Crafts

As a result, a diffusion potential arises. Diffuse potential, mechanism of origin and biological significance. Membrane diffusion potential

At the boundary of two dissimilar solutions, a potential difference always arises, which is called the diffusion potential. The emergence of such a potential is associated with the unequal mobility of cations and anions in solution. The magnitude of the diffusion potentials usually does not exceed several tens of millivolts, and they, as a rule, are not taken into account. However, with accurate measurements, special measures are taken to minimize them. The reasons for the appearance of the diffusion potential were shown by the example of two bordering solutions of copper sulfate of different concentrations. Ions Cu2 + and SO42- will diffuse across the interface from a more concentrated solution to a less concentrated one. The rates of movement of Cu2 + and SO42- ions are not the same: the mobility of SO42- ions is greater than the mobility of Cu2 +. As a result, an excess of negative SO42- ions appears at the interfaces of solutions on the side of the solution with a lower concentration, and an excess of Cu2 + occurs in a more concentrated one. A potential difference arises. The presence of an excess negative charge at the interface will slow down the movement of SO42- and accelerate the movement of Cu2 +. At a certain value of the potential, the velocities of SO42- and Cu2 + will become the same; the stationary value of the diffusion potential is established. The theory of diffusion potential was developed by M. Planck (1890), and later by A. Henderson (1907). The formulas they obtained for the calculation are complex. But the solution is simplified if the diffusion potential arises at the interface of two solutions with different concentrations of C1 and C2 of the same electrolyte. In this case, the diffusion potential is. Diffusion potentials arise during nonequilibrium diffusion processes, therefore they are irreversible. Their value depends on the nature of the boundary of two contacting solutions, on the size and their configuration. Accurate measurements use methods that minimize the value of the diffuse potential. For this purpose, an intermediate solution with the lowest possible mobility values U and V (for example, KCl and KNO3) is included between the solutions in the half-cells.

Diffuse potentials play an important role in biology. Their occurrence is not associated with metal electrodes. It is the interphase and diffusion potentials that generate biocurrents. For example, electric rays and eels create a potential difference of up to 450 V. Biopotentials are sensitive to physiological changes in cells and organs. This is the basis for the application of the methods of electrocardiography and electroencephalography (measurement of biocurrents of the heart and brain).

55. Interfluid phase potential, mechanism of origin and biological significance.

A potential difference also arises at the interface between immiscible liquids. Positive and negative ions in these solvents are distributed unevenly, their distribution coefficients do not coincide. Therefore, a potential jump occurs at the interface between the liquids, which prevents the unequal distribution of cations and anions in both solvents. In the total (total) volume of each phase, the amount of cations and anions is practically the same. It will differ only at the interface. This is the interfluid potential. Diffuse and interfluid potentials play an important role in biology. Their occurrence is not associated with metal electrodes. It is the interphase and diffusion potentials that generate biocurrents. For example, electric rays and eels create a potential difference of up to 450 V. Biopotentials are sensitive to physiological changes in cells and organs. This is the basis for the application of electrocardiography and electroencephalography methods (measurement of biocurrents of the heart and brain).

Diffusion potentials arise at the interface between two solutions. Moreover, it can be both solutions of different substances, and solutions of the same substance, only in the latter case they must necessarily differ from each other in their concentrations.

When two solutions come into contact, particles (ions) of dissolved substances interpenetrate in them due to the diffusion process.

The reason for the appearance of the diffusion potential in this case is the unequal mobility of ions of dissolved substances. If electrolyte ions have different diffusion rates, then faster ions gradually appear ahead of less mobile ones. As if two waves of differently charged particles are formed.

If solutions of the same substance are mixed, but with different concentrations, then a more dilute solution acquires a charge that coincides in sign with the charge of more mobile ions, and a less dilute one - a charge that coincides in sign with the charge of less mobile ions (Fig. 90).

Rice. 90. The emergence of a diffusion potential due to different ion velocities: I- "fast" ions, negatively charged;
II- "slow" ions, positively charged

The so-called diffusion potential arises at the interface between the solutions. It averages the speed of ion movement (slows down the "faster" ones and accelerates the "slower" ones).

Gradually, with the completion of the diffusion process, this potential decreases to zero (usually within 1-2 hours).

Diffusion potentials can also arise in biological objects when the cell membranes are damaged. In this case, their permeability is disturbed and electrolytes can diffuse from the cell into the tissue fluid or vice versa, depending on the difference in concentration on both sides of the membrane.

As a result of the diffusion of electrolytes, a so-called damage potential arises, which can reach values of the order of 30-40 mV. Moreover, the damaged tissue is most often charged negatively in relation to the undamaged one.

Diffusion potential occurs in galvanic cells at the interface between two solutions. Therefore, with accurate calculations of the emf galvanic circuits must be corrected for its value. To eliminate the influence of the diffusion potential, electrodes in galvanic cells are often connected to each other with a "salt bridge", which is a saturated solution of KCl.

Potassium and chlorine ions have almost the same mobility; therefore, their use makes it possible to significantly reduce the effect of the diffusion potential on the emf.

The diffusion potential can greatly increase if electrolyte solutions of different compositions or different concentrations are separated by a membrane that is permeable only for ions of a certain charge sign or type. Such potentials will be much more persistent and can persist for a longer time - they are called differently membrane potentials... Membrane potentials arise when ions are unevenly distributed on both sides of the membrane, depending on its selective permeability, or as a result of ion exchange between the membrane itself and the solution.

The principle of operation of the so-called ion-selective or membrane electrode.

The basis of such an electrode is a semi-permeable membrane obtained in a certain way, which has a selective ionic conductivity. A feature of the membrane potential is that electrons do not participate in the corresponding electrode reaction. Here, an exchange of ions takes place between the membrane and the solution.

Membrane electrodes with a solid membrane contain a thin membrane, on both sides of which there are different solutions containing the same detectable ions, but with different concentrations. From the inside, the membrane is washed standard solution with a precisely known concentration of the ions to be determined, from the outside - the analyzed solution with an unknown concentration of the ions to be determined.

Due to the different concentration of solutions on both sides of the membrane, ions are exchanged with the inner and outer sides of the membrane in a different way. This leads to the fact that a different electric charge is formed on different sides of the membrane, and as a result of this, a membrane potential difference arises.

When creating any electrode pair, a "salt bridge" is always used. The use of a "salt bridge" solves several problems that arise before researchers of electrochemical processes. One of these tasks is to increase the accuracy of determinations by eliminating or significantly reducing the diffusion potential ... Diffusion potential in galvanic cells occurs when solutions of different concentrations come into contact. The electrolyte from a solution with a higher concentration diffuses (passes) into a less concentrated solution. If the absolute velocities of movement of the cations and anions of the diffusing electrolyte are different, then the less concentrated solution acquires the potential of the “faster ions” charge sign, and the more concentrated solution acquires the potential of the opposite sign. To eliminate the diffusion potential, it is necessary to minimize the difference in the rates of movement of cations and anions of the diffusing electrolyte. For this, a saturated KCl solution was chosen, since absolute travel speeds K + and Cl ¯ are practically the same and have one of the highest values.

The appearance of a diffusion potential is also characteristic of biological systems... For example, if a cell is damaged, when the semipermeability of its membrane is disturbed, electrolyte begins to diffuse into or out of the cell. This creates a diffusion potential, which is referred to here as the "damage potential". Its value can reach 30 - 40 mV, the "damage potential" is stable for about one hour.

The value of the diffusion potential increases significantly if electrolyte solutions of different concentrations are separated by a membrane that allows only cations or anions to pass through. The selectivity of such membranes is due to their own charge. Membrane potentials are very stable and can persist for several months.

Potentiometry

Types of electrodes

For analytical and technical purposes, many different electrodes have been developed that form electrode pairs (elements).

There are two main types of electrode classification.

By chemical composition:

1. 1st kind electrodes - these are electrodes, the electrode reaction of which is reversible only by the cation or by the anion. For example, the electrodes forming the Jacobi-Daniel element are copper and zinc (see above).

2. 2nd kind electrodes - these are electrodes, the electrode reaction of which is reversible for two types of ions: both cations and anions.

3. Redox electrodes (Red - Ox) . The term "Red - Ox - electrode" is understood to mean such an electrode where all the elements of the half reaction (both the oxidized and reduced form) are in solution. Metal electrodes, immersed in a solution, do not participate in the reaction, but serve only as a carrier of electrons.

By appointment:

1. Reference electrodes .

Reference electrodes are such electrodes, the potential of which is known exactly, is stable over time and does not depend on the concentration of ions in the solution. These electrodes include: standard hydrogen electrode, calomel electrode and silver chloride electrode. Let's consider each electrode in more detail.

Standard hydrogen electrode.

This electrode is a closed vessel into which a platinum plate is inserted. The vessel is filled with a solution of hydrochloric acid, the activity of hydrogen ions in which is equal to 1 mol / l. Hydrogen gas is passed into a 1 atmosphere pressure vessel. Bubbles of hydrogen are adsorbed on a platinum plate, where they are dissociated into atomic hydrogen and oxidized.

Characteristics of a standard hydrogen electrode:

1.Electrode circuit: Pt (H 2) / H +

2.Electrode reaction: ½ H 2 - ē ↔ H +

As it is easy to see, this reaction is reversible only for the (H +) cation, therefore a standard hydrogen electrode is a type 1 electrode.

3. Calculation of the electrode potential.

The Nernst equation takes the form:

e H 2 / H + = e ° H 2 / N + RT ln a n +

nF (P n 2) 1/2

Because a n + = 1 mol / l, p n + = 1 atm, then ln a n + = 0, therefore

(R n 2) 1/2

e H 2 / H + = e ° H 2 / H +

Thus, at a n + = 1 mol / l and p (n 2) = 1 atm, the potential of the hydrogen electrode is zero and is called the "standard hydrogen potential".

Another example - calomel electrode(see picture)

It contains a paste containing calomel (Hg 2 Cl 2), mercury and potassium chloride. The paste is in pure mercury and is filled with potassium chloride solution. A platinum plate is immersed inside this system.

Electrode characteristics:

1.Scheme of electrode: Hg 2 Cl 2, Hg (Pt) / Cl¯

2. Two parallel reactions take place in this electrode:

Hg 2 Cl 2 ↔2Hg + + 2Cl¯

2 Hg + + 2ē → 2Hg

Hg 2 Cl 2 + 2ē → 2Hg + 2Cl¯ is the overall reaction.

It can be seen from the above equations that the calomel electrode is a type 2 electrode.

3. The potential of the electrode is determined by the Nernst equation, which, after appropriate transformations, takes the form:

e = e o - RT ln a Cl¯

Another important example is silver chloride electrode(see fig).

Here, the silver wire is covered with a layer of the hardly soluble AgCl salt and immersed in a saturated solution of potassium chloride.

Electrode characteristics:

1. Electrode diagram: Ag, AgCl / Cl¯

2. Electrode reactions: AgCl ↔ Ag + + Cl¯

Ag + + ē → Ag

AgCl + ē ↔ Ag + Cl¯ - total reaction.

As can be seen from this reaction, the formed metal is deposited on the wire, and the Cl¯ ions go into solution. The metal electrode acquires a positive charge, the potential of which depends on the concentration (activity) of Cl¯ ions.

3. The potential of the electrode is determined by the Nernst equation, which, after appropriate transformations, takes the already known form:

e = e o - RT ln a Cl¯

In silver chloride and calomel electrodes, the concentration of Cl¯ ions is kept constant and therefore their electrode potentials are known and constant over time.

2. Determination electrodes - these are electrodes, the potential of which depends on the concentration of any ions in the solution, therefore, the concentration of these ions can be determined by the value of the electrode potential.

Most often, the following are used as indicator electrodes: hydrogen, glass and quinhydrone electrodes.

Hydrogen electrode It is designed similarly to a standard hydrogen electrode, but if an acidic solution with the activity of H + ions is more than one is placed in the capacity of the hydrogen electrode, then a positive potential appears on the electrode, which is proportional to the activity (i.e., concentration) of protons. With a decrease in the concentration of protons, on the contrary, the electrode will be charged negatively. Therefore, by determining the potential of such an electrode, it is possible to calculate the pH of the solution in which it is immersed.

Characteristics of the electrode.

1. Electrode diagram: Pt (H 2) / H +

2. Electrode reaction: ½ H 2 - ē ↔ H +

3. e Н 2 / Н + = e o Н 2 / Н + + 0.059 log а Н +

Because n = 1 and e o H 2 / H += 0, then the Nernst equation takes the form:

e H2 / H + = 0.059 log a h + = - 0.059 pH pH = - e

0,059

Glass electrode is a silver plate coated with an insoluble silver salt, enclosed in a glass shell made of special glass, ending in a thin-walled conductive ball. The internal medium of the electrode is a hydrochloric acid solution. The potential of the electrode depends on the concentration of H + and is determined by the Nernst equation, which has the form:

e st = e about st + 0.059 lg a n +

Quinhydron electrode consists of a platinum plate immersed in a solution of quinhydrone - an equal molar mixture of quinone C 6 H 4 O 2 and hydroquinone C 6 H 4 (OH) 2, between which a dynamic equilibrium is rapidly established:

Since protons are involved in this reaction, the potential of the electrode depends on the pH.

Electrode characteristics:

1. Electrode circuit: Pt / H +, C 6 H 4 O 2, C 6 H 4 O 2-

2. Electrode response:

С 6 Н 4 (ОН) 2 - 2ē ↔ С 6 Н 4 О 2 + 2Н + -

redox process.

3. The potential of the electrode is determined by the Nernst equation, which after appropriate transformations takes the form:

fx. r = e about x. g. + 0.059 lg a H +

The quinhydrone electrode is used only for determining the pH of those solutions where this indicator is not more than 8. This is due to the fact that in an alkaline medium hydroquinone behaves like an acid and the value of the electrode potential ceases to depend on the concentration of protons.

Because in the quinhydrone electrode a plate made of a noble metal is immersed in a solution containing both an oxidized and a reduced form of one substance, then it can be considered as a typical "red - ox" - system.

The components of the redox system can be both organic and inorganic substances, for example:

Fe 3+ / Fe 2+ (Pt).

However, for organic matter, "Red - ox" - the electrodes are especially important because are the only way to form an electrode and determine its potential.

The values of electrode potentials arising on metal plates in red - ox - systems, can be calculated not only by the Nernst equation, but also by the Peters equation:

2 * 10 -4 C ox

e red-ox = e 0 red-ox + * T * lg;(V)

T- temperature, 0 K.

C ox and C red- the concentration of the oxidized and reduced forms of the substance, respectively.

e 0 red - ox is the standard redox potential that occurs in the system when the ratio of the concentrations of the oxidized and reduced forms of the compound is equal to 1.

In transfer cells, solutions of half cells of various qualitative and quantitative compositions come into contact with each other. In general, the mobility (diffusion coefficients) of ions, their concentration, and nature in the half-cells are different. The faster ion charges the layer on one side of the imaginary boundary of the layers with its sign, leaving the oppositely charged layer on the other side. Electrostatic attraction prevents the diffusion of individual ions from developing further. There is a separation of positive and negative charges at an atomic distance, which, according to the laws of electrostatics, leads to the appearance of a jump in the electric potential, called in this case diffusion potential Df and (synonyms - liquid potential, potential of a liquid connection, contact). However, diffusion-migration of the electrolyte as a whole continues at a certain gradient of forces, chemical and electrical.

As is known, diffusion is an essentially nonequilibrium process. Diffusion potential is a nonequilibrium component of the EMF (as opposed to electrode potentials). It depends on the physicochemical characteristics of individual ions and even on the device of contact between solutions: porous diaphragm, swab, thin section, free diffusion, asbestos or silk thread, etc. Its value cannot be accurately measured, but is estimated experimentally and theoretically with varying degrees of approximation.

For the theoretical assessment of Dph 0, various approaches of Dp4B are used. In one of them, called quasi-thermodynamic, the electrochemical process in the transfer cell is generally considered reversible, and diffusion is stationary. It is assumed that a certain transition layer is created at the boundary of solutions, the composition of which changes continuously from solution (1) to solution (2). This layer is mentally divided into thin sublayers, the composition of which, i.e., concentrations, and with them chemical and electrical potentials, change by an infinitely small amount in comparison with the neighboring sublayer:

The same ratios are maintained between subsequent sublayers, and so on until solution (2). Stationarity consists in the immutability of the picture in time.

Under the conditions for measuring the EMF, a diffusion transfer of charges and ions occurs between the sublayers, i.e., electrical and chemical work is performed, separable only mentally, as in the derivation of the equation of electrochemical potential (1.6). We consider the system to be infinitely large, and count on 1 equiv. substance and 1 Faraday charge carried by each type of participating ions:

On the right, there is a minus, because the work of diffusion is carried out in the direction of the decrease in the force - the gradient of the chemical potential; t; is the transfer number, that is, the fraction of the charge transferred by a given / -th type of ions.

For all participating ions and for the entire sum of sublayers that make up the transition layer from solution (1) to solution (2), we have:

Let us note on the left the definition of the diffusion potential as an integral value of the potential, continuously varying in the composition of the transition layer between solutions. Substituting | 1, = | φ + /? Γ1nr, and taking into account that (I, = const for p, T = const, we get:

Seeking a relationship between diffusion potential and ion characteristics such as transport numbers, charge and activity of individual ions. The latter, as is known, are not thermodynamically definable, which complicates the calculation of A (p D, requiring non-thermodynamic assumptions. Integration of the right-hand side of equation (4.12) is performed under various assumptions about the structure of the interface between solutions.

M. Planck (1890) considered the border to be sharp, the layer is thin. Integration under these conditions led to the Planck equation for Δφ 0, which turned out to be transcendental with respect to this quantity. Its solution is found by an iterative method.

Henderson (1907) derived his equation for Dph 0, proceeding from the assumption that a transition layer of thickness is created between the contacting solutions d, the composition of which varies linearly from solution (1) to solution (2), i.e.

Here WITH; is the concentration of the ion, x is the coordinate inside the layer. When integrating the right-hand side of expression (4.12), the following assumptions are made:

ion activity a, replaced by concentration C, (Henderson did not know any activities!);
the transfer numbers (ion mobility) are taken to be independent of the concentration and constant within the layer.

Then the general Henderson equation is obtained:

Zj,С „“, - charge, concentration and electrolytic mobility of the ion in solutions (1) and (2); the + and _ signs at the top refer to cations and anions, respectively.

The expression for the diffusion potential reflects the differences in the characteristics of ions on different sides of the boundary, i.e., in solution (1) and in solution (2). To estimate Δφ 0, it is the Henderson equation that is most often used, which is simplified in typical special cases of cells with transfer. In this case, various characteristics of ion mobility are used, associated with and, - ionic conductivities, transfer numbers (Table 2.2), i.e. values available from look-up tables.

Henderson's formula (4.13) can be written somewhat more compactly if we use ionic conductivity:

(here the designations of solutions 1 and 2 are replaced by "and", respectively).

A consequence of the general expressions (4.13) and (4.14) are some particular ones given below. It should be borne in mind that the use of concentrations instead of ionic activities and the characteristics of the mobility (electrical conductivity) of ions at infinite dilution makes these formulas very approximate (but the more accurate, the more dilute the solutions). In a more rigorous derivation, the dependences of the characteristics of mobility and transfer numbers on concentration are taken into account, and instead of concentrations, there are the activities of ions, which, with a certain degree of approximation, can be replaced by the average activities of the electrolyte.

Special cases:

For the border of two solutions of the same concentration of different electrolytes with a common ion of the type AX and BX, or AX and AY:

(Lewis - Sergeant formulas), where - the limiting molar electrical conductivity of the corresponding ions, A 0 - the limiting molar electrical conductivity of the corresponding electrolytes. For electrolytes type AX 2 and BX 2

WITH and WITH" the same electrolyte type 1: 1

where V) and A.® are the limiting molar electrical conductivities of cations and anions, t and r +- transfer numbers of anion and cation of the electrolyte.

For the boundary of two solutions of different concentrations WITH" and C "of the same electrolyte with cation charges z +, anions z ~, carry numbers t + and t_ respectively

For an electrolyte of the type Mn + Ag _, taking into account the condition of electroneutrality v + z + = -v_z_ and stoichiometric relationship C + = v + C and C_ = v_C, you can simplify this expression:

The above expressions for the diffusion potential reflect the differences in the mobility (transfer numbers) and the concentration of cations and anions on different sides of the solution boundary. The smaller these differences are, the smaller the value of Dph 0. This can be seen from Table. 4.1. The highest Dphi values (tens of mV) were obtained for acid and alkali solutions containing Н f and ОН ions, which have a uniquely high mobility. The smaller the difference in mobility, i.e. the closer to 0.5 the value t + and the less Df c. This is observed for electrolytes 6-10, which are called "equal conductive" or "equal transfer".

To calculate Dph 0, we used the limiting values of the electrical conductivity (and transfer numbers), but the real values of the concentrations. This introduces a certain error, which for 1 - 1 electrolytes (nos. 1 - 11) ranges from 0 to ± 3%, while for electrolytes containing ions with a charge change in ionic strength which

it is the multiply charged ions that make the greatest contribution.

The Dph 0 values at the boundaries of solutions of different electrolytes with the same anion and the same concentrations are given in Table. 4.2.

The conclusions about diffusion potentials made earlier for solutions of the same electrolytes of different concentrations (Table 4.1) are also confirmed in the case of different electrolytes of the same concentration (columns 1-3 of Table 4.2). Diffusion potentials turn out to be highest if electrolytes containing H + or OH ions are located on opposite sides of the boundary. "They are large enough for electrolytes containing ions, the transfer numbers of which in a given solution are far from 0.5.

The calculated Afr values are in good agreement with the measured ones, especially if we take into account both the approximations used in the derivation and application of equations (4.14a) and (4.14c), and the experimental difficulties (errors) when creating the boundary of liquids.

Table 41

Limiting ionic conductivity and electrical conductivity of aqueous solutions of electrolytes, transfer numbers and diffusion potentials,

calculated by the formulas (414g-414e) at for 25 ° C

Electrolyte	Cm cm mol	Cm? cm 2 mol	Cm cm 2 mol	Af s,






NH 4CI

NH 4NO 3
CH 3COOU

Have 2CaC1 2
1/2NcbSCX)
l / 3LaCl 3
1/2 CuS0 4
l / 2ZnS0 4

In practice, instead of quantifying the Afr value, most often resort to its elimination, i.e., bringing its value to a minimum (up to several millivolts) by switching on between contacting solutions electrolytic bridge("Key") filled with a concentrated solution of the so-called equal conducting electrolyte, i.e.

electrolyte, cations and anions of which have close mobility and, accordingly, ~ / + ~ 0.5 (Nos. 6-10 in Table 4.1). The ions of such an electrolyte, taken in a high concentration in relation to the electrolytes in the cell (in a concentration close to saturation), take on the role of the main charge carriers across the boundary of solutions. Due to the closeness of the mobility of these ions and their predominant concentration, Dpho -> 0 mV. This is illustrated by columns 4 and 5 of Table. 4.2. Diffusion potentials at the boundaries of NaCl and KCl solutions with concentrated KCl solutions are indeed close to 0. At the same time, at the boundaries of concentrated KCl solutions, even with dilute solutions of acid and alkali, D (р в is not equal to 0 and increases with increasing concentration of the latter.

Table 4.2

Diffusion potentials at the boundaries of solutions of different electrolytes, calculated using formula (4.14a) at 25 ° С

Liquid connection "1	exp. 6 ',	Liquid connection a), d>

ns1 o.1: kci od		HCI 1.0 \|\| KCl Sa,
		HC1 0.1CKS1 Sat
		HC1 0.01CKS1 &,
HC10.1: NaCl 0.1		NaCl 1.0 \|\| KCI 3.5
HCI 0.01 iNaCl 0.01		NaCl 0.11 \| KCI 3.5
HCI 0.01 ILiCl 0.01
KCI 0.1 iNaCl 0.1		KCI 0.1CKS1 Sat
KCI 0.01 iNaCl 0.01
KCI 0.01 iLiCl 0.01		NaOH 0.1CCS1 Sal
Kci o.oi: nh 4 ci o.oi		NaOH 1.0CCS1 Sat
LiCl 0.01: nh 4 ci 0.01		NaOH 1.0CKS1 3.5
LiCl 0.01 iNaCl 0.01		NaOH 0.1CKS1 0.1

Notes:

Concentrations in mol / L.

61 Measurements of EMF of cells with and without transfer; calculation taking into account average activity coefficients; see below.

Calculation using the Lewis - Sergeant equation (4L4a).

"KCl Sal is a saturated solution of KCl (~ 4.16 mol / L).

"Calculation according to Henderson's equation of the type (4.13), but using average activities instead of concentrations.

Diffusion potentials on each side of the bridge have opposite signs, which contributes to the elimination of the total Df 0, which in this case is called residual(residual) diffusion potential DDf and res.

The boundary of liquids, on which Df p is eliminated by the inclusion of an electrolytic bridge, is usually denoted (||), as is done in Table. 4.2.

Appendix 4B.

The voltage of an electrochemical system with a liquid interface between two electrolytes is determined by the difference in electrode potentials accurate to the diffusion potential.

Rice. 6.12. Elimination of diffusion potential with electrolytic bridges

Generally speaking, the diffusion potentials at the interface between two electrolytes can be quite significant and, in any case, often make the measurement results uncertain. Below are the values of diffusion potentials for some systems (the electrolyte concentration in kmol / m 3 is indicated in brackets):

Therefore, the diffusion potential must either be eliminated or accurately measured. Elimination of the diffusion potential is achieved by including an additional electrolyte with close values of the cation and anion mobilities into the electrochemical system. For measurements in aqueous solutions, saturated solutions of potassium chloride, potassium or ammonium nitrate are used as such an electrolyte.

Additional electrolyte is connected between the main electrolytes using electrolytic bridges (Fig. 6.12) filled with basic electrolytes. Then the diffusion potential between the main electrolytes, for example, in the case shown in Fig. 6.12, - between solutions of sulfuric acid and copper sulfate, is replaced by diffusion potentials at the boundaries of sulfuric acid - potassium chloride and potassium chloride - copper sulfate. At the same time, at the boundaries with potassium chloride, electricity is mainly carried by ions K + and C1 -, which are much more than ions of the main electrolyte. Since the mobilities of K + and C1 - ions in potassium chloride are practically equal to each other, the diffusion potential will also be small. If the concentrations of the main electrolytes are low, then with the help of additional electrolytes, the diffusion potential is usually reduced to values not exceeding 1 - 2 mV. So, in the experiments of Abbeg and Cumming it was established that the diffusion potential at the boundary of 1 kmol / m 3 LiCl - 0.1 kmol / m 3 LiCl is 16.9 mV. If additional electrolytes are included between the lithium chloride solutions, then the diffusion potential decreases to the following values:

Additional electrolyte Diffusion potential of the system, mV

NH 4 NO 3 (1 kmol / m 3) 5.0

NH 4 NO 3 (5 kmol / m 3) –0.2

NH 4 NO 3 (10 kmol / m 3) –0.7

KNO 3 (sat.) 2.8

KCl (sat.) 1.5

Elimination of diffusion potentials by incorporating an additional electrolyte with equal ion transfer numbers gives good results when measuring diffusion potentials in unconcentrated solutions with slightly different anion and cation mobilities. When measuring the stresses of systems containing solutions of acids or alkalis

Table 6.3. Diffusion potentials at the KOH - KCl and NaOH - KCl interface (according to V.G. Lokshtanov)

with very different speeds of movement of the cation and anion, one should be especially careful. For example, at the HC1 - KC1 (saturation) boundary, the diffusion potential does not exceed 1 mV, only if the concentration of the HC1 solution is below 0.1 kmol / m 3. Otherwise, the diffusion potential increases rapidly. A similar phenomenon is observed for alkalis (Table 6.3). So, the diffusion potential, for example, in the system

(-) (Pt) H 2 | KOH | KOH | H 2 (Pt) (+)

4.2 kmol / m 3 20.4 kmol / m 3

is 99 mV, and in this case it is impossible to achieve a significant reduction using the salt bridge.

To reduce the diffusion potentials to negligible values, Nernst suggested adding a large excess of some electrolyte indifferent for the given system to the contacting solutions. Then the diffusion of basic electrolytes will no longer lead to the emergence of a significant activity gradient at the interface, and, consequently, the diffusion potential. Unfortunately, the addition of an indifferent electrolyte changes the activity of the ions participating in the potential-determining reaction and leads to distortion of the results. Therefore, this method can only be used in those

in cases where the addition of an indifferent electrolyte cannot affect the change in activity or this change can be taken into account. For example, when measuring the system voltage Zn | ZnSO 4 | CuSO 4 | Cu, in which the concentration of sulfates is not lower than 1.0 kmol / m 3, the addition of magnesium sulfate to reduce the diffusion potential is quite acceptable, because the average ionic activity coefficients of zinc and copper sulfates will practically not change.

If, when measuring the voltage of an electrochemical system, diffusion potentials are not eliminated or must be measured, then first of all care should be taken to create a stable interface between the two solutions. A continuously renewing border is created by a slow directional movement of solutions parallel to each other. Thus, it is possible to achieve the stability of the diffusion potential and its reproducibility with an accuracy of 0.1 mV.

The diffusion potential is determined by the method of Cohen and Thombrock from measurements of the voltages of two electrochemical systems, and the electrodes of one of them are reversible to the salt cation, and the other to the anion. Let's say you need to determine the diffusion potential at the ZnSO 4 (a 1) / ZnSO 4 (a 2) interface. To do this, we measure the voltages of the following electrochemical systems (assume that a 1< < а 2):

1. (-) Zn | ZnSO 4 | ZnSO 4 | Zn (+)

2. (-) Hg | Hg 2 SO 4 (tv.), ZnSO 4 | ZnSO 4, Hg 2 SO 4 (tv.) | Hg (+)

System voltage 1

system 2

Considering that φ d 21 = - φ d 12, and subtracting the second equation from the first, we get:

When measurements are carried out at not very high concentrations, at which it is still possible to consider that = and = or that: =: the last two terms of the last equation cancel out and

The diffusion potential in system 1 can also be determined in a slightly different way, if instead of system 2 we use a double electrochemical system:

3. (-) Zn | ZnSO 4, Hg 2 SO 4 (tv.) | Hg - Hg | Hg 2 SO 4 (tv.), ZnSO 4 | Zn (+)