Metals include s-elements of groups 1 and 2, all d- and f-elements, as well as a number of p-elements of the main subgroups: 3 (except boron), 4 (Ge, Sn, Pb), 5 (Sb, Bi) and Ro. The most typical metal elements are located at the beginning of periods. Earlier we talked about the fact that highly delocalized bonds occur in metals. This is due to the fact that, due to the screening effect, the valence electrons in metal atoms are weaker attracted to the nucleus and the first ionization energies for them are relatively low. At our usual temperature (about 300 K), which is quite far from absolute zero, the energy of thermal motion is sufficient for the free movement of electrons throughout the metal.
Since the bond in metals is highly delocalized and extends over the entire crystal, metals have high plasticity, electrical and thermal conductivity. Silver and copper have the highest electrical and thermal conductivity, and mercury has the lowest. The latter is also the most fusible metal (-38.9 C). The most refractory metal is tungsten (3390 C). Such a large difference in the melting and boiling points is explained by the presence in metals, in addition to metallic bonds, of a certain proportion of covalent bonds, especially for transition elements with a large number of valence electrons.
Let's consider the electronic configurations of mercury and tungsten.
Hg – 5d 10 6s 2; W – 5d 4 6s 2 . The intermolecular interaction between mercury atoms is very small, so small that in general, at high density, due to the gravity of the atoms, it is the most fusible metal. Since all the sublevels in the mercury atom are filled, the formation of covalent bonds is generally impossible, and the metallic bond is quite weak, weaker than in alkali metals, which are generally the most fusible among all metals. On the contrary, the formation of four valence bonds at once is possible in the W atom. In addition, metallic bonding is the strongest of all 5d elements, and the atoms themselves are heavier than their electronic counterparts: Mo and Cr. The combination of these factors leads to the greatest refractoriness of tungsten.
The electronic configuration of osmium (5d 6 6s 2) is such that it lacks 4 electrons before completing the 5d sublevel, so it is most strongly capable of attracting electrons from neighboring atoms, which causes a shortening of the metal-metal bond. Therefore, osmium has the highest density (22.4 g/cm3).
IN pure form metals are relatively rare. Basically, these are chemically inert metals (gold, as well as platinum group metals - platinum, rhodium, iridium, osmium, etc.). Silver, copper, mercury, and tin can be found both in the native state and in the form of compounds. The remaining metals occur in the form of compounds called ores.
Metals are obtained from their compounds by reducing them from oxides. C, CO, active metals, hydrogen, and methane are used as reducing agents. If the ore is metal sulfide (ZnS, FeS 2), then it is first converted into oxide. The reduction of metals from their compounds by other metals is called metallothermy. Some metals are extracted from solutions of their salts by electrolysis, for example aluminum or sodium. All methods for obtaining metals from their compounds are based on redox processes.
The process of electron transfer in a redox half-reaction can be represented by the following general equation:
The process of electron transition corresponds to a change in the Gibbs energy equal to ∆G = –nFE, where F (Faraday’s constant, corresponds to the amount of electricity required to reduce or oxidize one mole of a substance) = 96500 C/mol, n is the number of electrons, E is the electrode potential, B is the voltage difference between the oxidizing agent and the reducing agent. On the other hand, ∆G = –RTlnK = –nFE; RTlnK = nFE. Hence E = RTlnK/nF. Since K = /, and 2.3lnK = logK, the dependence of the electrode potential on the concentrations of substances - participants in the electrode process - and on temperature is expressed by the following equation:
E = E 0 + log/ – Nernst equation.
At standard temperature (298 K) the equation takes the form:
E = E 0 + 0.059lg/
The concentration of the oxidizing agent is always given in the numerator, and the potential is always given for the reduction half-reaction: Ox + ne = Red.
At equilibrium concentrations of the oxidizing agent and the reducing agent equal to unity, E = E 0 is the standard electrode potential: this is the potential of a given electrode process at unit concentrations of all substances. Since the absolute value of standard electrode potentials cannot be determined, then the half-reaction potential is taken as the starting point: 2Н + + 2е = Н 2 . The potential of this electrode process is assumed to be 0 at unit concentrations of the hydrogen cation. The hydrogen electrode consists of a platinum plate, which is immersed in a solution of sulfuric acid with [H + ] = 1 mol/l and washed by a stream of H 2 under a pressure of 101325 Pa at 298 K.
The electrode potential is the EMF of a galvanic cell, which consists of the electrode under study and a standard hydrogen electrode. By arranging metals in increasing order of the magnitude of their electrode potentials, we obtain a number of standard electrode potentials of metals. It characterizes Chemical properties metals Each metal in the series displaces all subsequent metals from the solution of their salts. Metals in the row to the left of hydrogen displace it from acid solutions.
The potential of any redox reaction can be calculated based on the values of the half-reaction potentials.
Let's consider a simple example: Zn + 2HCl = ZnCl 2 + H 2. For this process, two half-reactions take place:
Zn 2+ + 2e = Zn 0 E 0 (Zn 2+ /Zn) = –0.76 B
2H + + 2e = H 2 0 E 0 (2H + /H 2) = 0.00 B
Since the potential of the second half-reaction is higher than the first, the second half-reaction will proceed from left to right, that is, towards the formation of hydrogen molecules. The first half-reaction will proceed from right to left, that is, towards the formation of zinc cations.
When considering the production of metals, we talked about the fact that a number of metals are reduced from their oxides by other, more active metals. For example, magnesium can reduce copper from copper(II) oxide. Let's compare two half-reactions:
Cu 2+ + 2e = Cu E 0 = +0.34 V
Mg 2+ + 2e = Mg E 0 = –2.36 V
The potential of the first half-reaction is higher than the second and it is the one that will proceed from left to right, and the second - from right to left.
Thus, to determine the direction of redox reactions, it is necessary to write down two half-reactions from the oxidized form to the reduced form and compare their potentials. A reaction with a higher potential will proceed from left to right, and one with a lower potential will proceed from right to left.
Almost all reactions of metals are redox processes and to determine their direction it is necessary, first of all, to take into account the potentials of each of the half-reactions in the redox process. But, besides, there are exceptions. For example, lead is insoluble in sulfuric acid, despite the fact that the potential of the Pb 2+ /Pb pair is –0.15 V. The fact is that lead sulfate is insoluble and its formation prevents further oxidation of lead.
Lecture 15.
Electrolysis.
Solutions and melts of electrolytes contain oppositely charged ions (cations and anions), which are in constant motion. If inert (graphite) electrodes are immersed in this kind of liquid, for example, in a melt of sodium chloride (melts at 801 0 C) and a constant electric current is passed, then the ions under the influence of an external electric field will move towards the electrodes, cations - towards the cathode, and anions - towards anode. Sodium cations, having reached the cathode, accept electrons from it and are reduced to metallic sodium:
Chloride ions are oxidized at the anode:
2Сl – – 2e = Cl 2 0
As a result, metallic sodium is released at the cathode, and molecular chlorine at the anode. The overall equation for the electrolysis of molten sodium chloride is as follows.
K: Na + + e = Na 0 2
A: 2Сl – – 2e = Cl 2 0 1
2Na + + 2Сl – electrolysis ® 2Na 0 + Cl 2 0
2NaСl = 2Na + Cl 2
This reaction is redox: an oxidation process occurs at the anode, and a reduction process occurs at the cathode.
The redox process occurring on the electrodes during the passage electric current through a melt or electrolyte solution is called electrolysis.
The essence of electrolysis is the implementation of chemical reactions using electrical energy. In this case, the cathode gives electrons to cations, and the anode accepts electrons from anions. The action of direct electric current is much stronger than the action of chemical reducing agents and oxidizing agents. It was through electrolysis that fluorine gas was first obtained.
Electrolysis was carried out in a solution of potassium fluoride in hydrofluoric acid. IN in this case Fluorine is released at the anode, and hydrogen is released at the cathode. Electrolysis is carried out in an electrolytic bath.
It is necessary to distinguish between the electrolysis of molten electrolytes and their solutions. In the latter case, water molecules can participate in the processes. For example, during the electrolysis of an aqueous solution of sodium chloride on inert (graphite) electrodes, water molecules are reduced at the cathode instead of sodium cations.
2H 2 O + 2e = H 2 + 2OH –
and chloride ions are oxidized at the anode:
2Сl – – 2e = Cl 2 0
As a result, hydrogen is released at the cathode, chlorine is released at the anode, and sodium hydroxide molecules accumulate in the solution. General equation electrolysis of an aqueous solution of sodium chloride has the form:
K: 2H 2 O + 2e = H 2 + 2OH –
A: 2Сl – – 2e = Cl 2 0
2H 2 O + 2Сl – = H 2 + Cl 2 + 2OH –
By the way, this is how industry produces hydroxides of all alkali and some alkaline earth metals, as well as aluminum.
What is the difference between the electrolysis of melts and aqueous solutions of electrolytes? Recovery processes at the cathode of aqueous solutions of electrolytes depend on the value of the standard electrode potentials of the metals, namely, they most often act as cations that are reduced at the cathode. There are three possible options here:
1. Metal cations that have a standard electrode potential are higher than that of hydrogen, that is, more than zero during electrolysis are completely reduced at the cathode (copper, silver, gold and others).
2. Metal cations having very small value standard electrode potential (from lithium to aluminum inclusive), are not reduced at the cathode, but water molecules are reduced.
3. Metal cations, whose standard electrode potential is less than that of hydrogen, but greater than that of aluminum, are reduced during electrolysis at the cathode along with water molecules.
If several metal cations are simultaneously present in an aqueous solution, then during electrolysis their release at the cathode proceeds in the order of decreasing algebraic value of the standard electrode potential of the corresponding metal. For example, when analyzing bronze type BrAZh or BrAZhMts (copper, aluminum, iron and manganese), you can select a certain current value, separate the copper onto an inert electrode (for example, platinum), pull out the electrode, weigh it and determine the copper content. Then separate the aluminum and determine its content. This is a good way to separate metals from positive value standard electrode potential.
All electrodes are divided into insoluble (inert) - carbon, graphite, platinum, iridium. Soluble - copper, silver, zinc, cadmium, nickel and others. The concept of a soluble electrode is important for the anode, since it is the one that is capable of dissolving during electrolysis. At the insoluble anode, during the electrolysis process, oxidation of anions or water molecules occurs. In this case, the anions of oxygen-free acids are quite easily oxidized. If anions of oxygen-containing acids are present in the solution, then water molecules are oxidized at the anode, releasing oxygen according to the reaction:
2H 2 O – 4e = O 2 + 4H +
During electrolysis, the soluble anode itself oxidizes, giving up electrons to the external electrical circuit and going into solution:
A: Me Û Me n+ + ne –
Let's look at examples of electrolysis of melts and electrolyte solutions.
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. On the surface of these metals, oxide films are formed that have 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 the main subgroup of group I of the periodic system, 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. From this it is clear 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 ions various metals 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, in contrast to the periodic system of elements, a series of metal stresses is not a reflection of a general pattern, on the basis of which it is possible to give a comprehensive Characteristic of the chemical properties of metals. A series of voltages characterizes only the redox ability Electrochemical system“metal - metal ion” under strictly defined conditions: the values given in it refer to an 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 more high ability accept electrons, so it will be cathodic or noble. On the other hand, there 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-Mg 2+ |
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 . The indicated numerical values of potentials and the order of metals in a series can vary to varying degrees depending on the purity of the metals, composition sea water, degree of aeration and surface condition of metals.
A number of standard electrode potentials quantitatively characterize the reducing ability of metal atoms and the oxidizing ability of their ions.
A number of standard electrode potentials make it possible to resolve the issue of the direction of spontaneous occurrence of redox reactions. As in the general case, any chemical reaction, the determining factor here is the sign of the change in the isobaric potential of the reaction. But this means that the first of these systems will act as a reducing agent, and the second as an oxidizing agent. With direct interaction of substances, the possible direction of the reaction will, of course, be the same as when it is carried out in a galvanic cell.
A number of standard electrode potentials make it possible to resolve the issue of the direction of spontaneous occurrence of redox reactions. As in the general case of any chemical reaction, the determining factor here is the sign of the change in the Gibbs energy of the reaction. But this means that the first of these systems will act as a reducing agent, and the second as an oxidizing agent. With direct interaction of substances, the possible direction of the reaction will, of course, be the same as when it is carried out in a galvanic cell.
A number of standard electrode potentials characterize the chemical properties of metals.
| Standard hydrogen electrode.| Galvanic circuit for measuring the standard electrode potential of a metal. |
A number of standard electrode potentials characterize the chemical properties of metals. It is used when considering the sequence of ion discharge during electrolysis, as well as when describing the general properties of metals.
A number of standard electrode potentials make it possible to resolve the issue of the direction of spontaneous occurrence of oxidizing and non-reducing reactions. As in the general case of any chemical reaction, the determining factor here is the change in the isobaric potential of the reaction. But this means that the first of these systems will act as a reducing agent, and the second - as an oxidizing agent. With direct interaction of substances, the possible direction of the reaction will, of course, be the same as when carried out in a galvanic cell.
A number of standard electrode potentials characterize the chemical properties of metals. It is used to determine the discharge sequence of ions during electrolysis, as well as to describe the general properties of metals. In this case, the values of standard electrode potentials quantitatively characterize the reducing ability of metals and the oxidizing ability of their ions.
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 NER 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), but 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 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, releasing hydrogen and forming 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 of 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.
