Active metals with water. Chemical properties of metals

Metals occupy the lower left corner of the Periodic Table. Metals belong to the families of s-elements, d-elements, f-elements and partially p-elements.

The most typical property of metals is their ability to donate electrons and become positively charged ions. Moreover, metals can only exhibit positive degree oxidation.

Me - ne = Me n +

1. Interaction of metals with non-metals.

A ) Interaction of metals with hydrogen.

Alkali and alkaline earth metals react directly with hydrogen, forming hydrides.

For example:

Ca + H 2 = CaH 2

Non-stoichiometric compounds with an ionic crystal structure are formed.

b) Interaction of metals with oxygen.

All metals except Au, Ag, Pt are oxidized by atmospheric oxygen.

Example:

2Na + O 2 = Na 2 O 2 (peroxide)

4K + O 2 = 2K 2 O

2Mg + O2 = 2MgO

2Cu + O 2 = 2CuO

c) Interaction of metals with halogens.

All metals react with halogens to form halides.

Example:

2Al + 3Br 2 = 2AlBr 3

These are mainly ionic compounds: MeHal n

d) Interaction of metals with nitrogen.

Alkali and alkaline earth metals interact with nitrogen.

Example:

3Ca + N2 = Ca3N2

Mg + N 2 = Mg 3 N 2 - nitride.

e) Interaction of metals with carbon.

Compounds of metals and carbon - carbides. They are formed by the interaction of melts with carbon. Active metals form stoichiometric compounds with carbon:

4Al + 3C = Al 4 C 3

Metals - d-elements form compounds of non-stoichiometric composition such as solid solutions: WC, ZnC, TiC - are used to produce superhard steels.

2. Interaction of metals with water.

Metals that have a more negative potential than the redox potential of water react with water.

Active metals react more actively with water, decomposing water and releasing hydrogen.

Na + 2H2O = H2 + 2NaOH

Less active metals slowly decompose water and the process is slowed down due to the formation of insoluble substances.

3. Interaction of metals with salt solutions.

This reaction is possible if the reacting metal is more active than the one in the salt:

Zn + CuSO 4 = Cu 0 ↓ + ZnSO 4

0.76 V., = + 0.34 V.

A metal with a more negative or less positive standard electrode potential displaces another metal from the solution of its salt.

4. Interaction of metals with alkali solutions.

Metals that produce amphoteric hydroxides or have high oxidation states in the presence of strong oxidizing agents can react with alkalis. When metals interact with alkali solutions, the oxidizing agent is water.

Example:

Zn + 2NaOH + 2H 2 O = Na 2 + H 2

1 Zn 0 + 4OH - - 2e = 2- oxidation

Zn 0 - reducing agent

1 2H 2 O + 2e = H 2 + 2OH - reduction

H 2 O - oxidizing agent

Zn + 4OH - + 2H 2 O = 2- + 2OH - + H 2

Metals with high oxidation states can interact with alkalis during fusion:

4Nb +5O 2 +12KOH = 4K 3 NbO 4 + 6H 2 O

5. Interaction of metals with acids.

These are complex reactions; the reaction products depend on the activity of the metal, the type and concentration of the acid, and the temperature.

Based on activity, metals are conventionally divided into active, medium activity and low activity.

Acids are conventionally divided into 2 groups:

Group I - acids with low oxidizing ability: HCl, HI, HBr, H 2 SO 4 (diluted), H 3 PO 4, H 2 S, the oxidizing agent here is H +. When interacting with metals, oxygen (H 2 ) is released. Metals with a negative electrode potential react with acids of the first group.

Group II - acids with high oxidizing ability: H 2 SO 4 (conc.), HNO 3 (diluted), HNO 3 (conc.). In these acids, the oxidizing agents are the acid anions: . The products of anion reduction can be very diverse and depend on the activity of the metal.

H 2 S - with active metals

H 2 SO 4 +6е S 0 ↓ - with metals of medium activity

SO 2 - with low-active metals

NH 3 (NH 4 NO 3) - with active metals

HNO 3 +4.5e N 2 O, N 2 - with medium activity metals

NO - with low-active metals

HNO 3 (conc.) - NO 2 - with metals of any activity.

If metals have variable valence, then with acids of group I the metals acquire a lower positive oxidation state: Fe → Fe 2+, Cr → Cr 2+. When interacting with acids of group II, the oxidation state is +3: Fe → Fe 3+, Cr → Cr 3+, and hydrogen is never released.

Some metals (Fe, Cr, Al, Ti, Ni, etc.) in solutions of strong acids, when oxidized, become covered with a dense oxide film, which protects the metal from further dissolution (passivation), but when heated, the oxide film dissolves and the reaction proceeds.

Slightly soluble metals with a positive electrode potential can dissolve in group I acids in the presence of strong oxidizing agents.

Moscow State Industrial University

Faculty of Applied Mathematics and Technical Physics

Department of Chemistry

Laboratory work

Chemical properties metals

Moscow 2012

Goal of the work. Studying properties s-, p-, d- metal elements (Mg, Al, Fe, Zn) and their compounds.

1. Theoretical part

All metals in their chemical properties are reducing agents, i.e. they donate electrons during a chemical reaction. Metal atoms relatively easily give up valence electrons and become positively charged ions.

1.1. Interaction of metals with simple substances

When metals interact with simple substances, non-metals usually act as oxidizing agents. Metals react with nonmetals to form binary compounds.

1. When interacting with oxygen metals form oxides:

2Mg + O 2 2MgO,

2Cu + O2 2CuO.

2. Metals react with halogens(F 2, Cl 2, Br 2, I 2) with the formation of salts of hydrohalic acids:

2Na + Br 2 = 2NaBr,

Ba + Cl 2 = BaCl 2,

2Fe + 3Cl 2 2FeCl3.

3. When metals interact with gray sulfides are formed (salts of hydrosulfide acid H 2 S):

4. C hydrogen Active metals react to form metal hydrides, which are salt-like substances:

2Na + H2 2NaH,

Ca+H2 CaH2.

In metal hydrides, hydrogen has an oxidation state of (-1).

Metals can also interact with other non-metals: nitrogen, phosphorus, silicon, carbon to form nitrides, phosphides, silicides, and carbides, respectively. For example:

3Mg + N 2 Mg 3 N 2,

3Ca+2P Ca 3 P 2 ,

2Mg + Si Mg2Si,

4Al + 3C Al 4 C 3 .

5. Metals can also interact with each other to form intermetallic compounds:

2Mg + Cu = Mg 2 Cu,

2Na + Sb = Na 2 Sb.

Intermetallic compounds(or intermetallic compounds) are compounds formed between themselves by elements that usually belong to metals.

1.2. Interaction of metals with water

The interaction of metals with water is an oxidation-reduction process in which the metal is a reducing agent and water acts as an oxidizing agent. The reaction proceeds according to the scheme:

Me + n H2O = Me(OH) n + n/2 H 2 .

Under normal conditions, alkali and alkaline earth metals react with water to form soluble bases and hydrogen:

2Na + 2H 2 O = 2NaOH + H 2,

Ca + 2H 2 O = Ca(OH) 2 + H 2.

Magnesium reacts with water when heated:

Mg + 2H 2 O Mg(OH) 2 + H 2 .

Iron and some other active metals react with hot water vapor:

3Fe + 4H 2 O Fe 3 O 4 + 4H 2 .

Metals with positive electrode potentials do not interact with water.

Do not interact with water 4 d-elements (except Cd), 5 d-elements and Cu (3 d-element).

1.3. Interaction of metals with acids

Based on the nature of their effect on metals, the most common acids can be divided into two groups.

1. Non-oxidizing acids: hydrochloric (hydrochloric, HCl), hydrobromic (HBr), hydroiodic (HI), hydrofluoric (HF), acetic (CH 3 COOH), dilute sulfuric (H 2 SO 4 (dil.)), dilute orthophosphoric (H 3 PO 4 (dil.)).

2. Oxidizing acids: nitric (HNO 3) in any concentration, concentrated sulfuric (H 2 SO 4 (conc.)), concentrated selenic (H 2 SeO 4 (conc.)).

Interaction of metals with non-oxidizing acids. The oxidation of metals by hydrogen ions H + in solutions of non-oxidizing acids occurs more vigorously than in water.

All metals that have a negative standard electrode potential, i.e. those in the electrochemical voltage range up to hydrogen displace hydrogen from non-oxidizing acids. The reaction proceeds according to the scheme:

Me + n H+=Me n + + n/2 H 2 .

For example:

2Al +6HCl = 2AlCl3 + 3H2,

Mg + 2CH 3 COOH = Mg(CH 3 COO) 2 + H 2,

2Ti + 6HCl = 2TiCl 3 + 3H 2.

Metals with variable oxidation states (Fe, Co, Ni, etc.) form ions in their lowest oxidation state (Fe 2+, Co 2+, Ni 2+ and others):

Fe + H 2 SO 4 (diluted) = FeSO 4 + H 2.

When some metals interact with non-oxidizing acids: HCl, HF, H 2 SO 4 (diluted), HCN, insoluble products are formed that protect the metal from further oxidation. Thus, the surface of lead in HCl (dil) and H 2 SO 4 (dil) is passivated by poorly soluble salts PbCl 2 and PbSO 4, respectively.

Interaction of metals with oxidizing acids. Sulfuric acid in a dilute solution is a weak oxidizing agent, but in a concentrated solution it is a very strong one. The oxidizing ability of concentrated sulfuric acid H 2 SO 4 (conc.) is determined by the SO 4 2 - anion, the oxidation potential of which is much higher than that of the H + ion. Concentrated sulfuric acid is a strong oxidizing agent due to sulfur atoms in the oxidation state (+6). In addition, a concentrated solution of H 2 SO 4 contains few H + ions, since in a concentrated solution it is weakly ionized. Therefore, when metals interact with H 2 SO 4 (conc.), hydrogen is not released.

Reacting with metals as an oxidizing agent, H 2 SO 4 (conc.) most often transforms into sulfur (IV) oxide (SO 2), and when interacting with strong reducing agents - into S or H 2 S:

Me + H 2 SO 4 (conc)  Me 2 (SO 4) n + H 2 O + SO 2 (S, H 2 S).

For ease of remembering, consider the electrochemical series of voltages, which looks like this:

Li, Rb, K, Cs, Ba, Sr, Ca, Na, Mg, Be, Al, Mn, Zn, Cr, Fe, Cd, Co, Ni, Sn, Pb, (H), Cu, Hg, Ag, Pt, Au.

In table 1. The products of the reduction of concentrated sulfuric acid when interacting with metals of various activities are presented.

Table 1.

Products of interaction of metals with concentrated

sulfuric acid

Cu + 2H 2 SO 4 (conc) = CuSO 4 + SO 2 + 2H 2 O,

4Mg + 5H 2 SO 4 (conc) = 4MgSO 4 + H 2 S + 4H 2 O.

For metals of intermediate activity (Mn, Cr, Zn, Fe), the ratio of reduction products depends on the acid concentration.

The general trend is: the higher the concentration H2SO4, the deeper the recovery proceeds.

This means that formally each sulfur atom from H 2 SO 4 molecules can take not only two electrons from the metal (and go into ), but also six electrons (and go to) and even eight (and go to ):

Zn + 2H 2 SO 4 (conc) = ZnSO 4 + SO 2 + 2H 2 O,

3Zn + 4H 2 SO 4 (conc) = 3ZnSO 4 + S + 4H 2 O,

4Zn + 5H 2 SO 4 (conc) = 4ZnSO 4 + H 2 S + 4H 2 O.

Lead reacts with concentrated sulfuric acid to form soluble lead (II) hydrogen sulfate, sulfur oxide (IV) and water:

Pb + 3H 2 SO 4 = Pb(HSO 4) 2 + SO 2 + 2H 2 O.

Cold H 2 SO 4 (conc) passivates some metals (for example, iron, chromium, aluminum), which allows the acid to be transported in steel containers. When heated strongly, concentrated sulfuric acid reacts with these metals:

2Fe + 6H 2 SO 4 (conc) Fe 2 (SO 4) 3 + 3SO 2 + 6H 2 O.

Interaction of metals with nitric acid. The oxidizing capacity of nitric acid is determined by the NO 3 - anion, the oxidation potential of which is much higher than that of H + ions. Therefore, when metals interact with HNO 3, hydrogen is not released. Nitrate ion NO 3 - containing nitrogen in the oxidation state (+ 5), depending on conditions (acid concentration, nature of the reducing agent, temperature) can accept from one to eight electrons. The reduction of the NO 3 anion can occur with the formation of various substances according to the following schemes:

NO 3  + 2H + + e = NO 2 + H 2 O,

NO 3  + 4H + + 3e = NO + 2H 2 O,

2NO 3  + 10H + + 8e = N 2 O + 5H 2 O,

2NO 3  + 12H + + 10e = N 2 + 6H 2 O,

NO 3  + 10H + + 8e = NH 4 + + 3H 2 O.

Nitric acid has oxidizing power at any concentration. All other things being equal, the following trends appear: the more active the metal reacts with the acid, and the lower the concentration of the nitric acid solution,the more deeply it is restored.

This can be illustrated with the following diagram:

, ,
,
,

Acid concentration

Metal activity

The oxidation of substances with nitric acid is accompanied by the formation of a mixture of products of its reduction (NO 2, NO, N 2 O, N 2, NH 4 +), the composition of which is determined by the nature of the reducing agent, temperature and concentration of the acid. Oxides NO 2 and NO predominate among the products. Moreover, when interacting with a concentrated solution of HNO 3, NO 2 is more often released, and with a diluted solution, NO is released.

Equations of redox reactions involving HNO 3 are drawn up conditionally, with the inclusion of only one reduction product, which is formed in larger quantities:

Me + HNO 3  Me (NO 3) n + H 2 O + NO 2 (NO, N 2 O, N 2, NH 4 +).

For example, in a gas mixture formed by the action of a sufficiently active metal, zinc (
= - 0.76 B) concentrated (68%) nitric acid, predominant - NO 2, 40% - NO; 20% – N 2 O; 6% - N 2. Very dilute (0.5%) nitric acid is reduced to ammonium ions:

Zn + 4HNO 3 (conc.) = Zn(NO 3) 2 + 2NO 2 + 2H 2 O,

3Zn + 8HNO 3 (40%) = 3Zn(NO 3) 2 + 2NO + 4H 2 O,

4Zn + 10HNO 3 (20%) = 4Zn(NO 3) 2 + N 2 O + 5H 2 O,

5Zn + 12HNO 3 (6%) = 5Zn(NO 3) 2 + N 2 + 6H 2 O,

4Zn + 10HNO3 (0.5%) = 4Zn(NO3)2 + NH4NO3 + 3H2O.

With low-active metal copper (
= + 0.34B) reactions proceed according to the following schemes:

Cu + 4HNO 3 (conc) = Cu(NO 3) 2 + 2NO 2 + 2H 2 O,

3Cu + 8HNO 3 (dil) = 3 Cu(NO 3) 2 + 2NO + 4H 2 O.

Almost all metals dissolve in concentrated HNO 3 except Au, Ir, Pt, Rh, Ta, W, Zr. And such metals as Al, Be, Bi, Co, Cr, Fe, Nb, Ni, Pb, Th, U, as well as stainless steels passivated with acid to form stable oxide films that adhere tightly to the metal surface and protect it from further oxidation. However, Al and Fe begin to dissolve when heated, and Cr is resistant to even hot HNO 3:

Fe + 6HNO3 Fe(NO 3) 3 + 3NO 2 + 3H 2 O.

Metals that are characterized by high degrees oxidation (+6, +7, +8), with concentrated nitric acid form oxygen-containing acids. In this case, HNO 3 is reduced to NO, for example:

3Re + 7HNO 3 (conc) = 3HReO 4 + 7NO + 2H 2 O.

In very dilute HNO 3 there are no longer any HNO 3 molecules, only H + and NO 3  ions exist. Therefore, a very dilute acid (~ 3-5%) interacts with Al and does not transfer Cu and other low-active metals into solution:

8Al + 30HNO 3 (very dilute) = 8Al(NO 3) 3 + 3NH 4 NO 3 + 9H 2 O.

A mixture of concentrated nitrogen and hydrochloric acid(1:3) is called aqua regia. It dissolves Au and platinum metals (Pd, Pt, Os, Ru). For example:

Au + HNO 3 (conc.) + 4HCl = H + NO + 2H 2 O.

These metals dissolve in HNO 3 and in the presence of other complexing agents, but the process proceeds very slowly.

The structure of metal atoms determines not only the characteristic physical properties simple substances - metals, but also their general chemical properties.

With great diversity, all chemical reactions of metals are redox and can be of only two types: combination and substitution. Metals are capable of donating electrons during chemical reactions, that is, being reducing agents and exhibiting only a positive oxidation state in the resulting compounds.

IN general view this can be expressed by the diagram:
Me 0 – ne → Me +n,
where Me is a metal - a simple substance, and Me 0+n is a metal, a chemical element in a compound.

Metals are capable of donating their valence electrons to non-metal atoms, hydrogen ions, and ions of other metals, and therefore will react with non-metals - simple substances, water, acids, salts. However, the reducing ability of metals varies. The composition of the reaction products of metals with various substances depends on the oxidizing ability of the substances and the conditions under which the reaction occurs.

At high temperatures, most metals burn in oxygen:

2Mg + O2 = 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 formed in some cases. Under normal conditions, alkali metals, as well as calcium, strontium, and barium, interact very actively with water. 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 a metal reacts with an acid, it is part of the resulting salt. When a metal interacts with acid solutions, it can be oxidized by hydrogen ions present in the solution. The abbreviated ionic equation can be written in general form as follows:

Me + nH + → Me n + + H 2

The anions of oxygen-containing acids, such as concentrated sulfuric and nitric, have stronger oxidizing properties than hydrogen ions. Therefore, those metals that are not able to be oxidized by hydrogen ions, for example, copper and silver, react with these acids.

When metals interact with salts, a substitution reaction occurs: electrons from the atoms of the replacing – more active metal – pass to the ions of the replaced – less active metal. Then the network replaces metal with metal in salts. These reactions are not reversible: if metal A displaces metal B from the salt solution, then metal B will not displace metal A from the 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 (activities) 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

Metals located to the left in this row are more active and are able to displace the following metals from salt solutions.

Hydrogen is included in the electrochemical voltage series of metals as the only non-metal that shares with metals general property- 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 = ZnCl 2 + H 2 + Q

Metals that come before hydrogen in the electrochemical voltage series displace it from solutions of many acids (hydrochloric, sulfuric, etc.), but all those following it, for example, copper, do not displace it.

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Metals mean a group of elements, which are presented in the form of the simplest substances. They have characteristic properties, namely high electrical and thermal conductivity, positive temperature coefficient of resistance, high ductility and metallic luster.

Note that out of 118 chemical elements, which were opened on this moment, metals should include:

• among the group of alkaline earth metals there are 6 elements;
• among alkali metals there are 6 elements;
• among transition metals 38;
• in the group of light metals 11;
• There are 7 elements among semimetals,
• 14 among lanthanides and lanthanum,
• 14 in the group of actinides and sea anemones,
• Beryllium and magnesium are outside the definition.

Based on this, 96 elements are classified as metals. Let's take a closer look at what metals react with. Because on the outside electronic level Most metals have a small number of electrons from 1 to 3, then in most of their reactions they can act as reducing agents (that is, they give up their electrons to other elements).

Reactions with the simplest elements

• Except for gold and platinum, absolutely all metals react with oxygen. Note also that the reaction occurs with silver at high temperatures, but silver(II) oxide is not formed at normal temperatures. Depending on the properties of the metal, oxides, superoxides and peroxides are formed as a result of reaction with oxygen.

Here are examples of each chemical education:

1. lithium oxide – 4Li+O 2 =2Li 2 O;
2. potassium superoxide – K+O 2 =KO 2;
3. sodium peroxide – 2Na+O 2 =Na 2 O 2.

In order to obtain an oxide from a peroxide, it must be reduced with the same metal. For example, Na 2 O 2 +2Na=2Na 2 O. With low- and medium-active metals, a similar reaction will occur only when heated, for example: 3Fe+2O 2 =Fe 3 O 4.

• Metals can react with nitrogen only with active metals, but when room temperature only lithium can interact, forming nitrides - 6Li+N 2 = 2Li 3 N, however, when heated, this occurs chemical reaction 2Al+N 2 =2AlN, 3Ca+N 2 =Ca 3 N 2.
• Absolutely all metals react with sulfur, as with oxygen, with the exception of gold and platinum. Note that iron can only react when heated with sulfur, forming sulfide: Fe+S=FeS
• Only active metals can react with hydrogen. These include metals of groups IA and IIA, except beryllium. Such reactions can only occur when heated, forming hydrides.

  Since the oxidation state of hydrogen is considered? 1, the metals in this case act as reducing agents: 2Na + H 2 = 2NaH.

• The most active metals also react with carbon. As a result of this reaction, acetylenides or methanides are formed.

Let's consider what metals react with water and what do they produce as a result of this reaction? Acetylenes, when reacting with water, will give acetylene, and methane will be obtained as a result of the reaction of water with methanides. Here are examples of these reactions:

1. Acetylene – 2Na+2C= Na 2 C 2 ;
2. Methane - Na 2 C 2 +2H 2 O=2NaOH+C 2 H 2.

Reaction of acids with metals

Metals can also react differently with acids. Only those metals that are in the series of electrochemical activity of metals up to hydrogen react with all acids.

Let's give an example of a substitution reaction that shows what metals react with. In another way, this reaction is called redox: Mg+2HCl=MgCl 2 +H 2 ^.

Some acids can also interact with metals that come after hydrogen: Cu+2H 2 SO 4 =CuSO 4 +SO 2 ^+2H 2 O.

Note that such a dilute acid can react with a metal according to the classical scheme shown: Mg + H 2 SO 4 = MgSO 4 + H 2 ^.

Metal ratio reaction equations:

• a) to simple substances: oxygen, hydrogen, halogens, sulfur, nitrogen, carbon;
• b) to complex substances: water, acids, alkalis, salts.

1. Metals include s-elements of groups I and II, all s-elements, p-elements of group III (except boron), as well as tin and lead (group IV), bismuth (group V) and polonium (group VI). Most metals have 1-3 electrons in their outer energy level. For atoms of d-elements, within periods, the d-sublevels of the pre-outer layer are filled from left to right.
2. The chemical properties of metals are determined by the characteristic structure of their outer electron shells.

Within a period, as the nuclear charge increases, the radii of atoms with the same number of electron shells decrease. The atoms of alkali metals have the largest radii. The smaller the radius of the atom, the greater the ionization energy, and the larger the radius of the atom, the less the ionization energy. Since metal atoms have the largest atomic radii, they are characterized mainly by low values ​​of ionization energy and electron affinity. Free metals exhibit exclusively reducing properties.

3) Metals form oxides, for example:

Only alkali and alkaline earth metals react with hydrogen, forming hydrides:

Metals react with halogens, forming halides, with sulfur - sulfides, with nitrogen - nitrides, with carbon - carbides.

With an increase in the algebraic value of the standard electrode potential of a metal E 0 in the voltage series, the ability of the metal to react with water decreases. So, iron reacts with water only at very high temperature:

Metals with a positive standard electrode potential, that is, those standing after hydrogen in the voltage series, do not react with water.

Reactions of metals with acids are characteristic. Metals with a negative E0 value displace hydrogen from solutions of HCl, H2S04, H3P04, etc.

A metal with a lower value of E 0 displaces a metal with great value E 0 from salt solutions:

The most important calcium compounds obtained industrially, their chemical properties and methods of production.

Calcium oxide CaO is called quicklime. It is obtained by burning limestone CaC0 3 --> CaO + CO, at a temperature of 2000° C. Calcium oxide has the properties of a basic oxide:

a) reacts with water to release large quantity heat:

CaO + H 2 0 = Ca (OH) 2 (slaked lime).

b) reacts with acids to form salt and water:

CaO + 2HCl = CaCl 2 + H 2 O

CaO + 2H + = Ca 2+ + H 2 O

c) reacts with acid oxides to form a salt:

CaO + C0 2 = CaC0 3

Calcium hydroxide Ca(OH) 2 is used in the form of slaked lime, lime milk and lime water.

Lime milk is a slurry formed by mixing excess slaked lime with water.

Lime water- a transparent solution obtained by filtering lime milk. Used in the laboratory to detect carbon (IV) monoxide.

Ca(OH) 2 + CO 2 = CaCO 3 + H 2 O

With prolonged passage of carbon monoxide (IV), the solution becomes transparent, as an acidic salt is formed, soluble in water:

CaC0 3 + C0 2 + H 2 O = Ca(HCO 3 ) 2

If the resulting transparent solution of calcium bicarbonate is heated, then turbidity occurs again, as a precipitate of CaC0 3 precipitates: