, is the spatial separation of oxidation and reduction processes. These processes, which cannot occur without each other, constitute the overall 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 the metal to the solution V 1 may be greater than the rate of reverse transition of ions from the solution to the 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 higher concentrations correspond to a higher velocity 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 the metal and the 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 will obtain a series of voltages known from a general chemistry course. 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 electric 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. In general, the chemical reaction that occurs in a copper-zinc cell 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.

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