active metals. Metals Metals in a chemical reaction
The structure of metal atoms determines not only the characteristic physical properties of simple substances - metals, but also their general chemical properties.
With a large variety, all chemical reactions of metals are redox and can only be of two types: compounds and substitutions. Metals are able to donate electrons during chemical reactions, that is, to be reducing agents, to show only a positive oxidation state in the compounds formed.
In general, this can be expressed by the scheme:
Me 0 - ne → Me + n,
where Me - metal - a simple substance, and Me 0 + n - metal chemical element in the compound.
Metals are able to donate their valence electrons to non-metal atoms, hydrogen ions, other metal ions, and therefore will react with non-metals - simple substances, water, acids, salts. However, the reducing ability of metals is different. The composition of the reaction products of metals with various substances also depends on the oxidizing ability of the substances and the conditions under which the reaction proceeds.
At high temperatures, most metals burn in oxygen:
2Mg + O 2 \u003d 2MgO
Only gold, silver, platinum and some other metals do not oxidize under these conditions.
Many metals react with halogens without heating. For example, aluminum powder, when mixed with bromine, ignites:
2Al + 3Br 2 = 2AlBr 3
When metals interact with water, hydroxides are sometimes formed. Alkali metals, as well as calcium, strontium, barium, interact very actively with water under normal conditions. The general scheme of this reaction looks like this:
Me + HOH → Me(OH) n + H 2
Other metals react with water when heated: magnesium when it boils, iron in water vapor when it boils red. In these cases, metal oxides are obtained.
If the metal reacts with an acid, then it is part of the resulting salt. When a metal interacts with acid solutions, it can be oxidized by the hydrogen ions present in that solution. The abbreviated ionic equation in general form can be written as follows:
Me + nH + → Me n + + H 2
Anions of such oxygen-containing acids, such as concentrated sulfuric and nitric acids, have stronger oxidizing properties than hydrogen ions. Therefore, those metals that are not able to be oxidized by hydrogen ions, such as copper and silver, react with these acids.
When metals interact with salts, a substitution reaction occurs: electrons from the atoms of the substituting - more active metal pass to the ions of the substituting - less active metal. Then the network replaces metal with metal in salts. These reactions are not reversible: if metal A displaces metal B from a salt solution, then metal B will not displace metal A from a salt solution.
In descending order of chemical activity, manifested in the reactions of displacement of metals from each other from aqueous solutions of their salts, metals are located in the electrochemical series of voltages (activity) of metals:
Li → Rb → K → Ba → Sr → Ca → Na→ Mg → Al → Mn → Zn → Cr → → Fe → Cd→ Co → Ni → Sn → Pb → H → Sb → Bi → Cu → Hg → Ag → Pd → Pt → Au
The metals located to the left of this row are more active and are able to displace the metals following them from salt solutions.
Hydrogen is included in the electrochemical series of voltages of metals, as the only non-metal that shares a common property with metals - to form positively charged ions. Therefore, hydrogen replaces some metals in their salts and can itself be replaced by many metals in acids, for example:
Zn + 2 HCl \u003d ZnCl 2 + H 2 + Q
Metals standing in the electrochemical series of voltages up to hydrogen displace it from solutions of many acids (hydrochloric, sulfuric, etc.), and all following it, for example, do not displace copper.
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Metals are active reducing agents with a positive oxidation state. Due to their chemical properties, metals are widely used in industry, metallurgy, medicine, and construction.
Metal activity
In reactions, metal atoms donate valence electrons and are oxidized. The more energy levels and fewer electrons a metal atom has, the easier it is for it to donate electrons and enter into reactions. Therefore, metallic properties increase from top to bottom and from right to left in the periodic table.
Rice. 1. Change in metallic properties in the periodic table.
The activity of simple substances is shown in the electrochemical series of metal voltages. To the left of hydrogen are active metals (activity increases towards the left edge), to the right - inactive.
The most active are alkali metals, which are in group I of the periodic table and are to the left of hydrogen in the electrochemical series of voltages. They react with many substances already at room temperature. They are followed by alkaline earth metals, which are included in group II. They react with most substances when heated. Metals that are in the electrochemical series from aluminum to hydrogen (medium activity) require additional conditions for entering into reactions.
Rice. 2. Electrochemical series of voltages of metals.
Some metals exhibit amphoteric properties or duality. Metals, their oxides and hydroxides react with acids and bases. Most metals only react with certain acids to replace the hydrogen and form a salt. The most pronounced dual properties show:
- aluminum;
- lead;
- zinc;
- iron;
- copper;
- beryllium;
- chromium.
Each metal is capable of displacing another metal to the right of it in the electrochemical series from salts. Metals to the left of hydrogen displace it from dilute acids.
Properties
Features of the interaction of metals with different substances are presented in the table of chemical properties of metals.
|
Reaction |
Peculiarities |
The equation |
|
With oxygen |
Most metals form oxide films. Alkali metals ignite spontaneously in the presence of oxygen. In this case, sodium forms peroxide (Na 2 O 2), the remaining metals of group I are superoxides (RO 2). When heated, alkaline earth metals spontaneously ignite, while metals of medium activity oxidize. Gold and platinum do not interact with oxygen |
4Li + O 2 → 2Li 2 O; 2Na + O 2 → Na 2 O 2; K + O 2 → KO 2; 4Al + 3O 2 → 2Al 2 O 3; 2Cu + O 2 → 2CuO |
|
With hydrogen |
Alkaline reacts at room temperature, while alkaline earth reacts when heated. Beryllium does not react. Magnesium additionally needs high pressure |
Sr + H 2 → SrH 2 ; 2Na + H 2 → 2NaH; Mg + H 2 → MgH 2 |
|
Only active metals. Lithium reacts at room temperature. Other metals - when heated |
6Li + N 2 → 2Li 3 N; 3Ca + N 2 → Ca 3 N 2 |
|
|
With carbon |
Lithium and sodium, the rest - when heated |
4Al + 3C → Al 3 C4; 2Li+2C → Li 2 C 2 |
|
Gold and platinum do not interact |
2K + S → K 2 S; Fe + S → FeS; Zn + S → ZnS |
|
|
with phosphorus |
When heated |
3Ca + 2P → Ca 3 P 2 |
|
With halogens |
Only inactive metals do not react, copper - when heated |
Cu + Cl 2 → CuCl 2 |
|
Alkali and some alkaline earth metals. When heated, in an acidic or alkaline environment, metals of medium activity react |
2Na + 2H 2 O → 2NaOH + H 2; Ca + 2H 2 O → Ca (OH) 2 + H 2; Pb + H 2 O → PbO + H 2 |
|
|
With acids |
Metals to the left of hydrogen. Copper dissolves in concentrated acids |
Zn + 2HCl → ZnCl 2 + 2H 2; Fe + H 2 SO 4 → FeSO 4 + H 2; Cu + 2H 2 SO 4 → CuSO 4 + SO 2 + 2H 2 O |
|
With alkalis |
Only amphoteric metals |
2Al + 2KOH + 6H 2 O → 2K + 3H 2 |
|
Active substitutes for less active metals |
3Na + AlCl 3 → 3NaCl + Al |
Metals interact with each other and form intermetallic compounds - 3Cu + Au → Cu 3 Au, 2Na + Sb → Na 2 Sb.
Application
The general chemical properties of metals are used to create alloys, detergents, and are used in catalytic reactions. Metals are present in batteries, electronics, and load-bearing structures.
The main fields of application are indicated in the table.
Rice. 3. Bismuth.
What have we learned?
From the 9th grade chemistry lesson, we learned about the basic chemical properties of metals. The ability to interact with simple and complex substances determines the activity of metals. The more active the metal, the easier it reacts under normal conditions. Active metals react with halogens, non-metals, water, acids, salts. Amphoteric metals interact with alkalis. Inactive metals do not react with water, halogens, and most non-metals. Briefly reviewed the application areas. Metals are used in medicine, industry, metallurgy, and electronics.
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This lesson is devoted to the study of the topic “General properties of metals. Metal connection. During the lesson, the general chemical properties of metals, the features of the metallic chemical bond will be considered. The teacher will explain the similarities between the chemical and physical properties of metals using a model of their internal structure.
Topic: Chemistry of metals
Lesson: General properties of metals. metal connection
Metals are characterized by common physical properties: they have a special metallic luster, high thermal and electrical conductivity, and ductility.
Metals also share some common chemical properties. It is important to remember that in chemical reactions, metals act as reducing agents: they donate electrons and increase their oxidation state. Consider some reactions in which metals participate.
INTERACTION WITH OXYGEN
Many metals can react with oxygen. Usually the products of these reactions are oxides, but there are exceptions, which you will learn about in the next lesson. Consider the interaction of magnesium with oxygen.
Magnesium burns in oxygen to form magnesium oxide:
2Mg + O 2 \u003d 2MgO
Rice. 1. Combustion of magnesium in oxygen
Magnesium atoms donate their outer electrons to oxygen atoms: two magnesium atoms donate two electrons each to two oxygen atoms. In this case, magnesium acts as a reducing agent, and oxygen acts as an oxidizing agent.
Metals react with halogens. The product of this reaction is a metal halide, such as chloride.
Rice. 2. Potassium combustion in chlorine
Potassium burns in chlorine to form potassium chloride:
2K + Cl 2 \u003d 2KCl
Two potassium atoms donate one electron to the chlorine molecule. Potassium, increasing the oxidation state, plays the role of a reducing agent, and chlorine, lowering the oxidation state, plays the role of an oxidizing agent.
Many metals react with sulfur to form sulfides. In these reactions, metals also act as reducing agents, while sulfur will act as an oxidizing agent. Sulfur in sulfides is in the -2 oxidation state, i.e. it lowers its oxidation state from 0 to -2. For example, when heated, iron reacts with sulfur to form iron (II) sulfide:
Rice. 3. Interaction of iron with sulfur
Metals can also react with hydrogen, nitrogen, and other non-metals under certain conditions.
Only active metals, such as alkali and alkaline earth, react with water without heating. During these reactions, alkali is formed and hydrogen gas is released. For example, calcium reacts with water to form calcium hydroxide and hydrogen, and a large amount of heat is released:
Ca + 2H 2 O \u003d Ca (OH) 2 + H 2
Less active metals, such as iron and zinc, react with water only when heated to form metal oxide and hydrogen. For example:
Zn + H 2 O \u003d ZnO + H 2
In these reactions, the oxidizing agent is the hydrogen atom, which is part of the water.
Metals to the right of hydrogen in the voltage series do not react with water.
You already know that metals that are in the series of voltages to the left of hydrogen react with acids. In these reactions, metals donate electrons and act as a reducing agent. The oxidizing agent is hydrogen cations formed in acid solutions. For example, zinc reacts with hydrochloric acid:
Zn + 2HCl \u003d ZnCl 2 + H 2
Otherwise, the reactions of metals with nitric and concentrated sulfuric acids proceed. Almost no hydrogen is released in these reactions. We will talk about such interactions in the next lessons.
A metal can react with a salt solution if it is more active than the metal in the salt. For example, iron replaces copper from copper (II) sulfate:
Fe + CuSO 4 \u003d FeSO 4 + Cu
Iron is a reducing agent, copper cations are an oxidizing agent.
Let's try to explain why metals have common physical and chemical properties. To do this, consider a model of the internal structure of the metal.
Metal atoms have relatively large radii and a small number of external electrons. These electrons are weakly attracted to the nucleus, therefore, in chemical reactions, metals act as reducing agents, donating electrons from the external energy level.
In the nodes of the crystal lattice of metals there are not only neutral atoms, but also metal cations, because outer electrons move freely in the crystal lattice. In this case, the atoms, donating electrons, become cations, and cations, accepting electrons, turn into electrically neutral atoms.
Rice. 4. Model of the internal structure of the metal
A chemical bond that is formed as a result of the attraction of metal cations to freely moving electrons is called metallic.
The electrical and thermal conductivity of metals is explained by the presence of free electrons, which can be carriers of electric current and heat carriers. The plasticity of the metal is explained by the fact that under mechanical action the chemical bond does not break, because. a chemical bond is established not between specific atoms and cations, but between all metal cations with all free electrons in a metal crystal.
1. Mikityuk A.D. Collection of tasks and exercises in chemistry. Grades 8-11 / A.D. Mikityuk. - M.: Ed. "Exam", 2009.
2. Orzhekovsky P.A. Chemistry: 9th grade: textbook. for general inst. / P.A. Orzhekovsky, L.M. Meshcheryakova, L.S. Pontak. - M.: AST: Astrel, 2007. (§23)
3. Orzhekovsky P.A. Chemistry: 9th grade: textbook for general education. inst. / P.A. Orzhekovsky, L.M. Meshcheryakova, M.M. Shalashova. - M.: Astrel, 2013. (§6)
4. Rudzitis G.E. Chemistry: inorgan. chemistry. Organ. chemistry: textbook. for 9 cells. / G.E. Rudzitis, F.G. Feldman. - M .: Education, JSC "Moscow textbooks", 2009.
5. Khomchenko I.D. Collection of problems and exercises in chemistry for high school. - M .: RIA "New Wave": Publisher Umerenkov, 2008.
6. Encyclopedia for children. Volume 17. Chemistry / Chapter. ed. V.A. Volodin, leading. scientific ed. I. Leenson. - M.: Avanta +, 2003.
Additional web resources
1. A single collection of digital educational resources (video experiences on the topic) ().
2. Electronic version of the journal "Chemistry and Life" ().
Homework
p.41 Nos. A1, A2 from P.A. Orzhekovsky’s Textbook. "Chemistry: 9th grade" (M.: Astrel, 2013).
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.
|
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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Restorative properties- These are the main chemical properties characteristic of all metals. They manifest themselves in interaction with a wide variety of oxidants, including oxidants from the environment. In general, 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 oxidative 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.
Everyone knows that iron in humid air is 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 oxidizer 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 +. 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 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
