Electrochemical activity series of metals (voltage range, range of standard electrode potentials) - sequence in which metals are arranged in order of increasing their standard electrochemical potentials φ 0, corresponding to the half-reaction of reduction of the metal cation Me n+: Me n+ + nē → Me
A number of voltages characterize the comparative activity of metals in redox reactions in aqueous solutions.
Story
The sequence of metals in the order of changes in their chemical activity in general outline was already known to alchemists. The processes of mutual displacement of metals from solutions and their surface deposition (for example, the displacement of silver and copper from solutions of their salts by iron) were considered as a manifestation of the transmutation of elements.
Later alchemists came close to understanding the chemical side of the mutual precipitation of metals from their solutions. Thus, Angelus Sala in his work “Anatomia Vitrioli” (1613) came to the conclusion that products chemical reactions consist of the same “components” that were contained in the original substances. Subsequently, Robert Boyle proposed a hypothesis about the reasons why one metal displaces another from solution based on corpuscular concepts.
In the era of the emergence of classical chemistry, the ability of elements to displace each other from compounds became an important aspect of understanding reactivity. J. Berzelius, based on the electrochemical theory of affinity, built a classification of elements, dividing them into “metalloids” (the term “non-metals” is now used) and “metals” and placing hydrogen between them.
The sequence of metals according to their ability to displace each other, long known to chemists, was especially thoroughly and comprehensively studied and supplemented by N. N. Beketov in the 1860s and subsequent years. Already in 1859, he made a report in Paris on the topic “Investigation of the phenomena of the displacement of some elements by others.” In this work, Beketov included a number of generalizations about the relationship between the mutual displacement of elements and their atomic weight, connecting these processes with “ the original chemical properties of elements - what is called chemical affinity". Beketov's discovery of the displacement of metals from solutions of their salts by hydrogen under pressure and the study of the reducing activity of aluminum, magnesium and zinc at high temperatures (metallothermy) allowed him to put forward a hypothesis about the connection between the ability of some elements to displace others from compounds with their density: lighter simple substances are able to displace more heavy (therefore this series often also called Beketov's displacement series, or just Beketov series).
Without denying Beketov’s significant merits in the development of modern ideas about the activity series of metals, the idea of him as the only creator of this series, existing in domestic popular and educational literature, should be considered erroneous. Numerous experimental data obtained in late XIX centuries, refuted Beketov’s hypothesis. Thus, William Odling described many cases of “reversal of activity.” For example, copper displaces tin from a concentrated acidified solution of SnCl 2 and lead from an acidic solution of PbCl 2 ; it is also capable of dissolving in concentrated hydrochloric acid with the release of hydrogen. Copper, tin and lead are in the series to the right of cadmium, but can displace it from a boiling, slightly acidified solution of CdCl 2.
The rapid development of theoretical and experimental physical chemistry pointed to another reason for the differences in the chemical activity of metals. With the development of modern concepts of electrochemistry (mainly in the works of Walter Nernst), it became clear that this sequence corresponds to the “series of voltages” - the arrangement of metals according to the value of standard electrode potentials. So instead of qualitative characteristics- the “propensity” of a metal and its ion to certain reactions - Nerst introduced an exact quantitative value characterizing the ability of each metal to pass into solution in the form of ions, as well as to be reduced from ions to the metal on the electrode, and the corresponding series was called range of standard electrode potentials.
Theoretical foundations
The values of electrochemical potentials are a function of many variables and therefore exhibit a complex dependence on the position of metals in the periodic table. Thus, the oxidation potential of cations increases with an increase in the atomization energy of the metal, with an increase in the total ionization potential of its atoms, and with a decrease in the hydration energy of its cations.
In the very general view It is clear that metals located at the beginning of the periods are characterized by low values of electrochemical potentials and occupy places on the left side of the voltage series. In this case, the alternation of alkali and alkaline earth metals reflects the phenomenon of diagonal similarity. Metals located closer to the middle of the periods are characterized by large potential values and occupy places in the right half of the row. A consistent increase in the electrochemical potential (from −3.395 V for the Eu 2+ /Eu [ ] to +1.691 V for the Au + /Au pair) reflects a decrease in the reducing activity of metals (the ability to donate electrons) and an increase in the oxidizing ability of their cations (the ability to gain electrons). Thus, the strongest reducing agent is metallic europium, and the strongest oxidizing agent is the gold cations Au+.
Hydrogen is traditionally included in the voltage series, since practical measurement of electrochemical potentials of metals is made using a standard hydrogen electrode.
Practical use of a range of voltages
A number of voltages are used in practice for comparative [relative] assessment of the chemical activity of metals in reactions with aqueous solutions of salts and acids and for assessment of cathodic and anodic processes during electrolysis:
- Metals to the left of hydrogen are stronger reducing agents than metals to the right: they displace the latter from salt solutions. For example, the interaction Zn + Cu 2+ → Zn 2+ + Cu is possible only in the forward direction.
- Metals in the row to the left of hydrogen displace hydrogen when interacting with aqueous solutions of non-oxidizing acids; most active metals(up to aluminum inclusive) - and when interacting with water.
- Metals in the series to the right of hydrogen do not interact with aqueous solutions of non-oxidizing acids under normal conditions.
- During electrolysis, metals to the right of hydrogen are released at the cathode; the reduction of moderately active metals is accompanied by the release of hydrogen; The most active metals (up to aluminum) cannot be isolated from aqueous salt solutions under normal conditions.
Table of electrochemical potentials of metals
| Metal | Cation | φ 0, V | Reactivity | Electrolysis (at the cathode): |
|---|---|---|---|---|
| Li+ | -3,0401 | reacts with water | hydrogen is released | |
| Cs+ | -3,026 | |||
| Rb+ | -2,98 | |||
| K+ | -2,931 | |||
| Fr+ | -2,92 | |||
| Ra 2+ | -2,912 | |||
| Ba 2+ | -2,905 | |||
| Sr 2+ | -2,899 | |||
| Ca2+ | -2,868 | |||
| Eu 2+ | -2,812 | |||
| Na+ | -2,71 | |||
| Sm 2+ | -2,68 | |||
| Md 2+ | -2,40 | reacts with aqueous solutions of acids | ||
| La 3+ | -2,379 | |||
| Y 3+ | -2,372 | |||
| Mg 2+ | -2,372 | |||
| Ce 3+ | -2,336 | |||
| Pr 3+ | -2,353 | |||
| Nd 3+ | -2,323 | |||
| Er 3+ | -2,331 | |||
| Ho 3+ | -2,33 | |||
| Tm 3+ | -2,319 | |||
| Sm 3+ | -2,304 | |||
| PM 3+ | -2,30 | |||
| Fm 2+ | -2,30 | |||
| Dy 3+ | -2,295 | |||
| Lu 3+ | -2,28 | |||
| Tb 3+ | -2,28 | |||
| Gd 3+ | -2,279 | |||
| Es 2+ | -2,23 | |||
| Ac 3+ | -2,20 | |||
| Dy 2+ | -2,2 | |||
| PM 2+ | -2,2 | |||
| Cf 2+ | -2,12 | |||
| Sc 3+ | -2,077 | |||
| Am 3+ | -2,048 | |||
| Cm 3+ | -2,04 | |||
| Pu 3+ | -2,031 | |||
| Er 2+ | -2,0 | |||
| Pr 2+ | -2,0 | |||
| Eu 3+ | -1,991 | |||
| Lr 3+ | -1,96 | |||
| Cf 3+ | -1,94 | |||
| Es 3+ | -1,91 | |||
| Th 4+ | -1,899 | |||
| Fm 3+ | -1,89 | |||
| Np 3+ | -1,856 | |||
| Be 2+ | -1,847 | |||
| U 3+ | -1,798 | |||
| Al 3+ | -1,700 | |||
| MD 3+ | -1,65 | |||
| Ti 2+ | -1,63 | competing reactions: both the release of hydrogen and the release of pure metal | ||
| Hf 4+ | -1,55 | |||
| Zr 4+ | -1,53 | |||
| Pa 3+ | -1,34 | |||
| Ti 3+ | -1,208 | |||
| Yb 3+ | -1,205 | |||
| No 3+ | -1,20 | |||
| Ti 4+ | -1,19 | |||
| Mn 2+ | -1,185 | |||
| V 2+ | -1,175 | |||
| Nb 3+ | -1,1 | |||
| Nb 5+ | -0,96 | |||
| V 3+ | -0,87 | |||
| Cr 2+ | -0,852 | |||
| Zn 2+ | -0,763 | |||
| Cr 3+ | -0,74 | |||
| Ga 3+ | -0,560 | |||
Metal stress range- this is a series of metals arranged in increasing order of their standard electrode potential (). The position of a metal in the voltage series indicates its redox abilities in relation to other metals and their cations for reactions occurring in electrolyte solutions, i.e., in reactions with salts and bases. And also with non-metals, if these reactions occur in aqueous solutions; in particular, such processes include the processes of corrosion of metals ().
In the series of voltages:
1) The reducing ability of metals decreases.
2) Oxidizing capacity increases. As a consequence of this, metals that are in the voltage series before hydrogen displace it from solutions of acids (not oxidizing agents).
3) Metals standing to the left in the series (having a lower potential) displace metals standing to the right (having a higher potential) from solutions of their salts.
4) Metals in the voltage range up to Mg (having ) displace hydrogen from water.
Thus, the value of the electrode potential determines the redox abilities of metals in relation to each other and in relation to H and the cations containing it in electrolytes.
Measurement of electrode potentials. A range of standard electrode potentials, hydrogen electrode.
The absolute value of the electrode potential is almost impossible to measure. In this regard, the electrode potential is measured by measuring the EMF of a galvanic cell composed of the electrode under study and the electrode potential, of which the potential is known. The standard electrode potential is determined by the value of the emf of a galvanic cell composed of the electrode under study and a standard hydrogen electrode, the potential of which is conventionally assumed to be zero.
Standard hydrogen electrode– This is a system located at normal conditions, consisting of a sponge plate into the pores of which hydrogen is pumped, placed in a one-molal solution of sulfuric acid H 2 SO 4 with C(H +) = 1 mol/kg
Standardizing the conditions and reproducing the potential of such an electrode is a difficult task, so this electrode is used for meteorological purposes. In laboratory practice, auxiliary electrodes are used to measure electrode potentials.
Example: calomel electrode - Hg,HgCl/Cl - ;
silver chlorine – Ag, AgCl/Cl - etc.
The potential of these electrodes is stably reproduced, that is, it retains its value during storage and operation.
What information can be obtained from a series of voltages?
A range of metal voltages are widely used in inorganic chemistry. In particular, the results of many reactions and even the possibility of their implementation depend on the position of a certain metal in the NER. Let's discuss this issue in more detail.
Interaction of metals with acids
Metals located in the voltage series to the left of hydrogen react with acids - non-oxidizing agents. Metals located in the ERN to the right of H interact only with oxidizing acids (in particular, with HNO 3 and concentrated H 2 SO 4).
Example 1. Zinc is located in the NER to the left of hydrogen, therefore, it is able to react with almost all acids:
Zn + 2HCl = ZnCl 2 + H 2
Zn + H 2 SO 4 = ZnSO 4 + H 2
Example 2. Copper is located in the ERN to the right of H; this metal does not react with “ordinary” acids (HCl, H 3 PO 4, HBr, organic acids), however, it interacts with oxidizing acids (nitric, concentrated sulfuric):
Cu + 4HNO 3 (conc.) = Cu(NO 3) 2 + 2NO 2 + 2H 2 O
Cu + 2H 2 SO 4 (conc.) = CuSO 4 + SO 2 + 2H 2 O
I would like to draw your attention to an important point: when metals interact with oxidizing acids, it is not hydrogen that is released, but some other compounds. You can read more about this!
Interaction of metals with water
Metals located in the voltage series to the left of Mg readily react with water already at room temperature with the release of hydrogen and the formation of an alkali solution.
Example 3. Sodium, potassium, calcium easily dissolve in water to form an alkali solution:
2Na + 2H 2 O = 2NaOH + H 2
2K + 2H 2 O = 2KOH + H 2
Ca + 2H 2 O = Ca(OH) 2 + H 2
Metals located in the voltage range from hydrogen to magnesium (inclusive) in some cases interact with water, but the reactions require specific conditions. For example, aluminum and magnesium begin to interact with H 2 O only after removing the oxide film from the metal surface. Iron does not react with water at room temperature, but does react with water vapor. Cobalt, nickel, tin, and lead practically do not interact with H 2 O, not only at room temperature, but also when heated.
The metals located on the right side of the ERN (silver, gold, platinum) do not react with water under any conditions.
Interaction of metals with aqueous solutions of salts
We will talk about reactions of the following type:
metal (*) + metal salt (**) = metal (**) + metal salt (*)
I would like to emphasize that the asterisks in this case do not indicate the oxidation state or the valency of the metal, but simply allow one to distinguish between metal No. 1 and metal No. 2.
To carry out such a reaction, three conditions must be met simultaneously:
- the salts involved in the process must be dissolved in water (this can be easily checked using the solubility table);
- the metal (*) must be in the stress series to the left of the metal (**);
- the metal (*) should not react with water (which is also easily verified by ESI).
Example 4. Let's look at a few reactions:
Zn + CuSO 4 = ZnSO 4 + Cu
K + Ni(NO 3) 2 ≠
The first reaction is easily feasible, all the above conditions are met: copper sulfate is soluble in water, zinc is in the NER to the left of copper, Zn does not react with water.
The second reaction is impossible because the first condition is not met (copper (II) sulfide is practically insoluble in water). The third reaction is not feasible, since lead is a less active metal than iron (located to the right in the ESR). Finally, the fourth process will NOT result in nickel precipitation because potassium reacts with water; the resulting potassium hydroxide can react with the salt solution, but this is a completely different process.
Thermal decomposition process of nitrates
Let me remind you that nitrates are salts nitric acid. All nitrates decompose when heated, but the composition of the decomposition products may vary. The composition is determined by the position of the metal in the stress series.
Nitrates of metals located in the NER to the left of magnesium, when heated, form the corresponding nitrite and oxygen:
2KNO 3 = 2KNO 2 + O 2
During the thermal decomposition of metal nitrates located in the voltage range from Mg to Cu inclusive, metal oxide, NO 2 and oxygen are formed:
2Cu(NO 3) 2 = 2CuO + 4NO 2 + O 2
Finally, during the decomposition of nitrates of the least active metals (located in the ERN to the right of copper), metal, nitrogen dioxide and oxygen are formed.
If from the entire series of standard electrode potentials we select only those electrode processes that correspond to the general equation
then we get a series of metal stresses. In addition to metals, this series will always include hydrogen, which allows you to see which metals are capable of displacing hydrogen from aqueous solutions of acids.
Table 19. Series of metal stresses
A number of stresses for the most important metals are given in table. 19. The position of a particular metal in the stress series characterizes its ability to undergo redox interactions in aqueous solutions under standard conditions. Metal ions are oxidizing agents, and metals in the form of simple substances are reducing agents. Moreover, the further a metal is located in the voltage series, the stronger the oxidizing agent in an aqueous solution are its ions, and vice versa, the closer the metal is to the beginning of the series, the stronger the reducing properties of a simple substance - the metal.
Electrode process potential
in a neutral environment it is equal to B (see page 273). Active metals at the beginning of the series, having a potential significantly more negative than -0.41 V, displace hydrogen from water. Magnesium displaces hydrogen only from hot water. Metals located between magnesium and cadmium generally do not displace hydrogen from water. Oxide films are formed on the surface of these metals, which have a protective effect.
Metals located between magnesium and hydrogen displace hydrogen from acid solutions. At the same time, protective films are also formed on the surface of some metals, inhibiting the reaction. Thus, the oxide film on aluminum makes this metal stable not only in water, but also in solutions of certain acids. Lead does not dissolve in sulfuric acid at its concentration below, since the salt formed when lead reacts with sulfuric acid is insoluble and creates a protective film on the metal surface. The phenomenon of deep inhibition of metal oxidation, due to the presence of protective oxide or salt films on its surface, is called passivity, and the state of the metal in this case is called a passive state.
Metals are capable of displacing each other from salt solutions. The direction of the reaction is determined by their relative position in the series of stresses. When considering specific cases of such reactions, it should be remembered that active metals displace hydrogen not only from water, but also from any aqueous solution. Therefore, the mutual displacement of metals from solutions of their salts practically occurs only in the case of metals located in the series after magnesium.
Beketov was the first to study in detail the displacement of metals from their compounds by other metals. As a result of his work, he arranged metals according to their chemical activity into a displacement series, which is the prototype of a series of metal stresses.
The relative position of some metals in the stress series and in the periodic table at first glance does not correspond to each other. For example, according to the position in the periodic table, the chemical activity of potassium should be greater than sodium, and sodium - greater than lithium. In the series of voltages, lithium is the most active, and potassium occupies a middle position between lithium and sodium. Zinc and copper, according to their position in the periodic table, should have approximately equal chemical activity, but in the voltage series, zinc is located much earlier than copper. The reason for this kind of inconsistency is as follows.
When comparing metals occupying one or another position in the periodic table, the ionization energy of free atoms is taken as a measure of their chemical activity - reducing ability. Indeed, when moving, for example, from top to bottom along main subgroup Group I periodic table the ionization energy of atoms decreases, which is associated with an increase in their radii (i.e., with a greater distance of outer electrons from the nucleus) and with increasing screening of the positive charge of the nucleus by intermediate electronic layers (see § 31). Therefore, potassium atoms exhibit greater chemical activity - they have stronger reducing properties - than sodium atoms, and sodium atoms exhibit greater activity than lithium atoms.
When comparing metals in a series of voltages, the work of converting a metal in a solid state into hydrated ions in an aqueous solution is taken as a measure of chemical activity. This work can be represented as the sum of three terms: the atomization energy - the transformation of a metal crystal into isolated atoms, the ionization energy of free metal atoms and the hydration energy of the resulting ions. Atomization energy characterizes the strength of the crystal lattice of a given metal. The energy of ionization of atoms - the removal of valence electrons from them - is directly determined by the position of the metal in the periodic table. The energy released during hydration depends on the electronic structure of the ion, its charge and radius.
Lithium and potassium ions, having the same charge but different radii, will create unequal electric fields. The field generated near small lithium ions will be stronger than the field near large potassium ions. It is clear from this that lithium ions will hydrate with the release of more energy than potassium ions.
Thus, during the transformation under consideration, energy is expended on atomization and ionization and energy is released during hydration. The lower the total energy consumption, the easier the entire process will be and the closer to the beginning of the stress series the given metal will be located. But of the three terms of the general energy balance, only one - the ionization energy - is directly determined by the position of the metal in the periodic table. Consequently, there is no reason to expect that the relative position of certain metals in the stress series will always correspond to their position in the periodic table. Thus, for lithium, the total energy consumption turns out to be less than for potassium, according to which lithium comes before potassium in the voltage series.
For copper and zinc, the energy expenditure for the ionization of free atoms and the energy gain during ion hydration are close. But metallic copper forms a stronger crystal lattice, than zinc, as can be seen from a comparison of the melting temperatures of these metals: zinc melts at , and copper only at . Therefore, the energy spent on the atomization of these metals is significantly different, as a result of which the total energy costs for the entire process in the case of copper are much greater than in the case of zinc, which explains the relative position of these metals in the stress series.
When passing from water to non-aqueous solvents, the relative positions of metals in the voltage series may change. The reason for this is that the solvation energy of different metal ions changes differently when moving from one solvent to another.
In particular, the copper ion is solvated quite vigorously in some organic solvents; This leads to the fact that in such solvents copper is located in the voltage series before hydrogen and displaces it from acid solutions.
Thus, unlike the periodic table of elements, the stress series of metals is not a reflection general patterns, on the basis of which it is possible to give versatile Characteristics chemical properties metals A series of voltages characterizes only the redox ability of the Electrochemical system “metal - metal ion” under strictly defined conditions: the values given in it refer to aqueous solution, temperature and unit concentration (activity) of metal ions.
In an electrochemical cell (galvanic cell), the electrons remaining after the formation of ions are removed through a metal wire and recombine with ions of another type. That is, the charge in the external circuit is transferred by electrons, and inside the cell, through the electrolyte in which the metal electrodes are immersed, by ions. This creates a closed electrical circuit.
The potential difference measured in an electrochemical cell is o is explained by the difference in the ability of each metal to donate electrons. Each electrode has its own potential, each electrode-electrolyte system is a half-cell, and any two half-cells form an electrochemical cell. The potential of one electrode is called the half-cell potential, and it determines the ability of the electrode to donate electrons. It is obvious that the potential of each half-element does not depend on the presence of another half-element and its potential. The half-cell potential is determined by the concentration of ions in the electrolyte and temperature.
Hydrogen was chosen as the “zero” half-element, i.e. it is believed that no work is done for it when an electron is added or removed to form an ion. The “zero” potential value is necessary to understand the relative ability of each of the two half-cells of the cell to give and accept electrons.
Half-cell potentials measured relative to a hydrogen electrode are called the hydrogen scale. If the thermodynamic tendency to donate electrons in one half of the electrochemical cell is higher than in the other, then the potential of the first half-cell is higher than the potential of the second. Under the influence of the potential difference, electron flow will occur. When two metals are combined, it is possible to determine the potential difference that arises between them and the direction of electron flow.
An electropositive metal has a higher ability to accept electrons, so it will be cathodic or noble. On the other side are electronegative metals, which are capable of spontaneously donating electrons. These metals are reactive and therefore anodic:
- → 0 → +
Al Mn Zn Fe Sn Pb H 2 Cu Ag Au
For example Cu gives up electrons more easily Ag, but worse than Fe . In the presence of a copper electrode, silver nonions will begin to combine with electrons, resulting in the formation of copper ions and the precipitation of metallic silver:
2 Ag + + Cu → Cu 2+ + 2 Ag
However, the same copper is less reactive than iron. When metallic iron comes into contact with copper nonates, it will precipitate and the iron will go into solution:
Fe + Cu 2+ → Fe 2+ + Cu.
We can say that copper is a cathode metal relative to iron and an anodic metal relative to silver.
The standard electrode potential is considered to be the potential of a half-cell of fully annealed pure metal as an electrode in contact with ions at 25 0 C. In these measurements, the hydrogen electrode acts as a reference electrode. In the case of a divalent metal, we can write down the reaction occurring in the corresponding electrochemical cell:
M + 2H +→ M 2+ + H 2.
If we arrange metals in descending order of their standard electrode potentials, we obtain the so-called electrochemical series of metal voltages (Table 1).
Table 1. Electrochemical series of metal voltages
|
Metal-ion equilibrium (unit activity) |
Electrode potential relative to the hydrogen electrode at 25°C, V (reduction potential) |
|
|
Noble or cathode |
Au-Au 3+ |
1,498 |
|
Pt-Pt 2+ |
||
|
Pd-Pd 2+ |
0,987 |
|
|
Ag-Ag+ |
0,799 |
|
|
Hg-Hg 2+ |
0,788 |
|
|
Cu-Cu 2+ |
0,337 |
|
|
H 2 -H + |
||
|
Pb-Pb 2+ |
0,126 |
|
|
Sn-Sn 2+ |
0,140 |
|
|
Ni-Ni 2+ |
0,236 |
|
|
Co-Co 2+ |
0,250 |
|
|
Cd-Cd 2+ |
0,403 |
|
|
Fe-Fe 2+ |
0,444 |
|
|
Cr-Cr 2+ |
0,744 |
|
|
Zn-Zn 2+ |
0,763 |
|
|
Active |
Al-Al 2+ |
1,662 |
|
Mg-Mg2+ |
2,363 |
|
|
Na-Na+ |
2,714 |
|
|
K-K+ |
2,925 |
For example, in a copper-zinc galvanic cell, there is a flow of electrons from zinc to copper. The copper electrode is the positive pole in this circuit, and the zinc electrode is the negative pole. The more reactive zinc loses electrons:
Zn → Zn 2+ + 2е - ; E °=+0.763 V.
Copper is less reactive and accepts electrons from zinc:
Cu 2+ + 2e - → Cu; E °=+0.337 V.
The voltage on the metal wire connecting the electrodes will be:
0.763 V + 0.337 V = 1.1 V.
Table 2. Stationary potentials of some metals and alloys in sea water in relation to a normal hydrogen electrode (GOST 9.005-72).
|
Metal |
Stationary potential, IN |
Metal |
Stationary potential, IN |
|
Magnesium |
1,45 |
Nickel (active co standing) |
0,12 |
|
Magnesium alloy (6% A l, 3 % Zn, 0,5 % Mn) |
1,20 |
Copper alloys LMtsZh-55 3-1 |
0,12 |
|
Zinc |
0,80 |
Brass (30 % Zn) |
0,11 |
|
Aluminum alloy (10% Mn) |
0,74 |
Bronze (5-10 % Al) |
0,10 |
|
Aluminum alloy (10% Zn) |
0,70 |
Red brass (5-10 % Zn) |
0,08 |
|
Aluminum alloy K48-1 |
0,660 |
Copper |
0,08 |
|
Aluminum alloy B48-4 |
0,650 |
Cupronickel (30% Ni) |
0,02 |
|
Aluminum alloy AMg5 |
0,550 |
Bronze "Neva" |
0,01 |
|
Aluminum alloy AMg61 |
0,540 |
Bronze Br. AZHN 9-4-4 |
0,02 |
|
Aluminum |
0,53 |
Stainless steel X13 (passive state) |
0,03 |
|
Cadmium |
0,52 |
Nickel (passive state) |
0,05 |
|
Duralumin and aluminum alloy AMg6 |
0,50 |
Stainless steel X17 (passive state) |
0,10 |
|
Iron |
0,50 |
Titan technical |
0,10 |
|
Steel 45G17Yu3 |
0,47 |
Silver |
0,12 |
|
Steel St4S |
0,46 |
Stainless steel 1X14ND |
0,12 |
|
Steel SHL4 |
0,45 |
Titanium iodide |
0,15 |
|
AK type steel and carbon steel |
0,40 |
Stainless steel Х18Н9 (passive state) and ОХ17Н7У |
0,17 |
|
Gray cast iron |
0,36 |
Monel metal |
0,17 |
|
Stainless steels X13 and X17 (active state) |
0,32 |
Stainless steel Х18Н12М3 (passive state) |
0,20 |
|
Nickel-copper cast iron (12-15% Ni, 5-7% Si) |
0,30 |
Stainless steel Х18Н10Т |
0,25 |
|
Lead |
0,30 |
Platinum |
0,40 |
|
Tin |
0,25 |
Note . Specified numeric values The potentials and order of metals in a series can vary to varying degrees depending on the purity of the metals, the composition of sea water, the degree of aeration and the state of the surface of the metals.