Diffuse potential. Diffusion and membrane potentials
At the boundary of two unequal 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 diffusion potentials usually does not exceed several tens of millivolts, and they are usually not taken into account. However, with accurate measurements, special measures are taken to reduce them as much as possible. The reasons for the occurrence of the diffusion potential were shown using the example of two adjacent solutions of copper sulfate of different concentrations. Cu2+ and SO42- ions 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 solution interfaces on the side of the solution with a lower concentration, and an excess of Cu2+ appears on the more concentrated side. A potential difference arises. The presence of excess negative charge at the interface will inhibit the movement of SO42- and accelerate the movement of Cu2+. At a certain potential, the rates of SO42- and Cu2+ will become the same; a stationary value of the diffusion potential will be established. The theory of diffusion potential was developed by M. Planck (1890), and subsequently by A. Henderson (1907). The calculation formulas they obtained are complex. But the solution is simplified if the diffusion potential arises at the boundary of two solutions with different concentrations C1 and C2 of the same electrolyte. In this case, the diffusion potential is equal. Diffusion potentials arise during nonequilibrium diffusion processes, therefore they are irreversible. Their magnitude depends on the nature of the boundary of two contacting solutions, on the size and their configuration. Accurate measurements use techniques that minimize the magnitude of the diffusion potential. For this purpose, an intermediate solution with the lowest possible mobility values of U and V (for example, KCl and KNO3) is included between solutions in half-elements.
Diffusion potentials play an important role in biology. Their occurrence is not associated with metal electrodes. It is interfacial and diffusion potentials that generate biocurrents. For example, in electric stingrays and eels, a potential difference of up to 450 V is created. Biopotentials are sensitive to physiological changes in cells and organs. This is the basis for the use of electrocardiography and electroencephalography methods (measurement of the biocurrents of the heart and brain).
55. Interfluid phase potential, mechanism of occurrence and biological significance.
A potential difference also arises at the boundary of contact of immiscible liquids. Positive and negative ions in these solvents are distributed unevenly, and their distribution coefficients do not coincide. Therefore, a potential jump occurs at the interface between liquids, which prevents the unequal distribution of cations and anions in both solvents. In the total (total) volume of each phase, the number of cations and anions is almost the same. It will differ only at the phase interface. This is the interfluid potential. Diffusion and interfluid potentials play an important role in biology. Their occurrence is not associated with metal electrodes. It is interfacial and diffusion potentials that generate biocurrents. For example, in electric stingrays and eels, a potential difference of up to 450 V is created. Biopotentials are sensitive to physiological changes in cells and organs. This is the basis for the use of electrocardiography and electroencephalography methods (measurement of the biocurrents of the heart and brain).
As already indicated, concentration chains are of great practical importance, 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 phases of different chemical compositions, 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.
c) Diffusion potential of the 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 of electric current through electrochemical systems is associated with a chemical transformation, 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.
When creating any electrode pair, a “salt bridge” is always used. The use of a “salt bridge” solves several problems that arise for 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 (transfers) into a less concentrated solution. If the absolute speeds of movement of the cations and anions of the diffusing electrolyte are different, then the less concentrated solution acquires the potential of the charge sign of the “faster ions,” 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 purpose, a saturated KCl solution was chosen, because absolute movement speeds K + and Cl ¯ almost identical and have one of the highest values.
The emergence of a diffusion potential is also characteristic of biological systems. For example, when a cell is damaged, when the semi-permeability of its membrane is disrupted, electrolyte begins to diffuse into or out of the cell. This creates a diffusion potential, which is called “damage potential” here. Its value can reach 30 - 40 mV, the “damage potential” is stable for approximately one hour.
The value of the diffusion potential increases significantly if solutions of electrolytes 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 to form electrode pairs (elements).
There are two main types of classification of electrodes.
By chemical composition:
1. Electrodes of the 1st kind - these are electrodes whose electrode reaction is reversible only with respect to the cation or anion. For example, the electrodes that form the Jacobi-Daniel element are copper and zinc (see above).
2. Electrodes 2 types - these are electrodes whose electrode reaction is reversible for two types of ions: cations and anions.
3. Redox electrodes (Red – Ox) . The term “Red – Ox – electrode” refers to an electrode where all the elements of the half-reaction (both oxidized and reduced forms) are in solution. Metal electrodes immersed in the solution do not participate in the reaction, but serve only as a carrier of electrons.
By purpose:
1. Reference electrodes .
Reference electrodes are electrodes whose potential is precisely known, is stable over time and does not depend on the concentration of ions in the solution. Such electrodes include: standard hydrogen electrode, calomel electrode and silver chloride electrode. Let's look at 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 1 mol/l. Hydrogen gas is passed into a vessel under a pressure of 1 atmosphere. Hydrogen bubbles are adsorbed on a platinum plate, where they are dissociated into atomic hydrogen and oxidized.
Characteristics of standard hydrogen electrode:
1.Electrode diagram: Pt(H 2) / H +
2. Electrode reaction: ½ Н 2 – ē ↔ Н +
As is easy to see, this reaction is reversible only for the cation (H +), therefore the standard hydrogen electrode is a type 1 electrode.
3.Calculation of electrode potential.
The Nernst equation takes the form:
eH 2 /H+ = e ° N 2 /N + RT ln a n +
nF (P n 2) 1/2
Because and n+ = 1 mol/l, р n+ = 1 atm, then ln a n+ = 0, That's why
(Rn 2) 1/2
eH 2 /H+ = e ° N 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 including calomel (Hg 2 Cl 2), mercury and potassium chloride. The paste is based on pure mercury and filled with a solution of potassium chloride. A platinum plate is immersed inside this system.
Electrode characteristics:
1. Electrode diagram: Hg 2 Cl 2, Hg(Pt) / Cl¯
2. Two parallel reactions occur in this electrode:
Hg 2 Cl 2 ↔2Hg + +2Cl¯
2 Hg + + 2ē →2Hg
Hg 2 Cl 2 + 2ē → 2Hg +2Cl¯ - total reaction.
From the above equations it is clear that the calomel electrode is a type 2 electrode.
3. The electrode potential is determined using 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 pic).
Here, the silver wire is coated with a layer of poorly soluble salt AgCl 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 resulting metal settles on the wire, and 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 electrode potential is determined using 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 maintained constant and therefore their electrode potentials are known and constant over time.
2. Definition electrodes - these are electrodes whose potential 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 is designed similarly to a standard hydrogen electrode, but if an acidic solution with an activity of H + ions greater than one is placed in the container of the hydrogen electrode, then a positive potential appears on the electrode, proportional to the activity (i.e., concentration) of protons. When the proton concentration decreases, on the contrary, the electrode will be negatively charged. Therefore, by determining the potential of such an electrode, it is possible to calculate the pH of the solution in which it is immersed.
Electrode characteristics.
1. Electrode diagram: Pt(H 2) / H +
2. Electrode reaction: ½ Н 2 – ē ↔ Н +
3. e H 2 /H+ = e o H 2 /H + + 0.059 lg a n+
n
Because n =1, and e o N 2 / H+= 0, then the Nernst equation takes the form:
e H2/H+ = 0.059 lg a n+ = - 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 solution of hydrochloric acid. The electrode potential depends on the concentration of H + and is determined by the Nernst equation, which has the form:
e st = e o st + 0.059 lg a n +
Quinhydrone 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 quickly established:
Since protons are involved in this reaction, the electrode potential depends on pH.
Electrode characteristics:
1. Electrode diagram: Pt / H +, C 6 H 4 O 2, C 6 H 4 O 2-
2. Electrode reaction:
C 6 H 4 (OH) 2 - 2ē ↔ C 6 H 4 O 2 + 2H + -
redox process.
3. The electrode potential is determined using the Nernst equation, which after appropriate transformations takes the form:
e x. g. = e o x. g. + 0.059 lg a H +
The quinhydrone electrode is used only to determine the pH of those solutions where this indicator is not more than 8. This is due to the fact that in an alkaline environment hydroquinone behaves like an acid and the value of the electrode potential ceases to depend on the concentration of protons.
Because in a quinhydrone electrode A plate of noble metal is immersed in a solution containing both the oxidized and reduced forms 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 substances, "red - ox" - electrodes are especially important because are the only way to form an electrode and determine its potential.
The magnitude of the 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 ;(IN)
T– temperature, 0 K.
C ox And C red– concentrations 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.
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. Eliminating diffusion potential using electrolytic bridges
Generally speaking, diffusion potentials at the interface of 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 parentheses):
In this regard, the diffusion potential must either be eliminated or accurately measured. Elimination of the diffusion potential is achieved by including an additional electrolyte with similar cation and anion mobilities into the electrochemical system. When making measurements in aqueous solutions, saturated solutions of potassium chloride, potassium nitrate or ammonium are used as such electrolytes.
An additional electrolyte is included between the main electrolytes using electrolytic bridges (Fig. 6.12) filled with the main 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 transferred by K + and C1 – ions, which are much more numerous than the ions of the main electrolyte. Since the mobilities of K + and C1 – ions in potassium chloride are almost equal to each other, the diffusion potential will 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. Thus, 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 equal to 16.9 mV. If additional electrolytes are included between 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 (saturated) 2.8
KCl (saturated) 1.5
Elimination of diffusion potentials by including an additional electrolyte with equal ion transfer numbers gives good results when measuring diffusion potentials in non-concentrated solutions with slightly different mobilities of the anion and cation. When measuring voltages of systems containing solutions of acids or alkalis
Table 6.3. Diffusion potentials at the interface of KOH – KCl and NaOH – KCl (according to V. G. Lokshtanov)
with very different rates of movement of the cation and anion, special care must be taken. For example, at the HC1 - KS1 (saturated) 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). Thus, the diffusion potential, for example in a 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 cannot be significantly reduced using a salt bridge.
To reduce diffusion potentials to negligible values, Nernst proposed adding a large excess of some indifferent electrolyte 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 distorted results. Therefore, this method can only be used in those
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 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, since 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 contact boundary between the two solutions. A continuously renewed boundary is created by slow directed movement of solutions parallel to each other. In this way, it is possible to achieve stability of the diffusion potential and its reproducibility with an accuracy of 0.1 mV.
The diffusion potential is determined by the Cohen and Tombrock method from voltage measurements of two electrochemical systems, the electrodes of one of them being reversible to the salt cation, and the other to the anion. Let's say we 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 (solid), ZnSO 4 | ZnSO 4, Hg 2 SO 4 (solid) | Hg(+)
System 1 voltage
systems 2
Considering that φ d 21 = – φ d 12, and subtracting the second equation from the first, we obtain:
When measurements are carried out at not very high concentrations, at which one can still assume that = and = or that : = : the last two terms of the last equation cancel and
The diffusion potential in system 1 can also be determined in a slightly different way, if instead of system 2 we use a dual electrochemical system:
3. (–) Zn | ZnSO 4, Hg 2 SO 4 (solid) | Hg - Hg | Hg 2 SO 4 (solid), ZnSO 4 | Zn(+)
System voltage
Therefore, the voltage difference between systems 1 and 3 will be expressed by the equation:
If, as before, the ratio of the activities of zinc ions is replaced by the ratio of the average ionic activities of the zinc salt, we obtain:
Since the last term of this equation can usually be calculated accurately, the value of the diffusion potential can be determined from the measurements of E p1 and E p 3.
The diffusion potential at the boundary of two different solutions is determined in a similar way. For example, if they want to determine the diffusion potential at the boundary of solutions of zinc sulfate and copper chloride, they create two electrochemical systems:
4. (–) Zn | ZnSO 4 | CuCl2 | Cu(+)
5. (–) Hg | Hg 2 Cl 2 (solid), CuCl 2 | ZnSO 4, Hg 2 SO 4 (solid) | Hg(+)
System voltage 4
systems 5
Hence
Naturally, the greater the number of terms included in the equation for the diffusion potential, the less likely the determination is to be highly accurate.
Related information.
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 transfer numbers (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 .
See what "DIFFUSION POTENTIAL" is in other dictionaries:
diffusion potential- – potential arising in a galvanic cell upon contact of electrolytes; due to different rates of ion diffusion. General chemistry: textbook / A. V. Zholnin ... Chemical terms
diffusion potential- - [L.G. Sumenko. English-Russian dictionary on information technology. M.: State Enterprise TsNIIS, 2003.] Topics information technology in general EN diffusion potential...
diffusion potential- difuzijos potencialas statusas T sritis Standartizacija ir metrologija apibrėžtis Potencialo pokytis, susidarantis dėl koncentracijų skirtumo kietųjų kūnų, tirpalų ir pan. sąlyčio riboje. atitikmenys: engl. diffusion potential vok.… … Penkiakalbis aiškinamasis metrologijos terminų žodynas
diffusion potential- difuzinis potencialas statusas T sritis chemija apibrėžtis Potencialo šuolis, atsirandantis tirpalų sąlyčio riboje. atitikmenys: engl. diffusion potential rus. diffusion potential... Chemijos terminų aiškinamasis žodynas
diffusion potential- difuzijos potencialas statusas T sritis fizika atitikmenys: engl. diffusion potential vok. Diffusion potential, n rus. diffusion potential, m pranc. potentiel de diffusion, m … Fizikos terminų žodynas
diffusion potential of spontaneous polarization in a well- diffusion potential Ed Electromotive force arising at the contact of solutions of different mineralization. [GOST 22609 77] Topics: geophysical research in wells General terms: processing and interpretation of geophysical results... ... Technical Translator's Guide
POTENTIAL- POTENTIAL. The amount of any type of energy can be expressed as the product of two different quantities, of which one characterizes the “level of energy” and determines the direction in which its transition should take place; so eg heavy body... ... Great Medical Encyclopedia
potential of spontaneous polarization in a well- spontaneous polarization potential Ups Potential created in the well by spontaneous polarization currents. Note The spontaneous polarization potential includes diffusion, diffusion absorption and filtration potentials. [GOST... ... Technical Translator's Guide
wave potential- – in classical polarography, the potential at which the diffusion current caused by the reduction of the substance on the indicator electrode reaches half of its maximum value. Dictionary of Analytical Chemistry... Chemical terms
Difference electrostatic potentials between the electrode and the electrolyte in contact with it. The emergence of electrical energy is due to space. separation of charges of opposite sign at the phase boundary and the formation of double... ... Chemical encyclopedia
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