What is a series of metal stresses. active metals. Alkali metals are considered the most active

The potential difference "electrode substance - solution" just serves as a quantitative characteristic of the ability of a substance (both metals andnon-metals) pass into solution in the form of ions, i.e. charactersby the OB ability of the ion and its corresponding substance.

This potential difference is calledelectrode potential.

However, direct methods for measuring such a potential differencedoes not exist, so we agreed to define them in relation tothe so-called standard hydrogen electrode, the potentialwhose value is conditionally taken as zero (often also calledreference electrode). The standard hydrogen electrode consists offrom a platinum plate immersed in an acid solution with conconcentration of ions H + 1 mol/l and washed by a jet of gaseoushydrogen under standard conditions.

The emergence of a potential at a standard hydrogen electrode can be imagined as follows. Gaseous hydrogen, being adsorbed by platinum, passes into the atomic state:

H22H.

Between atomic hydrogen formed on the surface of the plate, hydrogen ions in solution and platinum (electrons!) A state of dynamic equilibrium is realized:

H H + + e.

The overall process is expressed by the equation:

H 2 2H + + 2e.

Platinum does not take part in redox And process, but is only a carrier of atomic hydrogen.

If a plate of some metal, immersed in a solution of its salt with a concentration of metal ions equal to 1 mol / l, is connected to a standard hydrogen electrode, then a galvanic cell will be obtained. Electromotive force this element(EMF), measured at 25 ° C, and characterizes the standard electrode potential of the metal, usually denoted as E 0.

In relation to the H 2 / 2H + system, some substances will behave as oxidizing agents, others as reducing agents. At present, the standard potentials of almost all metals and many non-metals have been obtained, which characterize the relative ability of reducing or oxidizing agents to donate or capture electrons.

The potentials of the electrodes that act as reducing agents with respect to hydrogen have the “-” sign, and the “+” sign marks the potentials of the electrodes that are oxidizing agents.

If you arrange the metals in ascending order of their standard electrode potentials, then the so-called electrochemical voltage series of metals:

Li, Rb, K, Ba, Sr, Ca, N a, M g, A l, M n, Zn, C r, F e, C d, Co, N i, Sn, P b, H, Sb, V i , С u , Hg , А g , Р d , Р t , А u .

A series of stresses characterizes the chemical properties of metals.

1. The more negative the electrode potential of the metal, the greater its reducing ability.

2. Each metal is able to displace (restore) from salt solutions those metals that are in the series of metal stresses after it. The only exceptions are alkali and alkaline earth metals, which will not reduce other metal ions from solutions of their salts. This is due to the fact that in these cases, the reactions of interaction of metals with water proceed at a faster rate.

3. All metals having a negative standard electrode potential, i.e. located in the series of voltages of metals to the left of hydrogen, are able to displace it from acid solutions.

It should be noted that the presented series characterizes the behavior of metals and their salts only in aqueous solutions, since the potentials take into account the features of the interaction of one or another ion with solvent molecules. That is why the electrochemical series begins with lithium, while the more chemically active rubidium and potassium are located to the right of lithium. This is due to the exceptionally high energy of the lithium ion hydration process compared to other alkali metal ions.

The algebraic value of the standard redox potential characterizes the oxidative activity of the corresponding oxidized form. Therefore, a comparison of the values ​​of standard redox potentials allows us to answer the question: does this or that redox reaction proceed?

So, all half-reactions of oxidation of halide ions to free halogens

2 Cl - - 2 e \u003d C l 2 E 0 \u003d -1.36 V (1)

2 Br - -2e \u003d B r 2 E 0 \u003d -1.07 V (2)

2I - -2 e \u003d I 2 E 0 \u003d -0.54 V (3)

can be realized under standard conditions when lead oxide is used as an oxidizing agent ( IV ) (E 0 = 1.46 V) or potassium permanganate (E 0 = 1.52 V). When using potassium dichromate ( E0 = 1.35 V) only reactions (2) and (3) can be carried out. Finally, the use of nitric acid as an oxidizing agent ( E0 = 0.96 V) allows only a half-reaction with the participation of iodide ions (3).

Thus, the quantitative criterion for assessing the possibility of a particular redox reaction is positive value the difference between the standard redox potentials of the half-reactions of oxidation and reduction.

Restorative properties- These are the main chemical properties characteristic of all metals. They manifest themselves in interaction with a wide variety of oxidizing agents, including oxidizing agents from environment. In general, the interaction of a metal with oxidizing agents can be expressed by the scheme:

Me + Oxidizer" Me(+X),

Where (+X) is positive degree Me oxidation.

Examples of metal oxidation.

Fe + O 2 → Fe (+3) 4Fe + 3O 2 \u003d 2 Fe 2 O 3

Ti + I 2 → Ti(+4) Ti + 2I 2 = TiI 4

Zn + H + → Zn(+2) Zn + 2H + = Zn 2+ + H 2

If, from the whole series of standard electrode potentials, we single out only those electrode processes that correspond to the general equation

then we get a series of stresses of metals. In addition to metals, hydrogen is always included in this series, which makes it possible to see which metals are capable of displacing hydrogen from aqueous solutions of acids.

Table 19

A number of stresses for the most important metals are given in Table. 19. The position of a metal in a series of voltages characterizes its ability to redox interactions in aqueous solutions under standard conditions. Metal ions are oxidizing agents, and metals in the form of simple substances are reducing agents. At the same time, the further the metal is located in the series of voltages, 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 are exhibited by a simple substance - metal.

Electrode Process Potential

in a neutral medium it is B (see page 273). Active metals at the beginning of the series, having a potential much more negative than -0.41 V, displace hydrogen from water. Magnesium displaces hydrogen only from hot water. Metals located between magnesium and cadmium usually do not displace hydrogen from water. On the surface of these metals, oxide films are formed that 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, which inhibit the reaction. So, the oxide film on aluminum makes this metal resistant 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 during the interaction of lead 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 the passive state.

Metals are able to displace each other from salt solutions. The direction of the reaction is determined by their mutual position in a series of voltages. 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 row after magnesium.

The displacement of metals from their compounds by other metals was first studied in detail by Beketov. As a result of his work, he arranged the metals according to their chemical activity in a displacement series, which is the prototype of a series of metal stresses.

The mutual position of some metals in the series of voltages and in the periodic system at first glance does not correspond to each other. For example, according to the position in the periodic system, the reactivity of potassium must be greater than sodium, and sodium must be 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 system, should have approximately equal chemical activity, but in the series of voltages, zinc is located much earlier than copper. The reason for this kind of inconsistency is as follows.

When comparing metals occupying a particular position in the periodic system, the measure of their chemical activity - reducing ability - is taken as the value of the ionization energy of free atoms. Indeed, when passing, for example, from top to bottom along the main subgroup of group I periodic system the ionization energy of atoms decreases, which is associated with an increase in their radii (i.e., with a large distance of the outer electrons from the nucleus) and with an increasing screening of the positive charge of the nucleus by intermediate electron layers (see § 31). Therefore, potassium atoms exhibit greater chemical activity - they have stronger reducing properties - than sodium atoms, and sodium atoms are more active than lithium atoms.

When comparing metals in a series of voltages, the measure of chemical activity is taken as the work of converting a metal in a solid state into hydrated ions in an aqueous solution. This work can be represented as the sum of three terms: the energy of atomization - the transformation of a metal crystal into isolated atoms, the ionization energy of free metal atoms and the hydration energy of the formed ions. The atomization energy characterizes the strength of the crystal lattice of a given metal. The ionization energy of atoms - the detachment of valence electrons from them - is directly determined by the position of the metal in the periodic system. The energy released during hydration depends on the electronic structure of the ion, its charge and radius.

Lithium and potassium ions, which have 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. From this it is clear that lithium ions will hydrate with the release of more energy than potassium nones.

Thus, in the course of the transformation under consideration, energy is spent on atomization and ionization, and energy is released during hydration. The lower the total energy consumption, the easier the whole process will be and the closer to the beginning of the series of voltages the given metal will be located. But of the three terms of the total energy balance, only one - the ionization energy - is directly determined by the position of the metal in the periodic system. Consequently, there is no reason to expect that the mutual position of certain metals in a series of voltages will always correspond to their position in the periodic system. So, for lithium, the total energy consumption is less than for potassium, in accordance with which lithium is in the series of voltages before potassium.

For copper and zinc, the expenditure of energy for the ionization of free atoms and its gain during hydration of the ions are close. But metallic copper forms a stronger crystal lattice than zinc, which can be seen from a comparison of the melting points 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 voltage series.

When passing from water to non-aqueous solvents, the mutual position of metals in a series of voltages can change. The reason for this lies in the fact that the energy of solvation of ions of various metals varies in different ways when passing from one solvent to another.

In particular, the copper ion is very vigorously solvated in some organic solvents; this leads to the fact that in such solvents copper is located in a series of voltages up to hydrogen and displaces it from acid solutions.

Thus, in contrast to the periodic system of elements, a series of stresses in metals is not a reflection general patterns, on the basis of which it is possible to give a versatile Characteristic 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 ​​\u200b\u200bgiven in it refer to an aqueous solution, temperature and a unit concentration (activity) of metal ions.

Li, K, Ca, Na, Mg, Al, Zn, Cr, Fe, Pb, H 2 , Cu, Ag, Hg, Au

The further to the left the metal is in the series of standard electrode potentials, the stronger the reducing agent it is, the strongest reducing agent is metallic lithium, gold is the weakest, and, conversely, the gold (III) ion is the strongest oxidizing agent, lithium (I) is the weakest .

Each metal is able to restore from salts in solution those metals that are in a series of voltages after it, for example, iron can displace copper from solutions of its salts. However, it should be remembered that alkali and alkaline earth metals will interact directly with water.

Metals, standing in the series of voltages to the left of hydrogen, are able to displace it from solutions of dilute acids, while dissolving in them.

The reducing activity of a metal does not always correspond to its position in the periodic system, because when determining the place of a metal in a series, not only its ability to donate electrons is taken into account, but also the energy expended on the destruction of the metal crystal lattice, as well as the energy expended on the hydration of ions.

Interaction with simple substances

WITH oxygen most metals form oxides - amphoteric and basic:

4Li + O 2 \u003d 2Li 2 O,

4Al + 3O 2 \u003d 2Al 2 O 3.

Alkali metals, with the exception of lithium, form peroxides:

2Na + O 2 \u003d Na 2 O 2.

WITH halogens metals form salts of hydrohalic acids, for example,

Cu + Cl 2 \u003d CuCl 2.

WITH hydrogen the most active metals form ionic hydrides - salt-like substances in which hydrogen has an oxidation state of -1.

2Na + H 2 = 2NaH.

WITH gray metals form sulfides - salts of hydrosulfide acid:

WITH nitrogen some metals form nitrides, the reaction almost always proceeds when heated:

3Mg + N 2 \u003d Mg 3 N 2.

WITH carbon carbides are formed.

4Al + 3C \u003d Al 3 C 4.

WITH phosphorus - phosphides:

3Ca + 2P = Ca 3 P 2 .

Metals can interact with each other to form intermetallic compounds :

2Na + Sb = Na 2 Sb,

3Cu + Au = Cu 3 Au.

Metals can dissolve in each other at high temperature without interaction, forming alloys.

Alloys

Alloys are called systems consisting of two or more metals, as well as metals and non-metals that have characteristic properties inherent only in the metallic state.

The properties of alloys are very diverse and differ from the properties of their components, for example, in order to make gold harder and more suitable for making jewelry, silver is added to it, and an alloy containing 40% cadmium and 60% bismuth has a melting point of 144 °С, i.e. much lower than the melting point of its components (Cd 321 °С, Bi 271 °С).

The following types of alloys are possible:

Molten metals are mixed with each other in any ratio, dissolving in each other without limit, for example, Ag-Au, Ag-Cu, Cu-Ni and others. These alloys are homogeneous in composition, have high chemical resistance, conduct electric current;

The straightened metals are mixed with each other in any ratio, however, when cooled, they delaminate, and a mass is obtained, consisting of individual crystals of components, for example, Pb-Sn, Bi-Cd, Ag-Pb and others.

What information can be obtained from a series of voltages?

A number of metal stresses 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 some metal in the NRN. Let's discuss this issue in more detail.

The interaction of metals with acids

Metals that are in the series of voltages to the left of hydrogen react with acids - non-oxidizing agents. Metals located in the ERN to the right of H interact only with acids - oxidizing agents (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 \u003d ZnCl 2 + H 2

Zn + H 2 SO 4 \u003d 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 draw your attention to an important point: when metals interact with oxidizing acids, not hydrogen is released, but some other compounds. You can read more about this!

Interaction of metals with water

Metals located in the series of voltages to the left of Mg easily 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 \u003d 2NaOH + H 2

2K + 2H 2 O = 2KOH + H 2

Ca + 2H 2 O \u003d Ca (OH) 2 + H 2

Metals located in the range of voltages 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 the removal of the oxide film from the metal surface. Iron does not react with water at room temperature, but interacts with water vapor. Cobalt, nickel, tin, 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 circumstances.

Interaction of metals with aqueous solutions of salts

We will talk about the following types of reactions:

metal (*) + metal salt (**) = metal (**) + metal salt (*)

I would like to emphasize that the asterisks in this case do not indicate the degree of oxidation, not the valence of the metal, but simply allow us to distinguish between metal No. 1 and metal No. 2.

For such a reaction to occur, three conditions must be met simultaneously:

1. the salts involved in the process must be soluble in water (this is easy to check using the solubility table);
2. metal (*) must be in a series of voltages to the left of metal (**);
3. metal (*) should not react with water (which is also easily checked by ERN).

Example 4. Let's look at a few reactions:

Zn + CuSO 4 \u003d ZnSO 4 + Cu

K + Ni(NO 3) 2 ≠

The first reaction is easy to implement, all of the above conditions are met: copper sulfate is soluble in water, zinc is in the ERN 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 NRN). Finally, the fourth process will NOT result in nickel precipitation as potassium reacts with water; the resulting potassium hydroxide can react with a salt solution, but this is a completely different process.

The process of thermal decomposition of nitrates

Let me remind you that nitrates are salts of nitric acid. All nitrates decompose when heated, but the composition of the decomposition products may be different. The composition is determined by the position of the metal in the series of stresses.

Nitrates of metals located in the NER to the left of magnesium, when heated, form the corresponding nitrite and oxygen:

2KNO 3 \u003d 2KNO 2 + O 2

During the thermal decomposition of metal nitrates, located in a series of voltages from Mg to Cu inclusive, metal oxide, NO 2 and oxygen are formed:

2Cu(NO 3) 2 \u003d 2CuO + 4NO 2 + O 2

Finally, during the decomposition of nitrates of the least active metals (located in the NER to the right of copper), metal, nitrogen dioxide and oxygen are formed.