As already indicated, concentration chains have a large practical significance, since with their help it is possible to determine such important quantities as the activity coefficient and activity of ions, the solubility of slightly soluble salts, transfer numbers, etc. Such chains are practically easy to implement and the relationships connecting the EMF of the concentration chain with the activities of the ions are also simpler than for other chains. Let us recall that an electrochemical circuit containing the boundary of two solutions is called a transfer circuit and its diagram is depicted as follows:
Me 1 ½ solution (I) solution (II) ½ Me 2 ½ Me 1,
where the dotted vertical line indicates the existence of a diffusion potential between two solutions, which is galvani - the potential between points located in different directions chemical composition phases, and therefore cannot be accurately measured. The magnitude of the diffusion potential is included in the amount for calculating the EMF of the circuit:
The small value of the EMF of a concentration chain and the need for its accurate measurement make it especially important to either completely eliminate or accurately calculate the diffusion potential that arises at the boundary of two solutions in such a chain. Consider the concentration chain
Me½Me z+ ½Me z+ ½Me
Let us write the Nernst equation for each of the electrodes of this circuit:
for left 
for the right 
Let us assume that the activity of metal ions at the right electrode is greater than that at the left, i.e.
Then it is obvious that j 2 is more positive than j 1 and the emf of the concentration circuit (E k) (without diffusion potential) is equal to the potential difference j 2 – j 1.
Hence, 
, (7.84)
then at T = 25 0 C 
, (7.85)
where and are the molal concentrations of Me z + ions; g 1 and g 2 are the activity coefficients of Me z + ions, respectively, at the left (1) and right (2) electrodes.
a) Determination of average ionic activity coefficients of electrolytes in solutions
To most accurately determine the activity coefficient, it is necessary to measure the EMF of the concentration chain without transfer, i.e. when there is no diffusion potential.
Consider an element consisting of a silver chloride electrode immersed in a solution of HCl (molality C m) and a hydrogen electrode:
(–) Pt, H 2 ½HCl½AgCl, Ag (+)
Processes occurring on the electrodes:
(–) H 2 ® 2H + + 2
(+) 2AgCl + 2 ® 2Ag + 2Cl –
current-generating reaction H 2 + 2AgCl ® 2H + + 2Ag + 2Cl –
Nernst equation
for hydrogen electrode: 
( = 1 atm)
for silver chloride: 
It is known that
= 
(7.86)
Considering that the average ionic activity for HCl is
And 
,
where C m is the molal concentration of the electrolyte;
g ± – average ionic activity coefficient of the electrolyte,
we get 
(7.87)
To calculate g ± from EMF measurement data, it is necessary to know the standard potential of the silver chloride electrode, which in this case will also be the standard EMF value (E 0), since The standard potential of a hydrogen electrode is 0.
After transforming equation (7.6.10) we get
(7.88)
Equation (7.6.88) contains two unknown quantities j 0 and g ±.
According to the Debye–Hückel theory for dilute solutions of 1-1 electrolytes
lng ± = –A ,
where A is the coefficient of Debye’s limit law and, according to reference data for this case, A = 0.51.
Therefore, the last equation (7.88) can be rewritten as follows:
(7.89)
To determine, build a dependence graph 
from and extrapolate to C m = 0 (Fig. 7.19).
Rice. 7.19. Graph for determining E 0 when calculating g ± HCl solution
The segment cut off from the ordinate axis will be the value j 0 of the silver chloride electrode. Knowing , you can use the experimental values of E and the known molality for a solution of HCl (C m), using equation (7.6.88), to find g ±:
(7.90)
b) Determination of the solubility product
Knowledge of standard potentials makes it easy to calculate the solubility product of a sparingly soluble salt or oxide.
For example, consider AgCl: PR = L AgCl = a Ag + . a Cl –
Let us express L AgCl in terms of standard potentials, according to the electrode reaction
AgCl – AgCl+,
running on a type II electrode
Cl – / AgCl, Ag
And the reactions Ag + + Ag,
running on the I-type electrode with a current-generating reaction
Cl – + Ag + ®AgCl
; 
,
because j 1 = j 2 (electrode is the same) after transformation:
(7.91)
= PR
The values of standard potentials are taken from the reference book, then it is easy to calculate the PR.
V) Diffusion potential concentration chain. Definition of carry numbers
Consider a conventional concentration chain using a salt bridge to eliminate the diffusion potential
(–) Ag½AgNO 3 ½AgNO 3 ½Ag (+)
The emf of such a circuit without taking into account the diffusion potential is equal to:
(7.92)
Consider the same circuit without a salt bridge:
(–) Ag½AgNO 3 AgNO 3 ½Ag (+)
EMF of the concentration circuit taking into account the diffusion potential:
E KD = E K + j D (7.93)
Let 1 faraday of electricity pass through the solution. Each type of ion transfers a portion of this amount of electricity equal to its transport number (t + or t –). The amount of electricity that cations and anions will transfer will be equal to t +. F and t – . F accordingly. At the border of contact of two AgNO 3 solutions of different activities, a diffusion potential (j D) arises. Cations and anions, overcoming (j D), perform electrical work.
Per 1 mole:
DG = –W el = – zFj D = – Fj d (7.94)
In the absence of diffusion potential, ions perform only chemical work when crossing the solution boundary. In this case, the isobaric potential of the system changes:
Similarly for the second solution:
|
Then according to equation (7.6.18)
(7.99)
Let us transform expression (7.99), taking into account expression (7.94):
(7.100)
(7.101)
Transport numbers (t + and t –) can be expressed in terms of ionic conductivities:
;
Then 
(7.102)
If l – > l +, then j d > 0 (diffusion potential helps the movement of ions).
If l + > l – , then j d< 0 (диффузионный потенциал препятствует движению ионов, уменьшает ЭДС). Если l + = l – , то j д = 0.
If we substitute the value jd from equation (7.101) into equation (7.99), we obtain
E KD = E K + E K (t – – t +), (7.103)
after conversion:
E KD = E K + (1 + t – – t +) (7.104)
It is known that t + + t – = 1; then t + = 1 – t – and the expression
(7.105)
If we express ECD in terms of conductivity, we get:
E KD = 
(7.106)
By measuring ECD experimentally, it is possible to determine the transport numbers of ions, their mobility and ionic conductivity. This method is much simpler and more convenient than the Hittorf method.
Thus, using the experimental determination of various physicochemical quantities, it is possible to carry out quantitative calculations to determine the EMF of the system.
Using concentration chains, it is possible to determine the solubility of poorly soluble salts in electrolyte solutions, activity coefficient and diffusion potential.
Electrochemical kinetics
If electrochemical thermodynamics studies equilibria at the electrode-solution boundary, then measuring the rates of processes at this boundary and elucidating the laws to which they obey is the object of studying the kinetics of electrode processes or electrochemical kinetics.
Electrolysis
Faraday's laws
Since the passage electric current through electrochemical systems is associated with a chemical transformation, then there must be a certain relationship between the amount of electricity and the amount of reacted substances. This dependence was discovered by Faraday (1833-1834) and was reflected in the first quantitative laws of electrochemistry, called Faraday's laws.
Electrolysis – the occurrence of chemical transformations in an electrochemical system when an electric current from an external source is passed through it. By electrolysis it is possible to carry out processes whose spontaneous occurrence is impossible according to the laws of thermodynamics. For example, the decomposition of HCl (1M) into elements is accompanied by an increase in the Gibbs energy of 131.26 kJ/mol. However, under the influence of electric current this process can easily be carried out.
Faraday's first law.
The amount of substance reacted on the electrodes is proportional to the strength of the current passing through the system and the time of its passage.
Mathematically expressed:
Dm = keI t = keq, (7.107)
where Dm is the amount of reacted substance;
kе – some proportionality coefficient;
q – amount of electricity equal to the product of force
current I for time t.
If q = It = 1, then Dm = k e, i.e. the coefficient k e represents the amount of substance that reacts when a unit amount of electricity flows. The proportionality coefficient k e is called electrochemical equivalent . Since different quantities can be chosen as a unit of the amount of electricity (1 C = 1A. s; 1F = 26.8 A. h = 96500 K), then for the same reaction one should distinguish between electrochemical equivalents related to these three units : A. with k e, A. h k e and F k e.
Faraday's second law.
During the electrochemical decomposition of various electrolytes with the same amount of electricity, the content of the electrochemical reaction products obtained on the electrodes is proportional to their chemical equivalents.
According to Faraday's second law, at a constant amount of electricity passed, the masses of reacted substances are related to each other as their chemical equivalents A.
. (7.108)
If we choose the faraday as the unit of electricity, then
Dm 1 = F k e 1; Dm 2 = F k e 2 and Dm 3 = F k e 3, (7.109)
(7.110)
The last equation allows us to combine both Faraday's laws in the form of one general law, according to which an amount of electricity equal to one Faraday (1F or 96500 C, or 26.8 Ah) always electrochemically changes one gram equivalent of any substance, regardless of its nature .
Faraday's laws apply not only to aqueous and non-aqueous salt solutions at ordinary temperatures, but are also valid in the case of high-temperature electrolysis of molten salts.
Substance output by current
Faraday's laws are the most general and precise quantitative laws of electrochemistry. However, in most cases, a smaller amount of a given substance undergoes electrochemical change than calculated on the basis of Faraday's laws. So, for example, if you pass a current through an acidified solution of zinc sulfate, then when 1F of electricity passes, not 1 g-eq of zinc is usually released, but about 0.6 g-eq. If solutions of chlorides are subjected to electrolysis, then as a result of passing 1F electricity, not one, but a little more than 0.8 g-equiv of chlorine gas is formed. Such deviations from Faraday's laws are associated with the occurrence of side electrochemical processes. In the first of the examples discussed, two reactions actually occur at the cathode:
zinc precipitation reaction
Zn 2+ + 2 = Zn
and the reaction to form hydrogen gas
2Н + + 2 = Н 2
The results obtained during the release of chlorine will also not contradict Faraday’s laws, if we take into account that part of the current is spent on the formation of oxygen and, in addition, the chlorine released at the anode can partially go back into solution due to secondary chemical reactions, for example according to the equation
Cl 2 + H 2 O = HCl + HСlO
To take into account the influence of parallel, side and secondary reactions, the concept was introduced current output P . Current output is the portion of the amount of electricity flowing that accounts for a given electrode reaction.
R = (7.111)
or as a percentage
R = . 100 %, (7.112)
where q i is the amount of electricity spent on this reaction;
Sq i is the total amount of electricity passed.
So, in the first example, the current efficiency of zinc is 60%, and that of hydrogen is 40%. Often the expression for current efficiency is written in a different form:
R = . 100 %, (7.113)
where q p and q p are the amount of electricity, respectively calculated according to Faraday’s law and actually used for the electrochemical transformation of a given amount of substance.
You can also define the current output as the ratio of the amount of changed substance Dm p to that which would have to react if all the current were spent only on this reaction Dm p:
R = . 100 %. (7.114)
If only one of several possible processes is desired, then it is necessary that its current output be as high as possible. There are systems in which all the current is spent on just one electrochemical reaction. Such electrochemical systems are used to measure the amount of electricity passed and are called coulometers, or coulometers.
In transfer cells, half-cell solutions of different qualitative and quantitative composition come into contact with each other. The mobilities (diffusion coefficients) of ions, their concentrations and nature in half-cells generally differ. The faster ion charges the layer on one side of the imaginary layer boundary with its sign, leaving a layer charged in the opposite direction on the other side. Electrostatic attraction prevents the process of diffusion of individual ions from developing further. A separation of positive and negative charges occurs at an atomic distance, which, according to the laws of electrostatics, leads to a jump in the electrical potential, called in this case diffusion potential Df and (synonyms - liquid potential, potential of liquid connection, contact). However, diffusion-migration of the electrolyte generally continues under a certain gradient of forces, chemical and electrical.
As is known, diffusion is a significantly nonequilibrium process. Diffusion potential is a nonequilibrium component of the EMF (in contrast to electrode potentials). It depends on the physicochemical characteristics of individual ions and even on the contact device between solutions: porous diaphragm, tampon, 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 theoretical assessment of Df 0, various approaches Add4V are used. In one of them, called quasi-thermodynamic, the electrochemical process in a cell with transfer is generally considered reversible, and diffusion is considered stationary. It is assumed that a certain transition layer is created at the solution boundary, 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 infinitesimal amount compared to the neighboring sublayer:
The same relationships are maintained between subsequent sublayers, and so on until solution (2). Stationarity means that the picture remains unchanged over time.
Under the conditions of measuring the EMF, a diffusion transfer of charges and ions occurs between the sublayers, i.e., electrical and chemical work is performed, separated only mentally, as when deriving the electrochemical potential equation (1.6). We consider the system to be infinitely large, and count on 1 eq. substances and 1 Faraday of charge carried by each type of participating ions:
On the right is minus, because the work of diffusion is performed in the direction of the decrease in force - the gradient of the chemical potential; t;- transfer number, i.e. the fraction of charge transferred by a given type of ion.
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 that continuously changes in the composition of the transition layer between solutions. Substituting |1, = |ф +/?Г1пй, and taking into account that (I, =const at p,T= const, we get:
The sought relationship between diffusion potential and ion characteristics, such as transport numbers, charge and activities of individual ions. The latter, as is known, are thermodynamically indeterminable, which complicates the calculation of A(p D , requiring non-thermodynamic assumptions. Integration of the right side of equation (4.12) is carried out under various assumptions about the structure of the boundary between solutions.
M. Planck (1890) considered the boundary to be sharp and the layer to be thin. Integration under these conditions led to the receipt of Planck's equation for Df 0, which turned out to be transcendental with respect to this quantity. Its solution is found using the iterative method.
Henderson (1907) derived his equation for Df 0 based on the assumption that a transition layer of thickness is created between contacting solutions d, the composition of which changes linearly from solution (1) to solution (2), i.e.
Here WITH;- ion concentration, x - coordinate inside the layer. When integrating the right-hand side of expression (4.12), the following assumptions are made:
- ion activity A, replaced by concentrations C, (Genderson did not even know the activities!);
- transfer numbers (ion mobility) are assumed to be independent of concentration and constant within the layer.
Then it turns out general equation Henderson:
Zj, C„“, - 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 Df 0, the Henderson equation is most often used, which is simplified in typical special cases of cells with translation. In this case, various characteristics of ion mobility associated with And, - ionic electrical conductivities, transfer numbers (Table 2.2), i.e., values available from reference tables.
Henderson's formula (4.13) can be written somewhat more compactly if we use ionic conductivities:
(here the designations for solutions 1 and 2 are replaced by " and ", respectively).
Consequence 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 characteristics of the mobility (electrical conductivity) of ions at infinite dilution makes these formulas very approximate (but the more accurate the more diluted the solutions are). A more rigorous derivation takes into account the dependence of mobility characteristics and transfer numbers on concentration, and instead of concentrations there are ion activities, which, with a certain degree of approximation, can be replaced by average electrolyte activities.
Special cases:
For the boundary of two solutions of the same concentration of different electrolytes with a common ion of type AX and BX, or AX and AY:
(Lewis - Sergent formulas), where 
- limiting molar electrical conductivity of the corresponding ions, A 0 - 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 g +- 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 type M„+A g _, taking into account the electrical neutrality condition v + z + = -v_z_ and the stoichiometric ratio C + = v + C and C_ = v_C, we can simplify this expression:
The given expressions for the diffusion potential reflect the differences in mobility (transfer numbers) and concentrations of cations and anions on different sides of the solution boundary. The smaller these differences are, the smaller the value of Df 0. This can also be seen from Table. 4.1. The highest Dfi values (tens of mV) were obtained for solutions of acids and alkalis containing H f and OH ions, which have 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 “equally conducting” or “equally transferring”.
For calculations of Df 0, the limiting values of electrical conductivities (and transfer numbers) were used, but real concentration values were used. This introduces a certain error, which for 1 - 1 electrolytes (Nos. 1 - 11) ranges from 0 to ±3%, while for electrolytes containing ions with charge |r,|>2 the error should be greater, because the electrical conductivity changes with change in ionic strength 
which
It is the multiply charged ions that make the greatest contribution.
The values of Df 0 at the boundaries of solutions of different electrolytes with the same anion and the same concentrations are given in Table. 4.2.
Conclusions about diffusion potentials made earlier for solutions of identical electrolytes of different concentrations (Table 4.1) are confirmed in the case of different electrolytes of the same concentration (columns 1-3 of Table 4.2). Diffusion potentials are greatest if there are electrolytes containing H + or OH ions on opposite sides of the boundary." They are quite large for electrolytes containing ions whose transport numbers in a given solution are far from 0.5.
The calculated values of Aph agree well 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) in creating the liquid boundary.
Table 41
Limiting ionic conductivities and electrical conductivities aqueous solutions electrolytes, transport numbers and diffusion potentials,
calculated using formulas (414g-414e) at 
for 25 °C
|
Electrolyte |
cm cm mole |
Cm? cm 2 mol |
cm cm 2 mol |
Af s, |
||
|
N.H. 4C.I. |
||||||
|
N.H. 4NO 3 |
||||||
|
CH 3COOU |
||||||
|
U 2CaC1 2 |
||||||
|
1/2NcbSCX) |
||||||
|
l/3LaCl 3 |
||||||
|
1/2 CuS0 4 |
||||||
|
l/2ZnS0 4 |
In practice, most often, instead of quantitatively assessing the value of Afr, they resort to its elimination, i.e., bringing its value to a minimum (up to several millivolts) by including between contacting solutions electrolytic bridge(“key”) filled with a concentrated solution of the so-called equiconducting electrolyte, i.e.
electrolyte, the cations and anions of which have similar mobilities and, accordingly, ~ / + ~ 0.5 (Nos. 6-10 in Table 4.1). The ions of such an electrolyte, taken in a high concentration relative to the electrolytes in the cell (at a concentration close to saturation), take on the role of the main charge carriers across the solution boundary. Due to the proximity of the mobilities of these ions and their predominant concentration, Dfo -> 0 mV. This is illustrated by columns 4 and 5 of Table. 4.2. The diffusion potentials at the boundaries of NaCl and KCl solutions with concentrated solutions of KS1 are really close to 0. At the same time, at the boundaries of concentrated solutions of KS1, even with dilute solutions of acid and alkali, D(p in 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 °C
|
Liquid connection" 1 |
exp. 6', |
Liquid connection a), d> |
||
|
ns1 o.1 :kci od |
HCI 1.0||KClSa, |
|||
|
NS1 0.1TsKS1 Sat |
||||
|
NS1 0.01TsKS1&, |
||||
|
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 0.1TsKS1 Sat |
|||
|
KCI 0.01 iNaCl 0.01 |
||||
|
KCI 0.01 iLiCl 0.01 |
NaOH 0.1 TsKS1 Sal |
|||
|
Kci o.oi :nh 4 ci o.oi |
NaOH 1.0TsKS1 Sat |
|||
|
LiCl 0.01:nh 4 ci 0.01 |
NaOH 1.0TsKS1 3.5 |
|||
|
LiCl 0.01 iNaCl 0.01 |
NaOH 0.1 TsKS1 0.1 |
Notes:
Concentrations in mol/l.
61 EMF measurements 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 KC1 (~4.16 mol/l).
"Calculation using the Henderson equation like (4.13), but using average activities instead of concentrations.
The 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 at which Df r is eliminated by the inclusion of an electrolytic bridge is usually denoted (||), as is done in Table. 4.2.
Addendum 4B.
DIFFUSION POTENTIAL,
potential difference at the boundary of two contacting solutions of electrolytes. It is due to the fact that the rates of transfer of cations and anions across the boundary, caused by the difference in their electrochemical properties. potentials in solutions 1 and 2 are different. The presence of a D. point can cause an error in measuring the electrode potential, so efforts are made to calculate or eliminate the D. point. Accurate calculation is impossible due to the uncertainty of the coefficient. ion activity, as well as the lack of information about the distribution of ion concentrations in the boundary zone between adjacent solutions. If solutions of the same z are in contact, z - charging electrolyte (z - number of cations equal to the number of anions) decomp. concentrations and we can assume that the transfer numbers of anions and cations, respectively. t + and t_ do not depend on their activity, but the coefficient. The activities of anions and cations are equal to each other in both solutions, then D. p.
Where a 1 and a 2 - average activities of ions in solutions 1 and 2, T - abs. t-ra, R - , F - Faraday's constant. There are other approximate formulas for determining D. p. Reduce D. p. to a small value in the plural. cases, it is possible by separating solutions 1 and 2 with a “salt bridge” from the concentrate. solutions, cations and cut have approximately equal numbers transfer (KCl, NH 4 NO 3, etc.). Lit.: Fetter K., Electrochemical kinetics, trans. from German, M., 1967, p. 70-76; Rotinyan A. L., Tikhonov K. I., Shoshina I. A., Theoretical. L., 1981, p. 131-35. A. D. Davydov.
Chemical encyclopedia. - M.: Soviet Encyclopedia. Ed. I. L. Knunyants. 1988 .
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Diffusion potentials arise at the interface between two solutions. Moreover, these can be either solutions of different substances or solutions of the same substance, only in the latter case they must differ from each other in their concentrations.
When two solutions come into contact, particles (ions) of dissolved substances interpenetrate into them due to the process of diffusion.
The reason for the emergence of a diffusion potential in this case is the unequal mobility of the ions of dissolved substances. If the electrolyte ions have different diffusion rates, then the faster ions gradually appear ahead of the less mobile ones. It is as if two waves of differently charged particles are formed.
If solutions of the same substance are mixed, but with different concentrations, then the more dilute solution acquires a charge that coincides in sign with the charge of more mobile ions, and the less diluted solution acquires 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 speeds: I– “fast” nones, negatively charged; II– “slow” ions, positively charged
A so-called diffusion potential arises at the solution interface. It averages the speed of movement of ions (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 cell membranes are damaged. In this case, their permeability is disrupted 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, damaged tissue is most often charged negatively in relation to undamaged tissue.
The diffusion potential arises in galvanic cells at the interface between two solutions. Therefore, when accurately calculating the emf. galvanic circuits must necessarily introduce a correction for its value. To eliminate the influence of diffusion potential, electrodes in galvanic cells are often connected to each other by a “salt bridge”, which is a saturated solutionKCl.
Potassium and chlorine ions have almost identical mobilities, so their use makes it possible to significantly reduce the influence of the diffusion potential on the emf value.
The diffusion potential can greatly increase if solutions of electrolytes of different compositions or different concentrations are separated by a membrane that is permeable only to 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 the exchange of ions 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 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.
Solid membrane electrodes contain a thin membrane on either side of which there are different solutions containing the same detectable ions, but at different concentrations. On the inside, the membrane is washed by a standard solution with a precisely known concentration of the ions being determined, and on the outside by the analyzed solution with an unknown concentration of the ions being determined.
Due to the different concentrations of solutions on both sides of the membrane, ions are exchanged differently with the inner and outer sides of the membrane. This leads to the formation of different electric charge and as a result of this, a membrane potential difference arises.
Among ion-selective electrodes, the glass electrode, which is used to determine the pH of solutions, has become widespread.
The central part of the glass electrode (Fig. 91) is a ball made of special conductive hydrated glass. It is filled with an aqueous solution of HCl with a known concentration (0.1 mol/dm 3). An electrode of the second type is placed in this solution - most often silver chloride, which acts as a reference electrode. During measurements, a glass bead is dipped into the solution being analyzed, which contains a second reference electrode.
The operating principle of the electrode is based on the fact that in the glass structure, K + , Na + , Li + ions are replaced by H + ions by long-term soaking in an acid solution. In this way, the glass membrane can exchange its H + ions with internal and external solutions (Fig. 92). Moreover, different potentials arise on both sides of the membrane as a result of this process.
Rice. 91. Scheme of a glass electrode: 1 – glass ball (membrane); 2 – internal solution of HC1; 3 – silver chloride electrode; 4 – measured solution; 5 – metal conductor
Rice. 92. Glass electrode in a solution with an unknown concentration of H + ions (a) and diagram of ion exchange between two phases (b)
Using reference electrodes placed in the external and internal solutions, their difference is measured.
The potential on the inner side of the membrane is constant, so the potential difference of the glass electrode will depend only on the activity of hydrogen ions in the test solution.
The general circuit diagram, including a glass electrode and two reference electrodes, is shown in Fig. 93.
Rice. 93. Circuit diagram explaining the principle of operation of a glass electrode
A glass electrode has a number of significant advantages over a hydrogen electrode, which can also be used to measure the concentration of H + ions in a solution.
It is completely insensitive to various impurities in the solution, “is not poisoned by them,” it can be used if the analyzed liquids contain strong oxidizing and reducing agents, as well as in the widest range of pH values - from 0 to 12. The disadvantage of the glass electrode is its large fragility.
Diffusion potential is the potential difference that occurs at the interface between two unequal electrolyte solutions. It is caused by the diffusion of ions across the interface and causes inhibition of faster diffusing ions and acceleration of slower diffusing ions, whether cations or anions. Thus, soon the equilibrium potential at the interface is established and reaches a constant value, which depends on the number of ion transfers, the magnitude of their charge and the concentration of the electrolyte.
E.m.f. concentration chain (see)
expressed by the equation
is the sum of two electrode potentials and the diffusion potential. The algebraic sum of two electrode potentials is theoretically equal to
hence,
Let's assume that then
or in general for an electrode reversible with respect to the cation,
and for an electrode reversible with respect to the anion,
For electrodes that are reversible with respect to the cation, when if then the value is positive and is added to the sum of the electrode potentials; if then the value is negative and e. d.s. element in this case is less than the sum of the electrode potentials. Attempts have been made to eliminate the diffusion potential by introducing a salt bridge containing concentrated solution and other salts for which. In this case, since the solution is concentrated, diffusion is determined by the electrolyte of the salt bridge itself, and instead of the diffusion potential of the cell, we have two diffusion potentials acting in opposite directions and having a value close to zero. In this way, it is possible to reduce diffusion potentials, but it is almost impossible to completely eliminate them.