Li, K, Ca, Na, Mg, Al, Zn, Cr, Fe, Pb, H 2 , Cu, Ag, Hg, Au
The more to the left the metal is in the series of standard electrode potentials, the stronger the reducing agent 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, which 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.
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.
Main Chemical properties metals are presented in the table.
|
Reaction |
The equation |
Exception |
|
Alkali metals ignite spontaneously in air, interacting with oxygen |
K + O 2 → KO 2 |
Lithium reacts with oxygen only at high temperatures. |
|
Alkaline earth metals and aluminum form oxide films in air, and spontaneously ignite when heated. |
2Ca + O 2 → 2CaO |
|
|
React with simple substances to form salts |
Ca + Br 2 → CaBr 2; |
Aluminum does not react with hydrogen |
|
React violently with water, forming alkalis and hydrogen |
|
The reaction with lithium proceeds slowly. Aluminum reacts with water only after the removal of the oxide film. |
|
React with acids to form salts |
Ca + 2HCl → CaCl 2 + H 2; 2K + 2HMnO 4 → 2KMnO 4 + H 2 |
|
|
React with salt solutions, first reacting with water and then with salt |
2Na + CuCl 2 + 2H 2 O: 2Na + 2H 2 O → 2NaOH + H 2; |
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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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 particular metal in a series of stresses 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 only displaces hydrogen 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 a 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 in this case by their mutual position in the 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, during the transition, 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 large distance of external electrons from the nucleus) and with 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 electronic structure ion, its charge and radius.
Lithium and potassium ions, having the same charge but different radii, will create unequal electric fields around them. 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 solvation energy of different metal ions changes differently 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 of the general Regularity, on the basis of which it is possible to give a versatile Characteristic of the chemical properties of 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.
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. AT general view the interaction of a metal with oxidizing agents can be expressed by the scheme:
Me + Oxidizer" Me(+X),
Where (+X) is the positive oxidation state of Me.
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
Activity series of metals
The reducing properties of metals differ from each other. Electrode potentials E are used as a quantitative characteristic of the reducing properties of metals.
The more active the metal, the more negative its standard electrode potential E o.
Metals arranged in a row as their oxidizing activity decreases form a row of activity.
Activity series of metals
| Me | Li | K | Ca | Na | mg | Al | Mn | Zn | Cr | Fe | Ni | sn | Pb | H2 | Cu | Ag | Au |
| Mez+ | Li+ | K+ | Ca2+ | Na+ | Mg2+ | Al 3+ | Mn2+ | Zn2+ | Cr3+ | Fe2+ | Ni2+ | sn 2+ | Pb 2+ | H+ | Cu2+ | Ag+ | Au 3+ |
| E o ,B | -3,0 | -2,9 | -2,87 | -2,71 | -2,36 | -1,66 | -1,18 | -0,76 | -0,74 | -0,44 | -0,25 | -0,14 | -0,13 | 0 | +0,34 | +0,80 | +1,50 |
The reduction of a metal from a solution of its salt with another metal with a higher reducing activity is called cementation.. Cementation is used in metallurgical technologies.
In particular, Cd is obtained by reducing it from a solution of its salt with zinc.
Zn + Cd 2+ = Cd + Zn 2+
3.3. 1. Interaction of metals with oxygen
Oxygen is a strong oxidizing agent. It can oxidize the vast majority of metals exceptAuandPt . Metals in air come into contact with oxygen, therefore, when studying the chemistry of metals, attention is always paid to the features of the interaction of a metal with oxygen.
We all know that iron humid air covered with rust - hydrated iron oxide. But many metals in a compact state at a not too high temperature show resistance to oxidation, since they form thin protective films on their surface. These films of oxidation products do not allow the oxidizing agent to come into contact with the metal. The phenomenon of the formation of protective layers on the surface of the metal that prevent the oxidation of the metal is called metal passivation.
An increase in temperature promotes the oxidation of metals by oxygen. The activity of metals increases in the finely divided state. Most metals in powder form burn in oxygen.
s-metals
The greatest restorative activity is showns-metals. The metals Na, K, Rb Cs are capable of igniting in air, and they are stored in sealed vessels or under a layer of kerosene. Be and Mg are passivated at low temperatures in air. But when ignited, the Mg strip burns with a dazzling flame.
MetalsIIA-subgroups and Li, when interacting with oxygen, form oxides.
2Ca + O 2 \u003d 2CaO
4 Li + O 2 \u003d 2 Li 2 O
Alkali metals, other thanLi, when interacting with oxygen, they form not oxides, but peroxidesMe 2 O 2 and superoxidesMeO 2 .
2Na + O 2 \u003d Na 2 O 2
K + O 2 = KO 2
p-metals
Metals ownedp- to the block on air are passivated.
When burning in oxygen
- IIIA-subgroup metals form oxides of the type Me 2 O 3,
- Sn is oxidized to SNO 2 , and Pb - up to PbO
- Bi goes to Bi 2 O 3.
d-metals
Alld- period 4 metals are oxidized by oxygen. Sc, Mn, Fe are most easily oxidized. Particularly resistant to Ti, V, Cr corrosion.
When burned in oxygen of alld
When burned in oxygen of alld- elements of the 4th period, only scandium, titanium and vanadium form oxides in which Me is in the highest oxidation state, equal to group number. The remaining d-metals of the 4th period, when burned in oxygen, form oxides in which Me is in intermediate but stable oxidation states.
Types of oxides formed by d-metals of 4 periods during combustion in oxygen:
- Meo form Zn, Cu, Ni, Co. (at T>1000оС Cu forms Cu 2 O),
- Me 2 O 3, form Cr, Fe and Sc,
- MeO 2 - Mn and Ti
- V forms the highest oxide - V 2 O 5 .
When burned in oxygend-metals of 5 and 6 periods, as a rule, form higher oxides, the exceptions are the metals Ag, Pd, Rh, Ru.
Types of oxides formed by d-metals of 5 and 6 periods during combustion in oxygen:
- Me 2 O 3- form Y, La; Rh;
- MeO 2- Zr, Hf; Ir:
- Me 2 O 5- Nb, Ta;
- MeO 3- Mo, W
- Me 2 O 7- Tc, Re
- Meo 4 - Os
- MeO- Cd, Hg, Pd;
- Me 2 O- Ag;
The interaction of metals with acids
In acid solutions, the hydrogen cation is an oxidizing agent.. The H + cation can oxidize metals in the activity series to hydrogen, i.e. having negative electrode potentials.
Many metals, when oxidized, in acidic aqueous solutions, many turn into cationsMez + .
Anions of a number of acids are capable of exhibiting oxidizing properties that are stronger than H + . Such oxidizing agents include anions and the most common acids H 2 SO 4 andHNO 3 .
Anions NO 3 - exhibit oxidizing properties at any concentration in solution, but the reduction products depend on the concentration of the acid and the nature of the oxidized metal.
Anions SO 4 2- exhibit oxidizing properties only in concentrated H 2 SO 4 .
Oxidizer reduction products: H + , NO 3 - , SO 4 2 -
2H + + 2e - =H 2
SO 4
2-
from concentrated H 2 SO 4 SO 4
2-
+ 2e -
+ 4
H +
=
SO 2
+ 2
H 2
O
(possible also the formation of S, H 2 S)
NO 3 - from concentrated HNO 3 NO 3 - + e -
+2H+=
NO 2 + H 2 O
NO 3 - from diluted HNO 3 NO 3 - + 3e -
+4H+=NO + 2H 2 O
(It is also possible to form N 2 O, N 2, NH 4 +)
Examples of reactions of interaction of metals with acids
Zn + H 2 SO 4 (razb.) "ZnSO 4 + H 2
8Al + 15H 2 SO 4 (c.) "4Al 2 (SO 4) 3 + 3H 2 S + 12H 2 O
3Ni + 8HNO 3 (deb.) " 3Ni(NO 3) 2 + 2NO + 4H 2 O
Cu + 4HNO 3 (c.) "Cu (NO 3) 2 + 2NO 2 + 2H 2 O
Metal oxidation products in acidic solutions
Alkali metals form a cation of the Me + type, s-metals of the second group form cations Me 2+.
The p-block metals, when dissolved in acids, form the cations indicated in the table.
Metals Pb and Bi dissolve only in nitric acid.
| Me | Al | Ga | In | Tl | sn | Pb | Bi |
| Mez+ | Al 3+ | Ga3+ | In 3+ | Tl+ | sn 2+ | Pb 2+ | Bi 3+ |
| Eo,B | -1,68 | -0,55 | -0,34 | -0,34 | -0,14 | -0,13 | +0,317 |
All d-metals 4 periods except Cu , can be oxidized by ionsH+ in acid solutions.
Types of cations formed by d-metals 4 periods:
- Me 2+(form d-metals ranging from Mn to Cu)
- Me 3+ ( form Sc, Ti, V, Cr and Fe in nitric acid).
- Ti and V also form cations MeO 2+
In acidic solutions, H + can oxidize: Y, La, Cd.
In HNO 3 can dissolve: Cd, Hg, Ag. Hot HNO 3 dissolves Pd, Tc, Re.
In hot H 2 SO 4 dissolve: Ti, Zr, V, Nb, Tc, Re, Rh, Ag, Hg.
Metals: Ti, Zr, Hf, Nb, Ta, Mo, W are usually dissolved in a mixture of HNO 3 + HF.
In aqua regia (HNO 3 + HCl mixtures) Zr, Hf, Mo, Tc, Rh, Ir, Pt, Au and Os can be dissolved with difficulty). The reason for the dissolution of metals in aqua regia or in a mixture of HNO 3 + HF is the formation of complex compounds.
Example. The dissolution of gold in aqua regia becomes possible due to the formation of a complex -
Au + HNO 3 + 4HCl \u003d H + NO + 2H 2 O
Interaction of metals with water
The oxidizing properties of water are due H(+1).
2H 2 O + 2e -" H 2 + 2OH -
Since the concentration of H + in water is low, its oxidizing properties are low. Metals can dissolve in water E< - 0,413 B. Число металлов, удовлетворяющих этому условию, значительно больше, чем число металлов, реально растворяющихся в воде. Причиной этого является образование на поверхности большинства металлов плотного слоя оксида, нерастворимого в воде. Если оксиды и гидроксиды металла растворимы в воде, то этого препятствия нет, поэтому щелочные и щелочноземельные металлы энергично растворяются в воде. Alls- metals, other than Be and Mg easily soluble in water.
2 Na + 2 HOH = H 2 + 2 Oh -
Na reacts vigorously with water, releasing heat. Emitted H 2 may ignite.
2H 2 + O 2 \u003d 2H 2 O
Mg dissolves only in boiling water, Be is protected from oxidation by an inert insoluble oxide
p-block metals are less powerful reducing agents thans.
Among p-metals, the reducing activity is higher for metals of the IIIA subgroup, Sn and Pb are weak reducing agents, Bi has Eo > 0.
p-metals do not dissolve in water under normal conditions. When the protective oxide is dissolved from the surface in alkaline solutions, Al, Ga, and Sn are oxidized by water.
Among d-metals, they are oxidized by water when heated Sc and Mn, La, Y. Iron reacts with water vapor.
Interaction of metals with alkali solutions
In alkaline solutions, water acts as an oxidizing agent..
2H 2 O + 2e - \u003dH 2 + 2OH - Eo \u003d - 0.826 B (pH \u003d 14)
The oxidizing properties of water decrease with increasing pH, due to a decrease in the concentration of H +. Nonetheless, some metals that do not dissolve in water dissolve in alkali solutions, for example, Al, Zn and some others. main reason the dissolution of such metals in alkaline solutions is that the oxides and hydroxides of these metals are amphoteric, dissolve in alkali, eliminating the barrier between the oxidizing agent and the reducing agent.
Example. Dissolution of Al in NaOH solution.
2Al + 3H 2 O + 2NaOH + 3H 2 O \u003d 2Na + 3H 2
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 peculiarities 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 hydration process of lithium ions in comparison with ions of other alkali metals.
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 take place?
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.