Average activity of a strong electrolyte
It was tacitly assumed above that the /th component of the solution is an uncharged particle. If charged particles appear in the solution - ions, then between them arise (and prevail), in addition to the existing ones, the forces of electrostatic interaction. This is reflected in their thermodynamic characteristics.
As discussed in 1.2, in the thermodynamic description of the properties of systems containing charged particles, the main thermodynamic characteristic of an ion is electrochemical potential:
The activity of the electrolyte and its individual ions also has features of expression (due to the second of the prohibitions mentioned in 1.2). Let us have in solution 1 mol of a strong binary electrolyte MA of the valence type 1:1, completely dissociating into ions M "and A". Formally, the chemical potential of MA, which is absent in the solution in the form of molecules, will be formed from the electrochemical potentials of the ions:
(z M =z A = 1, but we leave the designation of the charge for the sake of generality of description).
Due to the electrical neutrality of the solution
i.e., the chemical potential of the electrolyte in solution is the sum of the chemical potentials of the ions, which, however, are thermodynamically indeterminate, because it is impossible to experimentally create a standard solution of ions of the same charge sign. But since the total effect of both types of ions is reflected in the properties of electrolyte solutions, they introduce average chemical potential of the electrolyte p ± MA conveying this total action:
The physical meaning of this quantity is the partial value of the Gibbs energy per 1 mole of an ion in a given system, regardless of whether it is a cation or anion. This value is directly related average electrolyte activity i ±MA (synonyms: medium ionic, geometric mean):
It is easy to see that 
The average activity of the electrolyte is equal to the product of the average electrolyte concentration times average activity factor. In this case, depending on the method of expressing the concentration, we obtain:
(MA index at y ± omitted).
The average concentration is related to the set electrolyte concentration and to ionic concentrations in the same way as the average activity is related to the electrolyte activity and ionic activities, for example,
The same reasoning applies to the average activities and average activity coefficients that we used for the activities and activity coefficients of the / "-th component - non-electrolyte (see 2.1.2). Just as in that case, a member of the type RT ny ± serves as an energy measure of the interaction of ions with each other and with solvent molecules, as well as the interaction of solvent molecules with each other, modified by the presence of ions. Also, when the concentration of MA tends to 0, the average activity coefficients for all scales of concentration tend to 1, i.e.
It should be taken into account that in the case of electrolytes, this situation arises in much more dilute solutions than in the case of non-electrolytes.
The state of a hypothetical solution is taken as the standard state of the electrolyte in solution, where the average activities a+ and the average activity coefficients y± are equal to 1; then p M d = Rmd, as in the case of a dissolved non-electrolyte. It is assumed that the partial molar enthalpies, volume, and heat capacities of the electrolyte in this hypothetical solution are the same as in the extremely dilute one, where all y ± = 1.
Exactly medium the activities and activity coefficients of the electrolyte are amenable to experimental determination by various methods, and from them the average chemical potential can be calculated.
Addition2A.
General case - electrolyte type M y + A y
In the general case, one electrolyte type M y + A. in solution, its chemical potential
because the v + z + F(p-v_ | G_| jP
As in the example with an electrolyte (electrolyte, el-te) of the MA type, we introduce the concepts average chemical potential And average activity:
Depending on the method of expressing concentrations, the average activity, ionic (y + , y_) and average (y ±) activity coefficients are related as follows:
If there is only one electrolyte in the solution M y + A y _, and its concentration is given in the molarity scale C or molality T, average concentrations C ± And t ± are expressed in terms of given electrolyte concentrations
(similar for C+).
It can be seen that before t ±(or C ±) multiplier appears
It has the following numerical values for common types of electrolytes:
If the solution contains a mixture of electrolytes with a common ion, the concentrations of this ion are summarized. In this case, the “acting” concentration of an individual ion is calculated taking into account its charge according to the formula m t zf(or C t zf), i.e., according to the same scheme,
as in the expression for ionic strength 
. Let there be
solution of 0.1 mol/kg NaCl and 0.3 mol/kg CaCl 2 . In him t a = 0.1-(-1) 2 Н- +2-0.3-(-1) 2 = 0.7 mol/kg, /?r ±NaC1 = (0.10.7) 1/2 = 0.27 mol/kg;
t Si = 0.3-22 = 1.2; tfi + c a ci 2 \u003d (1-2 0.7) 1/3 \u003d 0.84 mol / kg.
Relationship between mean molality t ± and average molarity C ±:
between the corresponding average activity coefficients:
Here Mj and M 2 are the molecular weights of the solvent and the solute, respectively; p 0 and p are the densities of the solvent and solution. In very dilute solutions, the difference between t ± And C ± can be neglected 
,
In chemical thermodynamics, to characterize the properties of systems that do not contain charged particles, and in which a change in composition occurs as a result of chemical reactions or phase transformations, the fundamental Gibbs equation is used, expressed through the characteristic function:
where is the chemical potential of the i-th reagent, and is an infinitesimal change in the amount of this reagent.
When considering phenomena in heterogeneous systems, it is necessary to take into account which phase the quantities , , , , belong to. Belonging to a certain phase is indicated by a superscript, for example - , . The equilibrium condition in a heterophase system at constant temperature and pressure (T and P - const) is the equality of chemical potentials , where and are the chemical potentials of a neutral substance in two coexisting phases.
If the component is a charged particle, then its state also depends on the magnitude of the electric field. When moving charged particles in phase in an electric field, the mass transfer of the component is associated with charge transfer. The fundamental Gibbs equation in this case should have the form:
where - is the internal potential of any phase, that is, the internal potential of that part of the system in which this particle is located.
The internal potential is the work of transferring a single negative imaginary charge from an infinitely distant point A, located in vacuum, to point B, located inside the conducting phase.
The term "imaginary" suggests that this unit charge reacts only to an external electric field and does not interact with the environment.
Since , where is the charge of the th ion, taking into account the sign of the charge; - Faraday constant, - number of moles i th substance, then after transformations we get:
All derivatives of the Gibbs energy with respect to generalized coordinates have the meaning of generalized forces. Therefore, is a generalized force in the phenomena of the transfer of charged particles in an electric field. By analogy with the chemical potential, for electrochemical systems, the value
called electrochemical potential .
When moving one mole of real charged particles (with charge ) from infinity in vacuum deep into the conducting phase (for example, phase ), the work expended consists of two parts: electrostatic equal and chemical, due to the interaction of real particles with this phase, that is, the chemical potential of the component in phase .
The fundamental Gibbs equation then becomes:
Consider the equilibrium at the phase boundary. Suppose that an electrochemical reaction occurs at the interface (electrode reaction)
where and is the stoichiometric coefficient i-th substance or ion (for the starting substances, the stoichiometric coefficients take negative values, and for the reaction products - positive), z- the total number of electrons involved in the electrode reaction (half-reaction).
When a chemical reaction proceeds, the amounts of individual reactants change in proportion to their stoichiometric coefficients in the reaction equation. Mutual proportionality of the values d n i can be expressed by a set of equations:
Thus, the redistribution of the quantities of all substances in the system can be expressed using a single variable, which is denoted x and called chemical variable. The differential of a chemical variable is determined using any of the above equations:
d n i = n I d x;
Then, taking into account this expression for d G we get
d G = – S d T + V d p + . (3.5)
At constant temperature and constant pressure, the equilibrium condition in the system is the minimum Gibbs energy. This means that for an equilibrium system
This equation summarizes the electrochemical potentials and stoichiometric coefficients of all participants in the electrochemical reaction, regardless of what phase they are in.
The equilibrium on the electrode is characterized by the equality of the electrochemical potentials of the components in all phases. In the case of their inequality, charged particles pass through the phase boundary, which is caused by the tendency of the system to thermodynamic equilibrium. As a result, the balance of electric charges in each phase is disturbed, the metal and the solution acquire an electric charge, and a potential jump occurs at their interface. In other words, between phases of the electrode, a certain potential difference arises, due to the nature of the components that make up the electrode, their concentrations and the values of external thermodynamic parameters.
This potential jump is called the galvanic potential (electrode potential) and is denoted by . The galvanic potential is determined by the difference between the internal potentials of both phases: .
Rice. The occurrence of a potential jump ( galvanic potential) at the phase and phase interface.
Let us consider the mechanism of the potential jump occurrence using the example of the most common electrodes with a metal-solution interface. There are such metals that if they are lowered into water or into a solution, then metal ions pass into a layer of water or solution adjacent to the surface of the metal according to the reaction.
This transition occurs if the electrochemical potential of the metal ion in the crystal is greater than the electrochemical potential of the solvated ion in solution. A metal can be thought of as consisting of positively charged metal ions and relatively free electrons. As the ions go into solution, the state of the system gradually changes. The metal acquires a negative charge, the value of which increases as the electrochemical reaction proceeds. As a result, the electrochemical potential of metal ions on the surface decreases. The number of ions in the solution increases, and their electrochemical potential increases due to the repulsion of like ions. As a result, the rate of transition of ions into the solution decreases, and the rate of the reverse process - the transition of ions from the solution to the metal increases. Finally, a state occurs in which the rates of both processes become the same, that is, equilibrium occurs in the system. In this case, the metal acquires a negative charge, which corresponds to a certain potential, and an excess of cations is formed in the solution, which are held at the surface of the metal electrode as a result of the action of electrostatic forces, and this layer corresponds to its own potential. These potentials are called internal potentials and are denoted by , where the index indicates which phase the potential belongs to. As a result, the so-called electrical double layer , which corresponds to a certain potential difference, called galvanic potential -
(For example, ).
To determine the magnitude of the galvanic potential that occurs at the phase boundary, it is also necessary to experimentally determine the difference in electrochemical potentials in these phases. Since then
It follows from the equation that the measurement of the galvanic potential between points in different phases is possible only if the chemical potentials of substances in different phases are equal, that is, at . In this case, we get:
Hence it follows that it is impossible to experimentally determine the galvanic potential at the interface between two phases of different composition.
The value of the galvanic potential depends on the properties of the phases that form the interface and on the concentration of ions in the solution.
In general, for an electrochemical reaction
flowing at the phase boundary, the equilibrium condition, in accordance with (3.6) will be written:
where and are the charges of the particles of the oxidized and reduced forms, and are the potentials of the phases containing the oxidized and reduced forms of the substance. After transforming the equation, we get:
According to the charge balance,
where is the total charge of the ions participating in the reaction in the phase containing the reduced form of the substance, and is the total charge of the ions participating in the reaction in the phase containing the oxidized form of the substance. For arbitrary liquid and solid solutions, the chemical potential i th component is expressed in terms of its activity by the equation . Given that - galvanic potential, we get:
Since the standard chemical potential of the component is equal to the value of its standard Gibbs energy, we get
j°, which is called
Where j° - standard electrode potential; R is the universal gas constant;
T–temperature, K; F is the Faraday constant; - the number of electrons involved in the electrode process; and - activity of the oxidized and reduced forms.
The resulting equation is called the Nernst equation. Standard electrode potential j° is a value specific to each electrode process, which also depends on temperature and the nature of the solvent. The standard electrode potential is equal to the electrode potential in which the ratio of the activities of all participants in the electrode reaction is equal to one. The Nernst equation relates the magnitude of the potential difference between the phase of the electrolyte solution and the phase of the conductor of the first kind with the activities of the components involved in the electrode reaction.
As an example of the establishment of electrochemical equilibrium, consider the simplest case - equilibrium at the interface of a metal with a solution containing ions of this metal. The following electrochemical reaction will take place on the electrode:
Equilibrium is established as a result of the transition of metal ions from the volume of the solution to the metal and vice versa, provided that equality (3.6) is fulfilled.
Combining all constant values into one value - j°, which is called standard electrode potential, we obtain an expression for the potential difference between the phases that make up the electrode:
The combination of constants R, F, and temperature (RT/F) often occurs in electrochemical equations; it has the dimension of stress. It is generally accepted to denote it as b 0 . Often the Nernst equation is written in terms of decimal logarithms. The transition to decimal logarithms is carried out by multiplying b 0 by ln10 = 2.3 (this product is denoted as b). At 298 K, the values of b 0 and b are respectively equal:
The values of the constant b at other temperatures can be easily calculated.
It should be noted that in the general case, when writing the Nernst equation, only those quantities that can vary remain under the logarithm. Thus, when writing the Nernst equation for different cases, it is necessary to observe several rules related to the application of expression (3.9) for various types of electrodes:
1. The activities of pure components that form a separate phase of constant composition (as a rule, these are solids) are taken equal to unity.
2. The activity of the solvent is taken equal to one.
3. Instead of the activities of gaseous substances, the equation includes the relative partial pressures of these gases over the solution. The pressure is given relative to the standard (1 bar \u003d 10 5 Pa), i.e. this quantity is dimensionless, although numerically it coincides with the partial pressure of the gas, expressed in bars.
The use of partial pressures is valid for the case of not very high pressures (on the order of several bar). In the case of high pressures, the fugacity of the gases must be used.
The chemical potential of the neutral component is a function of the temperature, pressure, and chemical composition of the phase in which it resides. The chemical potential is defined as follows:
where G - Gibbs free energy, A - Helmholtz free energy, U - internal energy, I - enthalpy, S - entropy, V - volume, T - temperature, pressure. In measurements, the difference in chemical potentials in various thermodynamic states is always determined, and never the absolute value of the chemical potential in a given state. However, when tabulating the results, it is convenient to assign a certain value to each thermodynamic state. This can be done by assigning an arbitrary value to the chemical potential in some state and determining its value in another state by comparison with the given standard state.
For example, the chemical potentials of pure elements at and pressure of one atmosphere can be taken equal to zero. As soon as the standard state is precisely established and the values of chemical potentials in other states are tabulated, the experimental results become unambiguous. We will return to this issue again when discussing data on electrochemical cells.
The electrochemical potential of an ion was introduced by Guggenheim, and the difference in its values in two phases was defined as the work on the reversible transfer of one gram ion from one phase to another at constant temperature and volume. It depends on the temperature, pressure, chemical composition and electrical state of the phase. It remains to be seen how well these independent variables are defined. Let us consider the following cases in which ion transport may appear:
1. Constant temperature and pressure, the same chemical composition of the phases. Differences between phases can only be electrical in nature.
a) For the transfer of one gram ion of component i from phase to phase a, the work of transfer is equal to
where the difference between the two phases can be characterized by the difference in the electric potentials of both phases (the second relation).
b) For the transfer of component 1 gram ions and component 2 gram ions, provided that
the work done is zero. Such electrically neutral combinations of ions do not depend on the electrical state of the phase, and this fact can be used to verify the above definition of the potential difference. Since for neutral combinations the total work of transfer will be equal to zero, so that equality (13-3) is satisfied, we have
If we apply equality (13-2) to the ionic component 1, then we can combine equalities (13-2) - (13-4) and express the difference
electrochemical potentials of ionic component 2 in the form
Therefore, the electric potential difference defined by equation (13-2) does not depend on which of the two charged components (1 or 2) is used in equation (13-2). In this sense, the electrical potential difference is defined correctly and coincides with the usual idea of the potential difference.
2. Constant temperature and pressure, different chemical compositions of both phases. When transferring neutral combinations of ions that satisfy equality (13-3), there is no dependence on the electrical state of any of the phases. Thus, the work of transfer will depend only on the difference in chemical compositions. The work of transfer of a charged component will still be given by the equality
but it can no longer be expressed simply in terms of electrical potential differences, since the chemical environment of the transferred component will be different in both phases.
It should be noted that a quantitative characteristic or measure of the difference in the electrical states of two phases with different chemical compositions has not yet been established. It is possible (and even reasonable for some computational purposes) to define such an electrical variable, but this is inevitably connected with an element of arbitrariness and is not essential for the consideration of thermodynamic phenomena. Several different ways of making this determination are discussed in Chap. 3. The usual definition of electric potential is based on electrostatics rather than thermodynamics, so the use of electrochemical potentials is more appropriate here.
Of interest is the question of the state of the phase, as well as whether both phases are in the same state. If two phases have different compositions, then the question of whether they are in the same electrical state is irrelevant from the point of view of thermodynamics. On the other hand, if both phases are chemically identical, then it is convenient to quantitatively describe their electrical states in a way that coincides with the usual definition of potential.
Electrode processes. The concept of potential jumps and electromotive force (EMF). Electrochemical circuits, galvanic cells. Standard hydrogen electrode, standard electrode potential. Classification of electrochemical circuits and electrodes.
9.1 Electrochemical systems. Electrode. Electrochemical potential. Absolute electrode potentials and electromotive force.
The mutual transformation of electrical and chemical forms of energy occurs in electrochemical systems, including:
conductors of the second kind - substances with ionic conductivity (electrolytes).
conductors of the first kind - substances with electronic conductivity.
At the interface between two phases, an electric charge is transferred, i.e. there is a potential jump ().
A system consisting of contacting conductors of the first and second kind is called electrode.
The processes occurring at the phase boundary of conductors of the I and II kind in the electrodes are calledelectrode processes .
The electrode is a system consisting of at least two phases.
Let us consider how a potential jump occurs - the electrode potential - at the interface between the metal and the salt solution of this metal. When a metal plate is immersed in a salt solution, some of the metal ions from the surface of the plate can go into a solution adjacent to the surface of the plate. The metal is charged negatively, and the resulting electrostatic forces prevent the further flow of this process. The system is in equilibrium. The reverse process of transition of metal cations from solution to the plate is also possible. These processes lead to the appearance of a double electric layer and a potential jump.
The direction of the metal ion transfer process is determined by the ratio of the electrochemical potentials of the ions ( 
) in the solution phase and the condensed phase. The process continues until the electrochemical potentials in the two phases are equalized.
The electrochemical potential consists of two terms
=
.
m chem. - chemical potential that characterizes the chemical response to a change in the environment of a given particle.
m el - the electrical component of the electrochemical potential or the potential energy of the electric field, which characterizes the response to the electric field.
For a certain kind of charged particles (i)
, Where
z i is the charge of the ion,
–internal potential, corresponding to the work of transfer of an elementary negative charge from infinity in vacuum deep into the phase.
Equilibrium of an electrochemical system characterized by the equality of electrochemical (rather than chemical) potentials of charged particles in different phases.
In the equilibrium system solution (I) / metal (II), we have:
.
In a non-equilibrium system, the work of transfer of one mol-equiv. ions from phase I to phase II is
.
Since then
In equilibrium, taking into account (1), we have:
,
Where 
– jump at the phase boundary (absolute electrode potential). Denote
,
Where 
- potential jump at the phase boundary at A i = 1
(standard electrode potential).
The standard potential is a value characteristic of a given electrode process. It depends on the temperature and the nature of the electrode. Then for an electrode of type Me Z+ /Me:
. (1)
A potential jump also occurs at the interface between two solutions, this is the diffusion potential 
.
In general terms (for any type of electrodes):
(2)
or for 298K
It should be remembered that if gases are involved in the electrode reaction, then the activity is assumed to be equal to the partial pressure; for the condensed phase of constant composition, A=1.
Equations (1), (2) are called Nernst equations
for the electrode potential. The electric potential difference can be experimentally measured only between two points of the same phase where μ i = const. When an elementary charge moves between two points that are in different phases, in addition to the electric one, work must be performed associated with a change in the chemical environment of the charge. The value of this chemical component of the work cannot be determined, therefore the absolute value of the electrode potential 
impossible to measure. Empirically, it is possible to determine only the magnitude of the EMF of a galvanic cell consisting of two electrodes.
Rules for recording electrodes and electrochemical circuits.
Systems consisting of two or more electrodes, connected in a special way and capable of producing electrical work, that is, serving as a source of electrical energy, are called galvanic cells.
Electromotive force of a galvanic cell(EMF GE) is the sum of jumps in electrode potentials at all phase boundaries in the equilibrium condition (the current in the external circuit is zero).
a) The following recording rules are accepted for electrodes: substances in solution are indicated to the left of the vertical bar, substances forming another phase (gas or solid) are indicated to the right.
If one phase contains several substances, then their characters are separated by commas.
For example,
.
The equation of the electrode reaction for a separate electrode is written in such a way that substances in the oxidized form and electrons are located on the left, and substances in the reduced form are on the right:
,
,
.
b) When recording galvanic cells, an electrode with a more negative potential is located on the left; the solutions of both electrodes are separated from each other by a vertical dotted line if they are in contact with each other, and by two solid lines if there is a salt bridge between the solutions, for example, a saturated KCl solution, with which the diffusion potential is eliminated. Thus, the positively charged electrode is always indicated on the right, and the negatively charged electrode is always indicated on the left.
As an example of an electrochemical circuit, consider a galvanic cell consisting of silver 
and copper 
electrodes. Schematically, the considered element is written in the following form:
where the solid vertical line denotes the metal–solution interface, and the vertical dotted line denotes the solution–solution interface.
As a result of the operation of the element on the copper electrode, the oxidation process will occur:
,
and on the silver electrode, the recovery process:
.
The processes of oxidation and reduction in a galvanic cell are spatially separated.
Electrode ,
on which it flows oxidation process, is called anode
(
).
The electrode on which flows recovery process, is called cathode
(
).
The reactions at the cathode and anode are called electrode reactions.
The total chemical process occurring in a galvanic cell consists of electrode processes and is expressed by the equation:
If electrode processes and a chemical reaction in a galvanic cell can be carried out in the forward (during the operation of the cell) and reverse (when electric current is passed through the cell) directions, then such electrodes and a galvanic cell are called reversible.
In what follows, only reversible electrodes and galvanic cells will be considered.
The mutual transformation of electrical and chemical forms of energy occurs in electrochemical systems, including:
ª conductors of the second kind - substances with ionic conductivity (electrolytes).
ª conductors of the first kind - substances with electronic conductivity.
At the interface between two phases, an electric charge is transferred, i.e. there is a potential jump ().
A system consisting of contacting conductors of the first and second kind is called electrode.
The processes occurring at the phase boundary of conductors of the I and II kind in the electrodes are calledelectrode processes .
The electrode is a system consisting of at least two phases.
Let us consider how a potential jump occurs - the electrode potential - at the interface between the metal and the salt solution of this metal. When a metal plate is immersed in a salt solution, some of the metal ions from the surface of the plate can go into a solution adjacent to the surface of the plate. The metal is charged negatively, and the resulting electrostatic forces prevent the further flow of this process. The system is in equilibrium. The reverse process of transition of metal cations from solution to the plate is also possible. These processes lead to the appearance of a double electric layer and a potential jump.
The direction of the process of transfer of metal ions is determined by the ratio of the electrochemical potentials of ions () in the solution phase and the condensed phase. The process continues until the electrochemical potentials in the two phases are equalized.
The electrochemical potential consists of two terms
m chem. - chemical potential that characterizes the chemical response to a change in the environment of a given particle.
m el - the electrical component of the electrochemical potential or the potential energy of the electric field, which characterizes the response to the electric field.
For a certain kind of charged particles (i)
z i is the charge of the ion,
– internal potential, corresponding to the work of transfer of an elementary negative charge from infinity in vacuum deep into the phase.
Equilibrium of an electrochemical system characterized by the equality of electrochemical (rather than chemical) potentials of charged particles in different phases.
In the equilibrium system solution (I) / metal (II), we have:
In a non-equilibrium system, the work of transfer of one mol-equiv. ions from phase I to phase II is
Since then
In equilibrium, taking into account (1), we have:
where is the jump at the interface (absolute electrode potential). Denote
where is the potential jump at the phase boundary at a i = 1 (standard electrode potential).
The standard potential is a value characteristic of a given electrode process. It depends on the temperature and the nature of the electrode. Then for an electrode of type Me Z+ /Me:
A potential jump also occurs at the interface between two solutions, this is the diffusion potential.
In general terms (for any type of electrodes):
or for 298K
It should be remembered that if gases are involved in the electrode reaction, then the activity is assumed to be equal to the partial pressure; for the condensed phase of constant composition, A=1.
Equations (1), (2) are called Nernst equations for the electrode potential. The electric potential difference can be experimentally measured only between two points of the same phase where μ i = const. When an elementary charge moves between two points that are in different phases, in addition to the electric one, work must be performed associated with a change in the chemical environment of the charge. The magnitude of this chemical component of the work cannot be determined, so the absolute value of the electrode potential cannot be measured. Empirically, it is possible to determine only the magnitude of the EMF of a galvanic cell consisting of two electrodes.
Rules for recording electrodes and electrochemical circuits.
Systems consisting of two or more electrodes, connected in a special way and capable of producing electrical work, that is, serving as a source of electrical energy, are called galvanic cells.
Electromotive force of a galvanic cell(EMF GE) is the sum of jumps in electrode potentials at all phase boundaries in the equilibrium condition (the current in the external circuit is zero).
a) The following recording rules are accepted for electrodes: substances in solution are indicated to the left of the vertical bar, substances forming another phase (gas or solid) are indicated to the right.
If one phase contains several substances, then their characters are separated by commas.
For example,
The equation of the electrode reaction for a separate electrode is written in such a way that substances in the oxidized form and electrons are located on the left, and substances in the reduced form are on the right:
b) When recording galvanic cells, an electrode with a more negative potential is located on the left; the solutions of both electrodes are separated from each other by a vertical dotted line if they are in contact with each other, and by two solid lines if there is a salt bridge between the solutions, for example, a saturated KCl solution, with which the diffusion potential is eliminated. Thus, the positively charged electrode is always indicated on the right, and the negatively charged electrode is always indicated on the left.
Electrode , on which it flows oxidation process, is called anode ().
The electrode on which flows recovery process, is called cathode ().
The reactions at the cathode and anode are called electrode reactions.
The total chemical process occurring in a galvanic cell consists of electrode processes and is expressed by the equation:
If the electrode processes and the chemical reaction in a galvanic cell can be carried out in the forward (during the operation of the cell) and reverse (when electric current is passed through the cell) directions, then such electrodes and the galvanic cell are called reversible.
In what follows, only reversible electrodes and galvanic cells will be considered.
