active metals. Active metals The electrochemical series of voltages of metals is built in order

Metals that react easily are called active metals. These include alkali, alkaline earth metals and aluminium.

Position in the periodic table

The metallic properties of the elements weaken from left to right in Mendeleev's periodic table. Therefore, elements of groups I and II are considered the most active.

Rice. 1. Active metals in the periodic table.

All metals are reducing agents and easily part with electrons at the external energy level. Active metals have only one or two valence electrons. In this case, the metallic properties are enhanced from top to bottom with an increase in the number of energy levels, because. the farther an electron is from the nucleus of an atom, the easier it is for it to separate.

Alkali metals are considered the most active:

• lithium;
• sodium;
• potassium;
• rubidium;
• cesium;
• francium.

The alkaline earth metals are:

• beryllium;
• magnesium;
• calcium;
• strontium;
• barium;
• radium.

You can find out the degree of activity of a metal by the electrochemical series of metal voltages. The more to the left of hydrogen an element is located, the more active it is. The metals to the right of hydrogen are inactive and can only interact with concentrated acids.

Rice. 2. Electrochemical series of voltages of metals.

The list of active metals in chemistry also includes aluminum, located in group III and to the left of hydrogen. However, aluminum is located on the border of active and medium active metals and does not react with certain substances under normal conditions.

Properties

Active metals are soft (can be cut with a knife), light, and have a low melting point.

The main chemical properties of metals are presented in the table.

Active metals easily react, therefore, in nature they are found only in mixtures - minerals, rocks.

Rice. 3. Minerals and pure metals.

What have we learned?

Active metals include elements of groups I and II - alkali and alkaline earth metals, as well as aluminum. Their activity is due to the structure of the atom - a few electrons are easily separated from the external energy level. These are soft light metals that quickly react with simple and complex substances, forming oxides, hydroxides, salts. Aluminum is closer to hydrogen and its reaction with substances requires additional conditions - high temperatures, destruction of the oxide film.

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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.

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

FROM 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.

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

Cu + Cl 2 \u003d CuCl 2.

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

2Na + H 2 = 2NaH.

FROM gray metals form sulfides - salts of hydrosulfide acid:

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

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

FROM carbon carbides are formed.

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

FROM 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.

Electrochemical activity series of metals(a series of voltages, a series of standard electrode potentials) - a sequence in which metals are arranged in order of increasing their standard electrochemical potentials φ 0 corresponding to the metal cation reduction half-reaction Me n+ : Me n+ + nē → Me

Practical use of the metal activity series

A number of voltages are used in practice for a comparative assessment of the chemical activity of metals in reactions with aqueous solutions of salts and acids and for the 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; the most active metals (up to and including aluminum) - and when interacting with water.
• Metals in the row 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 metals of moderate activity is accompanied by the release of hydrogen; the most active metals (up to aluminum) cannot be isolated from aqueous solutions of salts under normal conditions.

Alkali metals are considered the most active:

• lithium;
• sodium;
• potassium;
• rubidium;
• cesium;
• francium.

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