Standard potentials of redox reactions. A range of standard electrode potentials

The potential jump at the metal-solution interface, as well as the potential difference between two points in different phases, cannot be measured experimentally. Since only the magnitude of the EMF of an electrochemical circuit can be measured experimentally, only the relative values ​​of the so-called electrode potentials can be experimentally determined, i.e. The EMF of a circuit composed of a given electrode and some standard electrode, the potential of which is conventionally assumed to be zero. Such a standard electrode, or reference electrode, is a reversible hydrogen electrode - a glass vessel filled with a strong acid solution (HCl or H 2 SO 4) with a concentration of hydrogen ions [H + ] = 1 mol / l, in which a platinum plate coated with platinum black (powdered platinum deposited on its surface), capable of adsorbing the supplied gaseous hydrogen at a pressure of 1 atm (Fig. 4).

This electrode corresponds to a reversible process, which can be written as

2H + +2ē ↔ N 2 ,

By connecting another half-cell to a hydrogen electrode in a galvanic cell, it is possible to determine the EMF of this galvanic cell, and from it the relative standard electrode potential of a given galvanic couple. For example, in a galvanic cell Zn 0 /Zn 2+ //2H + /H 2, the EMF determined by a voltmeter is 0.76V (see Fig. 5).

The “+” sign of the electrode potential corresponds to the movement of ions from the solution to the electrode in the element, where the electrode in question is connected to the hydrogen electrode, and the movement of electrons along the external circuit from the hydrogen electrode. The “–” sign is placed in front of the value of the electrode potential when ions and electrons move in the opposite direction.

Since in our example, an increase in the concentration of Zn 2+ ions and a decrease in the concentration of H + ions were experimentally established, the values ​​of the electrode potential of the zinc electrode should be given with the “–” sign.

With respect to a standard hydrogen electrode, it is possible to determine the potentials not only of Me/Me n+ pairs, but also of pairs composed of any reducing agent and its oxidized form and any oxidizing agent and its reduced form.

5.4. Redox potentials

Let us consider such electrodes, the reactions on which are not associated with the release of simple substances from the electrolyte or the dissolution of simple substances in it, but are associated with a change in the valence of ions in solution. A chemical reaction accompanied by the transfer of electrons between the molecules of two substances participating in the reaction can be written in the following form:

For example: Oxidized 1 + n 1 ē↔ Restore 1 - restored form;

Restore 2 - n 2 ē ↔ Oxid. 2 - oxidized form.

Therefore, we should not talk about a separate oxidizing agent and reducing agent, but about redox systems, the components of which are the oxidized and reduced forms of the same compound.

The value of the redox potential (ORP) must be indicated for a pair: oxidized and reduced forms. It is denoted by φ, V (Volt) - φ oxidized form / reduced form. The numerator of the index is the oxidized form, and the denominator is the reduced form.

For example, one usually writes
;
;

ORP is a value that characterizes the redox ability of substances.

When experimentally determining the relative values ​​of the redox potential of various pairs, it should be taken into account that their value depends not only on the strength of the oxidizing agent and reducing agent included in a given pair, but also on the ratio of their concentrations (activities). To obtain comparable results, it is necessary to make their concentrations the same, for example, equal to 1 mol/l or 1 g-ion/l, and combine different redox pairs with the same standard pair (standard hydrogen electrode, which is a pair of 2H + / H 2 at an H + concentration of 1 g-ion/l) (see Figs. 4 and 6).

Any oxidizing agent, by gaining electrons, passes into its reducing form, and the reducing agent, by donating electrons, passes into the oxidized form. For example:

Fe 3+ + ē = Fe 2+ - restored form;

2 H + + 2ē =H 2 - oxidized form.

The negative pole of such an element is a standard hydrogen electrode, the positive pole is a platinum electrode.

At the first stage, the process of donating electrons by hydrogen molecules to platinum takes place, i.e. the reaction of their oxidation to hydrogen cations:

H 2 – 2 ē ↔2H +

The electrons released in this case flow through the conductor to the platinum electrode, where they are joined by Fe 3+ ions, which are reduced to Fe 2+:

2Fe 3+ + 2 ē ↔ 2Fe 2+

Adding term by term both written equations, we obtain the general equation of the reaction that occurs during the operation of this element:

2 Fe 3+ +H 2 ↔ 2 Fe 2+ + 2H +

The EMF of this element turns out to be 0.77V, because. it is the difference between the standard potentials of both pairs can be written:

EMF=
=0.77V;

because the value φ 0 / 2Н + /Н2 is conditionally taken as 0, then
= +0.77V.

The plus sign shows that this pair, when combined with a standard hydrogen electrode, plays the role of a positive pole and the result obtained for this pair
the value of the standard potential (+0.77V) is a measure of the ability of Fe 3+ ions to take electrons from the H 2 molecule, i.e. oxidize them to H + ions.

The higher the standard redox potential of a given pair, the stronger the oxidizing agent is its oxidized form and the weaker reducing agent is the reduced form.

When any two redox pairs are combined, the stronger of the two oxidizing agents takes electrons from the stronger reducing agent, and a weaker reducing agent and oxidizing agent are formed.

φ values ​​for various redox systems measured under standard conditions
(temperature 298 K, pressure 101.3 kPa, concentrations of oxidized and reduced forms equal to one: 1 mol/l or 1 g-ion/l) are given in reference tables (see Appendix 3).

The direction of the redox reaction is such that a weaker oxidizing agent and reducing agent are obtained from a stronger oxidizing agent and a reducing agent. Relationship between quantities And
is expressed by the Nernst formula

or (5.1)

(5.2)

where T is the absolute temperature (273+t°), K;

F – Faraday number – 96485 cells/mol;

R – gas constant – 8.31 J/(mol K);

n is the number of electrons accepted by the oxidizing agent or donated by the reducing agent;

a Ox is the active concentration of the oxidizing agent;

a Red is the active concentration of the reducing agent;

a and b are the coefficients in front of the oxidizer and reductant.

If we substitute the values ​​of R, F into formula (5.2), taking into account that for dilute solutions the activities of ions are approximately equal to their concentrations, then for 25 ° C the Nernst equation will have the following form:

, (5.3),

where and are the concentrations of the oxidizing agent and reducing agent, mol/l.

If hydrogen ions H + take part in the reaction, then their concentration affects the ORP value:

, (5.4)

where c is the coefficient before H + in the OVR ion-molecular equation.

For example:

(5.7)

The smaller the value of the redox potential, the stronger the reducing properties are characterized by the reduced form of the redox system and the weaker oxidizing properties are characterized by the oxidized form. Conversely, the more positive the value of the redox potential, the stronger the oxidizing properties are characterized by the oxidized form and the weaker reducing properties are exhibited by the reduced form in the OVR.

For example, when comparing the standard values ​​of the following ORP (systems):

And
Let us determine between which components of these systems a reaction can occur. Since the value
>
,That Fe 3+ will exhibit stronger oxidizing properties than WITHu 2+ , A Cu 0 - stronger restorative properties than Fe 2+ . Hence Cu 0 And Fe 3+ may react in the following way. Let us compose the OVR molecular equation based on the scheme, for this, positively charged ions must be combined with negatively charged ones, so that the desired neutral compound is obtained. As you can see, there are no negatively charged ions in the scheme itself, it is necessary to think about which anions can be used. The choice is made for the following reasons: the substance obtained by combining ions must be stable and soluble. For the scheme under consideration, such ions can be chloride or sulfate ions. Chloride ions are the most convenient. There are no more cations on the left side of the diagram, so no other anions are needed. The same anions must be present in the reaction products, therefore we will combine the cations of the right side with chloride ions:, Cu 2+ ions oxidize ions Fe 2+ they cannot, i.e. the reverse direction of this reaction is impossible.

High potential oxidizers are capable of oxidizing any of the lower potential reducing agents. Yes, ion
in an acidic environment, having a can oxidize reducing agents:

To predict the direction of the OVR, you need to find
(or
reactions).

If
(or
) is greater than zero, the reaction proceeds from left to right.

For a galvanic cell, the following form of writing is accepted (for example, Daniel's cell):

Zn | ZnSO 4 || CuSO4 | Cu,

where is the vertical line | denotes the phase boundary, and the double vertical line || - salt bridge. The electrode at which oxidation occurs is called anode; the electrode at which reduction occurs is called cathode. It is customary to write a galvanic cell so that the anode is on the left.

Electrode half-reactions are usually written as reduction reactions (Table 12.1), so the total reaction in the galvanic cell is written as the difference between the reactions on the right and left electrodes:

Right electrode: Cu 2+ + 2e = Cu

Left electrode: Zn 2+ + 2e = Zn

General reaction: Cu 2+ + Zn = Cu + Zn 2+

Potential E electrode is calculated by Nernst formula:

Where a Ox and a Red - activities of the oxidized and reduced forms of the substance participating in the half-reaction; E o- standard capacity electrode (at a Ox = a Red=1); n- the number of electrons involved in the half-reaction; R- gas constant; T- absolute temperature; F is the Faraday constant. At 25°C

The standard electrode potentials of the electrodes are measured relative to the standard hydrogen electrode, the potential of which is assumed to be zero. The values ​​of some standard electrode potentials are given in Table 12.1.

Electromotive force ( EMF) element is equal to the potential difference of the right and left electrodes:

E= E P - E L.

If the EMF of the element is positive, then the reaction (as it is written in the element) proceeds spontaneously. If the EMF is negative, then the reverse reaction occurs spontaneously.

Standard EMF is equal to the difference of standard potentials:

For the Daniell element, the standard emf is

E o = E o(Cu2+ /Cu) - E o(Zn 2+ / Zn) \u003d +0.337 - (-0.763) \u003d +1.100 V.

The emf of the element is related to G reaction occurring in the element:

G = - nFE.

.

The equilibrium constant of the reaction occurring in the Daniell element is equal to

= 1.54 . 10 37 .

Knowing EMF temperature coefficient, other thermodynamic functions can be found:

H = G + T S = - nFE + .

Table 12.1. Standard electrode potentials at 25 o C.

(More details can be found in
base on redox potentials

Electrode	Electrode reaction
Li+/Li	Li + + e = Li	-3.045
K + /K	K + + e = K	-2.925
Ba2+ /Ba	Ba 2+ + 2e = Ba	-2.906
Ca2+ /Ca	Ca 2+ + 2e = Ca	-2.866
Na+/Na	Na + + e = Na	-2.714
La3+ /La	La 3+ + 3e = La	-2.522
Mg2+ /Mg	Mg 2+ + 2e = Mg	-2.363
Be 2+ /Be	Be 2+ + 2e = Be	-1.847
A1 3+ /A1	Al 3+ + 3e = Al	-1.662
Ti2+ /Ti	Ti 2+ + 2e = Ti	-1.628
Zr 4+ /Zr	Zr 4+ + 4e = Zr	-1.529
V2+ /V	V 2+ + 2e = V	-1.186
Mn2+ /Mn	Mn 2+ + 2e = Mn	-1.180
WO 4 2- /W	WO 4 2- + 4H 2 O + 6e \u003d W + 8OH -	-1.05
Se2- /Se	Se + 2e = Se 2-	-0.77
Zn2+ /Zn	Zn 2+ + 2e = Zn	-0.763
Cr 3+ /Cr	Cr 3+ + 3e = Cr	-0.744
Ga 3+ /Ga	Ga 3+ + 3e = Ga	-0.529
S2-/S	S + 2e = S 2-	-0.51
Fe 2+ /Fe	Fe 2+ + 2e = Fe	-0.440
Cr 3+ ,Cr 2+ /Pt	Cr 3+ + e = Cr 2+	-0.408
Cd 2+ /Cd	Cd 2+ + 2e = Cd	-0.403
Ti 3+ , Ti 2+ /Pt	Ti 3+ + e = Ti 2+	-0.369
Tl + /Tl	Tl + + e = Tl	-0.3363
Co2+ /Co	Co 2+ + 2e = Co	-0.277
Ni2+ /Ni	Ni 2+ + 2e = Ni	-0.250
Mo3+ /Mo	Mo 3+ + 3e = Mo	-0.20
Sn 2+ /Sn	Sn 2+ + 2e = Sn	-0.136
Pb 2+ /Pb	Pb 2+ + 2e = Pb	-0.126
Ti 4+ , ​​Ti 3+ /Pt	Ti 4+ +e = Ti 3+	-0.04
D + /D 2 , Pt	D + + e \u003d 1 / 2 D 2	-0.0034
H + /H 2 , Pt	H + + e \u003d 1 / 2 H 2	0.000
Ge2+ /Ge	Ge 2+ + 2e = Ge	+0.01
Br - /AgBr/Ag	AgBr + e = Ag + Br -	+0.0732
Sn 4+ , ​​Sn 2+ /Pt	Sn4+ + 2e = Sn2+	+0.15
Cu 2+ , Cu + /Pt	Cu 2+ + e \u003d Cu +	+0.153
Cu 2+ /Cu	Cu 2+ + 2e = Cu	+0.337
Fe(CN) 6 4- , Fe(CN) 6 3- /Pt	Fe(CN) 6 3- + e = Fe(CN) 6 4-	+0.36
OH - /O 2, Pt	l / 2 O 2 + H 2 O + 2e \u003d 2OH -	+0.401
Cu+/Cu	Cu + + e = Cu	+0.521
J - /J 2 , Pt	J 2 + 2e = 2J -	+0.5355
Te4+ /Te	Te 4+ + 4e = Te	+0.56
MnO 4 - , MnO 4 2- /Pt	MnO 4 - + e \u003d MnO 4 2-	+0.564
Rh2+ /Rh	Rh2+ /Rh	+0.60
Fe 3+ , Fe 2+ /Pt	Fe 3+ + e \u003d Fe 2+	+0.771
Hg 2 2+ /Hg	Hg 2 2+ + 2e = 2Hg	+0.788
Ag+/Ag	Ag + + e = Ag	+0.7991
Hg2+ /Hg	Hg 2+ + 2e = Hg	+0.854
Hg 2+ , Hg + /Pt	Hg 2+ + e = Hg +	+0.91
Pd2+ /Pd	Pd 2+ + 2e = Pd	+0.987
Br - /Br 2 , Pt	Br 2 + 2e \u003d 2Br -	+1.0652
Pt 2+ /Pt	Pt 2+ + 2e = Pt	+1.2
Mn 2+ , H + /MnO 2 , Pt	MnO 2 + 4H + + 2e \u003d Mn 2+ + 2H 2 O	+1.23
Cr 3+ , Cr 2 O 7 2- , H + /Pt	Cr 2 O 7 2- + 14H + + 6e = 2Cr 3+ + 7H 2 O	+1.33
Tl 3+ , Tl + /Pt	Tl 3+ + 2e = Tl +	+1.25
Cl - /Cl 2 , Pt	Cl 2 + 2e \u003d 2Cl -	+1.3595
Pb 2+ , H + /PbO 2 , Pt	PbO 2 + 4H + + 2e \u003d Pb 2+ + 2H 2 O	+1.455
Au 3+ /Au	Au 3+ + 3e = Au	+1.498
MnO 4 - , H + /MnO 2 , Pt	MnO 4 - + 4H + + 3e \u003d MnO 2 + 2H 2 O	+1.695
Ce 4+ , ​​Ce 3+ /Pt	Ce4+ + e = Ce3+	+1.61
SO 4 2-, H + / PbSO 4, PbO 2, Pb	PbO 2 + SO 4 2- + 4H + + 2e = PbSO 4 + 2H 2 O	+1.682
Au+/Au	Au + + e = Au	+1.691
H - /H 2 , Pt	H 2 + 2e \u003d 2H -	+2.2
F - /F 2 , Pt	F 2 + 2e \u003d 2F -	+2.87

Cu 2+ + 2e = Cu Go= -nFEo\u003d -2 (96485 C. mol -1) (+0.337 V) \u003d -65031 J. mol -1.

Cu + + e = Cu Go= -nFEo\u003d - (96485 C. mol -1) (+0.521 V) \u003d -50269 J. mol -1.

Subtracting, we get:

Cu 2+ + e \u003d Cu + Go= -nFEo\u003d -3 (96485 C. mol -1) E o\u003d -14762 J. mol -1,

where E o= +0.153 V.

Example 12-2. Draw a diagram of the galvanic cell in which the reaction takes place.

Ag | AgBr| Br - || Ag + | Ag

Right electrode: Ag + + e = Ag E o= 0.7792 V

Left electrode: AgBr + e = Ag + Br - E o= 0.0732 V

General reaction: Ag + + Br - = AgBr E o= 0.7260 V

Go= -nFEo\u003d - (96485 C. mol -1) (0.7260 V) \u003d -70.05 kJ. mol -1

= 1.872 . 10 12

1/K= a(ag +) . a(Br-)= m(ag +) . m(Br -) . () 2 = m 2 () 2

Hence, setting = 1, we obtain m= 7.31 . 10 -7 mol. kg -1

Example 12-3. H reaction Pb + Hg 2 Cl 2 \u003d PbCl 2 + 2Hg occurring in a galvanic cell is -94.2 kJ. mol -1 at 298.2 K. The EMF of this element increases by 1.45 . 10 -4 V when the temperature rises by 1K. Calculate the EMF of the element and S at 298.2 K.

2. 96485 . 1.45. 10 -4 \u003d 28.0 (J. mol -1. K -1).

G = H - T S = -nFE, where

1. In an acidic environment there should be no ions on either the left or right side Equalization is carried out due to ions and water molecules.

2. In an alkaline environment neither on the left nor on the right side should there be ions. Equalization is carried out due to ions and water molecules.

3. In a neutral environment there should not be any ions or on the left side. However, they may appear among the reaction products on the right side.

4. Consider how the proposed schemes work on specific examples.

5. Task. Complete the equation for the reaction between potassium bichromate and hydrochloric acid.

6. The ion contains chromium in its highest oxidation state, therefore, it can only act as an oxidizing agent. According to the scheme, we compose a half-reaction, given that the medium is acidic (HCl).
Recovery half-reaction:

7. Ions can only be oxidized, because Chlorine has the lowest oxidation state. We compose the oxidation half-reaction:

9. We summarize first the left and then the right parts of the half-reactions, not forgetting first multiply multiplier for the coefficient, if it is in front of the formula.

11. Got a reduced ionic equation.

12. Add the missing cations or anions, taking into account that the number of ions added to the right and left sides of the ionic equation should be the same.

13. In this case, the source of ions ─ was salt, therefore, with each mole, 2 moles of ions enter the solution. They do not take part in the reaction, therefore, they must go unchanged to the right side of the equation. Together with 14 mol of ions, 14 mol of ions are introduced into the solution. Of these, 6 participates in the reaction as a reducing agent, and the remaining 8, like ions, remain unchanged after the reaction, i.e. added to the right side.

14. As a result, we get:

16. After that, you can combine the ions into formulas of real substances:

40. Quantitative characteristics of redox transitions. Electrode potentials of metals. Galvanic cell. Hydrogen electrode and hydrogen potential zero. Standard conditions and standard half-reaction potential. Tables of standard reduction potentials. Using tabular data to assess the possibility of OVR.

Electrode potentials- the difference in electrical potential between the electrode and the electrolyte in contact with it.

The occurrence of the electrode potential is due to the transfer of charged particles through the phase boundary, spec. ion adsorption. The magnitude of the electrode potential in an uneven state depends on the nature and composition of the contacting phases.

The electrode potential is a constant value at a given temperature if a metal plate is immersed in a solution of its salt with the activity of metal ions. This potential is called standard electrode potential.

Galvanic cell- a chemical source of electric current based on the interaction of two metals and / or their oxides in an electrolyte, leading to the appearance of an electric current in a closed circuit. Named after Luigi Galvani. The conversion of chemical energy into electrical energy occurs in galvanic cells.

Standard hydrogen electrode- an electrode used as a reference electrode in various electrochemical measurements and in galvanic cells. A hydrogen electrode (HE) is a metal plate or wire that absorbs gaseous hydrogen well (usually platinum or palladium is used), saturated with hydrogen (at atmospheric pressure) and immersed in an aqueous solution containing hydrogen ions. The potential of the plate depends on the concentration of H + ions in the solution. The electrode is a standard against which the electrode potential of the determined chemical reaction is measured. At a hydrogen pressure of 1 atm, a proton concentration in the solution of 1 mol/l, and a temperature of 298 K, the SE potential is assumed to be 0 V. When assembling a galvanic cell from the SE and the electrode to be determined, the reaction proceeds reversibly on the platinum surface:

2Н + + 2e − = H 2

that is, either hydrogen reduction or its oxidation occurs - this depends on the potential of the reaction taking place on the electrode being determined. By measuring the EMF of a galvanic electrode under standard conditions (see above), the standard electrode potential of the chemical reaction being determined is determined.

SE is used to measure the standard electrode potential of an electrochemical reaction, to measure the concentration (activity) of hydrogen ions, as well as any other ions. VE is also used to determine the solubility product, to determine the rate constants of some electrochemical reactions.

Scheme of a standard hydrogen electrode:

1. Platinum electrode.

2. Hydrogen gas supplied.

3. Acid solution (usually HCl), in which the concentration of H + = 1 mol/l.

4. A water seal that prevents the ingress of oxygen from the air.

5. An electrolytic bridge (consisting of a concentrated solution of KCl) that allows you to connect the second half of the galvanic cell.

The normal electrode potential makes it possible to evaluate the thermodynamic activity of various chemicals, but at present there are no methods to measure its absolute value. In this regard, the electrodes are characterized by the so-called standard electrode potential, which is (according to Nernst's proposal) the difference between the normal potentials of the considered and standard hydrogen electrodes, determined at 25 ° C (298 K). With this approach, the standard electrode potential of hydrogen is conditionally taken equal to zero. Then the standard potential of the substance, the electrode potential of which under the specified conditions is more negative than the potential of the standard hydrogen electrode, is considered negative. If the electrode potential of the substance is less negative than the potential of the standard hydrogen electrode, the standard potential of the substance is considered positive.

Electrochemical activity series of metals (voltage range, a range of standard electrode potentials) - the sequence in which the metals are arranged in order of increasing their standard electrochemical potentials φ 0 corresponding to the reduction half-reaction of the metal cation Me n+ : Me n+ + nē → Me

A number of stresses characterize the comparative activity of metals in redox reactions in aqueous solutions.

A number of voltages are used in practice for a comparative [relative] assessment of the chemical activity of metals in reactions with aqueous solutions of salts and acids and for assessing cathodic and anodic processes during electrolysis:

The metals to the left are stronger reducing agents than the metals to the right: they displace the latter from salt solutions. For example, the interaction Zn + Cu 2+ → Zn 2+ + Cu is possible only in the forward direction.

Metals to the left of hydrogen in the row displace hydrogen when interacting with aqueous solutions of non-oxidizing acids; the most active metals (up to and including aluminum) - and when interacting with water.

Metals in the row to the right of hydrogen do not interact with aqueous solutions of non-oxidizing acids under normal conditions.

During electrolysis, metals to the right of hydrogen are released at the cathode; the reduction of metals of moderate activity is accompanied by the release of hydrogen; the most active metals (up to aluminum) cannot be isolated from aqueous solutions of salts under normal conditions.

41. Redox equilibrium in solutions. Nernst equation. Electrolysis. Electrochemical energy sources. Corrosion as an electrochemical process. Electrolysis of solutions and melts. Electrolytic production of metals. Faraday's law. The practical significance of electrolysis.

Electrolysis- the process of separate oxidation and reduction on the electrodes, carried out due to the flow of current from an external source. Anode = oxidation, positively charged, cathode = reduction, negatively charged.

Faraday's Law: the mass of the substance released during electrolysis is directly proportional to the amount of electricity passed through the solution. Equal amounts of electricity contribute to the release of equivalent masses from various chemical compounds.

m=(M*I*t)/(n*F)

The practical significance of electrolysis

The phenomenon of electrolysis is widely used in modern industry. In particular, electrolysis is one of the methods for the industrial production of aluminum, hydrogen, as well as chlorine, sodium hydroxide. A large amount of metals is extracted from ores and processed by electrolysis. Also, electrolysis is the main process by which chemical current sources function.

Electrolysis is used in wastewater treatment.

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

In addition, there is a method of calculation that in many cases is simpler, more direct, and sometimes more accurate. For this, the standard electrode potentials of reduction reactions in an aqueous medium are used, published in tables of physicochemical quantities.

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

For example, the standard emf of a Harned cell

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

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

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

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

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

The total reaction of this element is:

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

with reaction

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

with reaction

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

Therefore:

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

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

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

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

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

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

Rice. 7.1.Schematic representation of a galvanic cell

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

Sn 4+ + Sn 2+

Rice. 7.2.Occurrence of electrode potential

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

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

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

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

Rice. 7.3 Standard hydrogen electrode

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

2H + +  H 2 (Pt)

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

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

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

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

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

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

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

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

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

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

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

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

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

Fe 3+ + 3Fe

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

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

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

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

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

2Fe 3+ + 2I -  2Fe 2+ + I 2

=
-
= 0.771 - 0.536 = 0.235V

This reaction can proceed spontaneously in the forward direction.