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 the environment. In general, the interaction of a metal with oxidizing agents can be expressed by the following scheme:
Me + Oxidizing agent" Me(+X),
Where (+X) is the positive oxidation state of Me.
Examples of metal oxidation.
Fe + O 2 → Fe(+3) 4Fe + 3O 2 = 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
Metal activity series
The reducing properties of metals differ from each other. Electrode potentials E are used as a quantitative characteristic of the reduction properties of metals.
The more active the metal, the more negative its standard electrode potential E o.
Metals arranged in a row as their oxidative activity decreases form an activity series.
Metal activity series
| Me | Li | K | Ca | Na | Mg | Al | Mn | Zn | Cr | Fe | Ni | Sn | Pb | H 2 | Cu | Ag | Au |
| Me z+ | Li+ | K+ | Ca2+ | Na+ | Mg 2+ | Al 3+ | Mn 2+ | Zn 2+ | Cr 3+ | Fe 2+ | Ni 2+ | Sn 2+ | Pb 2+ | H+ | Cu 2+ | 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 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 exposed to air come into contact with oxygen, so when studying the chemistry of metals, one always pays attention to the peculiarities of the interaction of the metal with oxygen.
Everyone knows that iron in humid air becomes covered with rust - hydrated iron oxide. But many metals in a compact state at not too high temperatures exhibit resistance to oxidation, since they form thin protective films on their surface. These films of oxidation products prevent the oxidizing agent from contacting the metal. The phenomenon of formation of protective layers on the surface of a metal that prevent oxidation of the metal is called passivation of the metal.
An increase in temperature promotes the oxidation of metals with oxygen. The activity of metals increases in a finely crushed state. Most metals in powder form burn in oxygen.
s-metals
Show the greatest reducing activitys-metals. Metals Na, K, Rb Cs can ignite 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 tape burns with a blinding flame.
MetalsIIA-subgroups and Li, when interacting with oxygen, form oxides.
2Ca + O2 = 2CaO
4 Li + O 2 = 2 Li 2 O
Alkali metals, exceptLi, when interacting with oxygen, they form not oxides, but peroxidesMe 2 O 2 and superoxidesMeO 2 .
2Na + O 2 = Na 2 O 2
K + O 2 = KO 2
p-metals
Metals belonging top- the block is passivated in air.
When burning in oxygen
- metals of the IIIA subgroup form oxides of the type Me 2 O 3,
- Sn is oxidized to SnO 2 , and Pb - up to PbO
- Bi goes to Bi2O3.
d-metals
Alld-period 4 metals are oxidized by oxygen. Sc, Mn, Fe are most easily oxidized. Particularly resistant to corrosion are Ti, V, Cr.
When burned in oxygen of alld
When burned in oxygen of alld-of period 4 elements, only scandium, titanium and vanadium form oxides in which Me is in the highest oxidation state, equal to group number. The remaining period 4 d-metals, when burned in oxygen, form oxides in which Me is in intermediate but stable oxidation states.
Types of oxides formed by period 4 d-metals upon combustion in oxygen:
- MeO form Zn, Cu, Ni, Co. (at T>1000°C Cu forms Cu 2 O),
- Me 2 O 3, form Cr, Fe and Sc,
- MeO 2 - Mn, and Ti,
- V forms a higher oxide - V 2 O 5 .
When burned in oxygend-metals of periods 5 and 6, as a rule, form higher oxides, the exceptions are the metals Ag, Pd, Rh, Ru.
Types of oxides formed by d-metals of periods 5 and 6 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;
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 up to hydrogen, i.e. having negative electrode potentials.
Many metals, when oxidized, transform into cations in acidic aqueous solutionsMe z + .
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 .
NO 3 - anions exhibit oxidizing properties at any concentration in solution, but the reduction products depend on the concentration of the acid and the nature of the metal being oxidized.
SO 4 2- anions exhibit oxidizing properties only in concentrated H 2 SO 4.
Reduction products of oxidizing agents: H + , NO 3 - , SO 4 2 -
2Н + + 2е - =H 2
SO 4
2-
from concentrated H 2 SO 4 SO 4
2-
+ 2e -
+ 4
H +
=
SO 2
+ 2
H 2
O
(formation of S, H 2 S is also possible)
NO 3 - from concentrated HNO 3 NO 3 - + e -
+ 2H + =
NO 2 + H 2 O
NO 3 - from dilute HNO 3 NO 3 - + 3e -
+4H+=NO+2H2O
(formation of N 2 O, N 2, NH 4 + is also possible)
Examples of reactions between metals and acids
Zn + H 2 SO 4 (diluted) " ZnSO 4 + H 2
8Al + 15H 2 SO 4 (k.) " 4Al 2 (SO 4) 3 + 3H 2 S + 12H 2 O
3Ni + 8HNO 3 (dil.) " 3Ni(NO 3) 2 + 2NO + 4H 2 O
Cu + 4HNO 3 (k.) " Cu(NO 3) 2 + 2NO 2 + 2H 2 O
Products of metal oxidation in acidic solutions
Alkali metals form a Me + type cation, s-metals of the second group form cations Me 2+.
When dissolved in acids, p-block metals form the cations indicated in the table.
The metals Pb and Bi are dissolved only in nitric acid.
| Me | Al | Ga | In | Tl | Sn | Pb | Bi |
| Mez+ | Al 3+ | Ga 3+ | 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 of 4 periods, except Cu , can be oxidized by ionsH+ in acidic solutions.
Types of cations formed by period 4 d-metals:
- 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.
The following can dissolve in HNO 3: Cd, Hg, Ag. Pd, Tc, Re dissolve in hot HNO 3.
The following dissolve in hot H 2 SO 4: 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 (a mixture of HNO 3 + HCl) 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 = H + NO + 2H 2 O
Interaction of metals with water
The oxidizing properties of water are due to H(+1).
2H 2 O + 2e -" N 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, except Be and Mg easily dissolve in water.
2 Na + 2 HOH = H 2 + 2 OH -
Na reacts vigorously with water, releasing heat. The released H2 may ignite.
2H 2 +O 2 =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 in 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 with water, Al, Ga and Sn are oxidized.
Among d-metals, they are oxidized by water when Sc and Mn, La, Y are heated. Iron reacts with water vapor.
Interaction of metals with alkali solutions
In alkaline solutions, water acts as an oxidizing agent..
2H 2 O + 2e - =H 2 + 2OH - Eo = - 0.826 B (pH = 14)
The oxidizing properties of water decrease with increasing pH, due to a decrease in H + concentration. Nevertheless, some metals that do not dissolve in water dissolve in alkali solutions, for example, Al, Zn and some others. The main reason for the dissolution of such metals in alkaline solutions is that the oxides and hydroxides of these metals exhibit amphotericity and 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 = 2Na + 3H 2
Li, K, Ca, Na, Mg, Al, Zn, Cr, Fe, Pb, H 2 , Cu, Ag, Hg, Au
The further to the left a metal is in the series of standard electrode potentials, the stronger the reducing agent it is; the strongest reducing agent is lithium metal, gold is the weakest, and, conversely, gold (III) ion is the strongest oxidizing agent, lithium (I) is the weakest .
Each metal is capable of reducing from salts in solution those metals that are in the series of stresses after it; for example, iron can displace copper from solutions of its salts. However, remember that alkali and alkaline earth metals will react directly with water.
Metals standing in the voltage series to the left of hydrogen are capable of displacing it from solutions of dilute acids and dissolving in them.
The reducing activity of a metal does not always correspond to its position in the periodic table, because when determining a metal’s place 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’s 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 = 2Li 2 O,
4Al + 3O 2 = 2Al 2 O 3.
Alkali metals, with the exception of lithium, form peroxides:
2Na + O 2 = Na 2 O 2.
WITH halogens metals form salts of hydrohalic acids, for example,
Cu + Cl 2 = CuCl 2.
WITH hydrogen the most active metals form ionic hydrides - salt-like substances in which hydrogen has an oxidation state of -1.
2Na + H2 = 2NaH.
WITH gray metals form sulfides - salts of hydrogen sulfide acid:
WITH nitrogen Some metals form nitrides; the reaction almost always occurs when heated:
3Mg + N2 = Mg3N2.
WITH carbon carbides are formed:
4Al + 3C = Al 3 C 4.
WITH phosphorus – phosphides:
3Ca + 2P = Ca 3 P 2 .
Metals can interact with each other, forming intermetallic compounds :
2Na + Sb = Na 2 Sb,
3Cu + Au = Cu 3 Au.
Metals can dissolve into each other at high temperatures without reacting, 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 for gold to become 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 °C, i.e. much lower than the melting point of its components (Cd 321 °C, Bi 271 °C).
The following types of alloys are possible:
Molten metals are mixed with each other in any ratio, dissolving in each other indefinitely, for example, Ag-Au, Ag-Cu, Cu-Ni and others. These alloys are homogeneous in composition, have high chemical resistance, and conduct electric current;
The straightened metals are mixed with each other in any ratio, but when cooled they separate, 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 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 would like to draw your attention to an important point: when metals interact with oxidizing acids, it is not hydrogen that is released, but some other compounds. You can read more about this!
Interaction of metals with water
Metals located in the voltage series to the left of Mg readily react with water already at room temperature, 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.
Electrochemical activity series of metals(a series of voltages, a series of standard electrode potentials) - a sequence in which metals are arranged in increasing order of their standard electrochemical potentials φ 0, corresponding to the half-reaction of reduction of the metal cation Me n+: Me n+ + nē → Me
Practical use of the activity series of metals
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 series to the right of hydrogen do not interact with aqueous solutions of non-oxidizing acids under normal conditions.
- During electrolysis, metals to the right of hydrogen are released at the cathode; the reduction of moderately active metals is accompanied by the release of hydrogen; The most active metals (up to aluminum) cannot be isolated from aqueous salt solutions under normal conditions.
Alkali metals are considered the most active:
- lithium;
- sodium;
- potassium;
- rubidium;
- cesium;
- French
Metals that react easily are called active metals. These include alkali, alkaline earth metals and aluminum.
Position in the periodic table
The metallic properties of elements decrease from left to right in the 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 outer energy level. Active metals have only one or two valence electrons. In this case, metallic properties increase from top to bottom with increasing number of energy levels, because The further 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;
- French
Alkaline earth metals include:
- beryllium;
- magnesium;
- calcium;
- strontium;
- barium;
- radium.
The degree of activity of a metal can be determined by the electrochemical series of metal voltages. The further to the left of hydrogen an element is located, the more active it is. Metals to the right of hydrogen are inactive and can only react 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 on the border of active and intermediately active metals and does not react with some 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.
|
Reaction |
The equation |
Exception |
|
Alkali metals spontaneously ignite in air when 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 is slow. Aluminum reacts with water only after removing the oxide film |
|
React with acids to form salts |
Ca + 2HCl → CaCl 2 + H 2; 2K + 2HMnO 4 → 2KMnO 4 + H 2 |
|
|
Interact 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, so 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 determined by 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, and 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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