Inorganic substances are simple and complex. Simple substances are divided into metals (K, Na, Li) and non-metals (O, Cl, P). Complex substances are divided into oxides, hydroxides (bases), salts and acids.
oxides
oxides- compounds of a chemical element (metal or non-metal) with oxygen (oxidation state -2), while oxygen is associated with a less electronegative element.
Allocate:
1. Acid oxides- oxides showing acid properties. Formed by non-metals and oxygen. Examples: SO3, SO2, CO2, P2O5, N2O5.
2. Amphoteric oxides- oxides, which can exhibit both basic and acidic properties (this property is called amphoteric). Examples: Al2O3, CrO3, ZnO, BeO, PbO.
3. Basic oxides- metal oxides, while the metals exhibit an oxidation state of +1 or +2. Examples: K2O, MgO, CaO, BaO, Li2O, Na2O.
4. Non-salt-forming oxides- practically do not react, do not have the corresponding acids and hydroxides. Examples: CO, NO.
Chemical properties basic oxides
1. Interaction with water
Only oxides of alkali and alkaline earth metals enter into the reaction, the hydroxides of which form a soluble base
basic oxide + water → alkali
K2O + H2O → 2KOH
CaO + H2O → Ca(OH)2
2. Interaction with acid
basic oxide + acid → salt + water
MgO + H2SO4 → MgSO4 + H2O
Na2O + H2S(ex) → 2NaHS + H2O
MgO(ex) + HCl → Mg(OH)Cl
3. Interaction with acidic or amphoteric oxides
basic oxide + acidic/ amphoteric oxide→ salt
In this case, the metal in the basic oxide becomes a cation, and the acid/amphoteric oxide becomes an anion (acid residue). Reactions between solid oxides occur when heated. Water-insoluble basic oxides do not interact with gaseous acidic oxides.
BaO + SiO2 (t) → BaSiO3
K2O + ZnO (t) → K2ZnO2
FeO + CO2 ≠
4. Interaction with amphoteric hydroxides
basic oxide + amphoteric hydroxide→ salt + water
Na2O + 2Al(OH)3 (t) → 2NaAlO2 + 3H2O
5. Temperature decomposition of noble metal oxides and mercury
2Ag2O (t) → 4Ag + O2
2HgO (t) → 2Hg + O2
6. Interaction with carbon (C) or hydrogen (H2) at high temperature.
When reducing oxides of alkali, alkaline earth metals and aluminum in this way, not the metal itself is released, but its carbide.
FeO + C (t) → Fe + CO
3Fe2O3 + C (t) → 2Fe3O4 + CO
CaO + 3C (t) → CaC2 + CO
CaO + 2H2 (t) → CaH2 + H2O
7. Active metals reduce the less active of their oxides at high temperature
CuO + Zn (t) → ZnO + Cu
8. Oxygen oxidizes lower oxides to higher ones.
Oxides of alkali and alkaline earth metals are converted into peroxides
4FeO + O2 (t) → 2Fe2O3
2BaO + O2 (t) → 2BaO2
2NaO + O2 (t) → 2Na2O2
Chemical properties of acid oxides
1. Interaction with water
acid oxide + water → acid
SO3+ H2O → H2SO4
SiO2 + H2O ≠
Some oxides do not have the corresponding acids, in which case a disproportionation reaction occurs
2NO2 + H2O → HNO3 + HNO2
3NO2 + H2O (t) → 2HNO3 + NO
2ClO2 + H2O → HClO3 + HClO2
6ClO2 + 3H2O (t) → 5HClO3 + HCl
Depending on the number of water molecules attached to P2O5, three different acids are formed - metaphosphoric HPO3, pyrophosphoric H4P2O7 or orthophosphoric H3PO4.
P2O5 + H2O → 2HPO3
P2O5 + 2H2O → H4P2O7
P2O5 + 3H2O → 2H3PO4
Chromium oxide corresponds to two acids - chromic H2CrO4 and dichromic H2Cr2O7(III)
CrO3 + H2O → H2CrO4
2CrO3 + H2O → H2Cr2O7
2. Interaction with bases
acid oxide + base → salt + water
Insoluble acidic oxides react only during fusion, while soluble oxides react under normal conditions.
SiO2 + 2NaOH (t) → Na2SiO3 + H2O
With an excess of oxide, an acid salt is formed.
CO2(ex) + NaOH → NaHCO3
P2O5(ex) + 2Ca(OH)2 → 2CaHPO4 + H2O
P2O5(ex) + Ca(OH)2 + H2O → Ca(H2PO4)2
With an excess of base, a basic salt is formed.
CO2 + 2Mg(OH)2(g) → (MgOH)2CO3 + H2O
Oxides that do not have the corresponding acids enter into a disproportionation reaction and form two salts.
2NO2 + 2NaOH → NaNO3 + NaNO2 + H2O
2ClO2 + 2NaOH → NaClO3 + NaClO2 + H2O
CO2 reacts with some amphoteric hydroxides (Be(OH)2, Zn(OH)2, Pb(OH)2, Cu(OH)2) to form a basic salt and water.
CO2 + 2Be(OH)2 → (BeOH)2CO3↓ + H2O
CO2 + 2Cu(OH)2 → (CuOH)2CO3↓ + H2O
3. Interaction with basic or amphoteric oxide
acid oxide + basic/amphoteric oxide → salt
Reactions between solid oxides occur during fusion. Amphoteric and water-insoluble basic oxides interact only with solid and liquid acidic oxides.
SiO2 + BaO (t) → BaSiO3
3SO3 + Al2O3 (t) → Al2(SO4)3
4. Interaction with salt
acidic not volatile oxide+ salt (t) → salt + acid volatile oxide
SiO2 + CaCO3 (t) → CaSiO3 + CO2
P2O5 + Na2CO3 → 2Na3PO4 + 2CO2
5. Acid oxides do not react with acids, but P2O5 reacts with anhydrous oxygen-containing acids.
This produces HPO3 and the anhydride of the corresponding acid
P2O5 + 2HClO4(anhydrous) → Cl2O7 + 2HPO3
P2O5 + 2HNO3(anhydrous) → N2O5 + 2HPO3
6. Enter into redox reactions.
1. Recovery
At high temperatures, some non-metals can reduce oxides.
CO2 + C (t) → 2CO
SO3 + C → SO2 + CO
H2O + C (t) → H2 + CO
Magnesium is often used to reduce nonmetals from their oxides.
CO2 + 2Mg → C + 2MgO
SiO2 + 2Mg (t) → Si + 2MgO
N2O + Mg (t) → N2 + MgO
2. Lower oxides are converted into higher ones when interacting with ozone (or oxygen) at high temperature in the presence of a catalyst
NO + O3 → NO2 + O2
SO2 + O3 → SO3 + O2
2NO2 + O3 → N2O5 + O2
2CO + O2 (t) → 2CO2
2SO2 + O2 (t, kat) → 2SO3
P2O3 + O2 (t) → P2O5
2NO + O2 (t) → 2NO2
2N2O3 + O2 (t) → 2N2O4
3. Oxides also enter into other redox reactions
SO2 + NO2 → NO + SO3 4NO2 + O2 + 2H2O → 4HNO3
2SO2 + 2NO → N2 + 2SO3 2N2O5 → 4NO2 + O2
SO2 + 2H2S → 3S↓ + 2H2O 2NO2 (t)→ 2NO + O2
2SO2 + O2 + 2H2O → 2H2SO4 3N2O + 2NH3 → 4N2 + 3H2O
2CO2 + 2Na2O2 → 2Na2CO3 + O2 10NO2 +8P → 5N2 + 4P2O5
N2O + 2Cu (t) → N2 + Cu2O
2NO + 4Cu (t) → N2 + 2Cu2O
N2O3 + 3Cu (t) → N2 + 3CuO
2NO2 + 4Cu (t) → N2 + 4CuO
N2O5 + 5Cu (t) → N2 + 5CuO
Chemical properties of amphoteric oxides
1. Do not interact with water
amphoteric oxide + water ≠
2. Interaction with acids
amphoteric oxide + acid → salt + water
Al2O3 + 3H2SO4 → Al2(SO4)3 + 3H2O
With an excess of a polybasic acid, an acid salt is formed
Al2O3 + 6H3PO4(ex) → 2Al(H2PO4)3 + 3H2O
With an excess of oxide, a basic salt is formed
ZnO(ex) + HCl → Zn(OH)Cl
Double oxides form two salts
Fe3O4 + 8HCl → FeCl2 + 2FeCl3 + 4H2O
3. Interaction with acid oxide
amphoteric oxide + acid oxide → salt
Al2O3 + 3SO3 → Al2(SO4)3
4. Interaction with alkali
amphoteric oxide + alkali → salt + water
When fused, an average salt and water are formed, and in solution - a complex salt
ZnO + 2NaOH(tv) (t) → Na2ZnO2 + H2O
ZnO + 2NaOH + H2O → Na2
5. Interaction with the basic oxide
amphoteric oxide + basic oxide (t) → salt
ZnO + K2O (t) → K2ZnO2
6. Interaction with salts
amphoteric oxide + salt (t) → salt + volatile acidic oxide
Amphoteric oxides displace volatile acidic oxides from their salts during fusion
Al2O3 + K2CO3 (t) → KAlO2 + CO2
Fe2O3 + Na2CO3 (t) → 2NaFeO2 + CO2
Chemical properties of bases
Bases are substances that contain a metal cation and a hydroxide anion. Bases are soluble (alkalis - NaOH, KOH, Ba(OH)2) and insoluble (Al2O3, Mg(OH)2).
1. Soluble base + indicator → color change
When an indicator is added to a base solution, its color changes:
Colorless phenolphthalein - raspberry
Purple litmus - blue
Methyl orange - yellow
2. Interaction with acid (neutralization reaction)
base + acid → salt + water
According to the reaction, medium, acidic or basic salts can be obtained. With an excess of a polybasic acid, an acid salt is formed, with an excess of a polyacid base, a basic salt.
Mg(OH)2 + H2SO4 → MGSO4 + 2H2O
Mg(OH)2 + 2H2SO4 → MG(HSO4)2 + 2H2O
2Mg(OH)2 + H2SO4 → (MgOH)2SO4 + 2H2O
3. Interaction with acid oxides
base + acid oxide → salt + water
6NH4OH + P2O5 → 2(NH4)3PO4 + 3H2O
4. Interaction of alkali with amphoteric hydroxide
alkali + amphoteric hydroxide → salt + water
In this reaction, amphoteric hydroxide exhibits acidic properties. During the reaction in the melt, an average salt and water are obtained, and in a solution, a complex salt. Hydroxides of iron (III) and chromium (III) dissolve only in concentrated alkali solutions.
2KOH(tv) + Zn(OH)2 (t) → K2ZnO2 + 2H2O
KOH + Al(OH)3 → K
3NaOH(conc) + Fe(OH)3 → Na3
5. Interaction with amphoteric oxide
alkali + amphoteric oxide → salt + water
2NaOH(s) + Al2O3 (t) → 2NaAlO2 + H2O
6NaOH + Al2O3 + 3H2O → 2Na3
6. Interaction with salt
An ion exchange reaction takes place between the base and the salt. It occurs only when a precipitate is formed or when gas is released (during the formation of NH4OH).
A. Reaction between a soluble base and a soluble acid salt
soluble base + soluble acid salt → medium salt + water
If the salt and base are formed by different cations, then two middle salts are formed. In the case of acidic ammonium salts, an excess of alkali leads to the formation of ammonium hydroxide.
Ba(OH)2 + Ba(HCO3)2 → 2BaCO3↓ + 2H2O
2NaOH(ex) + NH4HS → Na2S + NH4OH + H2O
B. Reaction of a soluble base with a soluble intermediate or basic salt.
Several scenarios are possible
soluble base + soluble intermediate/basic salt → insoluble salt↓ + base
→ salt + insoluble base↓
→ salt + weak electrolyte NH4OH
→ no reaction
Reactions occur between soluble bases and an average salt only if the result is an insoluble salt, or an insoluble base, or a weak electrolyte NH4OH
NaOH + KCl ≠ no reaction
If the original salt is formed by a polyacid base, with a lack of alkali, a basic salt is formed
Under the action of alkalis on salts of silver and mercury (II), not their hydroxides are released, which dissolve at 25 ° C, but insoluble oxides Ag2O and HgO.
7. Decomposition at temperature
basic hydroxide (t) → oxide + water
Ca(OH)2 (t) → CaO + H2O
NaOH(t)≠
Some bases (AgOH, Hg(OH)2 and NH4OH) decompose even at room temperature
LiOH (t) → Li2O + H2O
NH4OH (25C) → NH3 + H2O
8. Interaction of alkali and transition metal
alkali + transition metal → salt + H2
2Al + 2KOH + 6H2O → 2K +3H2
Zn + 2NaOH(tv) (t) → Na2ZnO2 + H2
Zn + 2NaOH + 2H2O → Na2 + H2
9. Interaction with non-metals
Alkalis interact with some non-metals - Si, S, P, F2, Cl2, Br2, I2. In this case, two salts are often formed as a result of disproportionation.
Si + 2KOH + H2O → K2SiO3 + 2H2
3S + 6KOH (t) → 2K2S + K2SO3 + 3H2O
Cl2 +2KOH(conc) → KCl + KClO + H2O (for Br, I)
3Cl2 + 6KOH(conc) (t)→ 5KCl + KClO3 +3H2O (for Br, I)
Cl2 + Ca(OH)2 → CaOCl2 + H2O
4F2 + 6NaOH(dec) → 6NaF + OF2 + O2 + 3H2O
4P + 3NaOH + 3H2O → 3NaH2PO2 + PH3
Hydroxides with reducing properties are able to be oxidized by oxygen
4Fe(OH)2 + O2 + 2H2O → 4Fe(OH)3 (=Cr)
Chemical properties of acids
1. Changing the color of the indicator
soluble acid + indicator → color change
Violet litmus and methyl orange turn red, phenolphthalein becomes transparent
2. Interaction with bases (neutralization reaction)
acid + base → salt + water
H2SO4 + Mg(OH)2 → MgSO4 + 2H2O
3. Interaction with the basic oxide
acid + basic oxide → salt + water
2HCl + CuO → CuCl2 + H2O
4. Interaction with amphoteric hydroxides with the formation of medium, acidic or basic salts
acid + amphoteric hydroxide → salt + water
2HCl + Be(OH)2 → BeCl2 + 2H2O
H3PO4() + Zn(OH)2 → ZNHPO4 + 2H2O
HCl + Al(OH)3() → Al(OH)2Cl + H2O
5. Interaction with amphoteric oxides
acid + amphoteric oxide → salt + water
H2SO4 + ZnO → ZnSO4 + H2O
6. Interaction with salts
General reaction scheme: acid + salt → salt + acid
An ion exchange reaction occurs, which goes to completion only in the case of gas formation or precipitation.
For example: HCl + AgNO3 → AgCl↓ + HNO3
2HBr + K2SiO3 → 2KBr + H2SiO3↓
A. Reaction with a salt of a more volatile or weak acid to form a gas
HCl + NaHS → NaCl + H2S
B. Interaction strong acid and salts of a strong or medium acid to form an insoluble salt
strong acid + strong/medium acid salt → insoluble salt + acid
Non-volatile orthophosphoric acid displaces strong but volatile hydrochloric and nitric acids from their salts, provided that an insoluble salt is formed
B. The interaction of an acid with a basic salt of the same acid
acid1 + basic salt of acid1 → medium salt + water
HCl + Mg(OH)Cl → MgCl2 + H2O
D. The interaction of a polybasic acid with an average or acid salt of the same acid to form an acid salt of the same acid containing more hydrogen atoms
polybasic acid1 + medium/acid salt of acid1 → acid salt of acid1
H3PO4 + Ca3(PO4)2 → 3CaHPO4
H3PO4 + CaHPO4 → Ca(H2PO4)2
E. Interaction of hydrosulfide acid with salts of Ag, Cu, Pb, Cd, Hg with the formation of insoluble sulfide
acid H2S + salt Ag, Cu, Pb, Cd, Hg → Ag2S/CuS/PbS/CdS/HgS↓ + acid
H2S + CuSO4 → CuS↓ + H2SO4
E. Reaction of an acid with an average or complex salt with an amphoteric metal in an anion
a) in the case of a lack of acid, an average salt and amphoteric hydroxide are formed
acid + average/complex salt in amphoteric metal in anion → average salt + amphoteric hydroxide
b) in the case of an excess of acid, two medium salts and water are formed
acid + average/complex salt with amphoteric metal in anion → average salt + average salt + water
G. In some cases, acids with salts enter into redox reactions or complex formation reactions:
H2SO4(conc) and I‾/Br‾ (H2S and I2/SO2 and Br2 products)
H2SO4(conc) and Fe² + (SO2 and Fe³ + products)
HNO3 dil/conc and Fe² + (NO/NO2 and Fe³ + products)
HNO3 dil/conc and SO3²‾/S²‾ (NO/NO2 and SO4²‾/S or SO4²‾ products)
HClconc and KMnO4/K2Cr2O7/KClO3 (products Cl2 and Mn² + /Cr² + /Cl‾)
3. Interaction of concentrated sulfuric acid with a solid salt
Non-volatile acids can displace volatiles from their solid salts.
7. Interaction of acid with metal
A. Interaction of an acid with metals standing in a row before or after hydrogen
acid + metal to H2 → metal sel in the minimum oxidation state + H2
Fe + H2SO4(dil) → FeSO4 + H2
acid + metal after H2 ≠ no reaction
Cu + H2SO4(dil) ≠
B. Interaction of concentrated sulfuric acid with metals
H2SO4(conc) + Au, Pt, Ir, Rh, Ta ≠ no reaction
H2SO4(conc) + alkali/alkaline earth metal and Mg/Zn → H2S/S/SO2 (depending on conditions) + metal sulfate at maximum oxidation state + H2O
Zn + 2H2SO4(conc) (t1) → ZnSO4 + SO2 + 2H2O
3Zn + 4H2SO4(conc) (t2>t1)→ 3ZnSO4 + S↓ + 4H2O
4Zn + 5H2SO4(conc) (t3>t2) → 4ZnSO4 + H2S + 4H2O
H2SO4(conc) + other metals → SO2 + metal sulfate in maximum oxidation state + H2O
Cu + 2H2SO4(conc) (t) → CuSO4 + SO2 + 2H2O
2Al + 6H2SO4(conc) (t) → Al2(SO4)3 + 3SO2 + 6H2O
B. Interaction of concentrated nitric acid with metals
HNO3(conc) + Au, Pt, Ir, Rh, Ta, Os ≠ no reaction
HNO3(conc) + Pt ≠
HNO3(conc) + alkali/alkaline earth metal → N2O + metal nitrate at maximum oxidation state + H2O
4Ba + 10HNO3(conc) → 4Ba(NO3)2 + N2O + 5H2O
HNO3(conc) + other metals at temperature → NO2 + metal nitrate at maximum oxidation state + H2O
Ag + 2HNO3(conc) → AgNO3 + NO2 + H2O
It interacts with Fe, Co, Ni, Cr and Al only when heated, since under normal conditions these metals are passivated by nitric acid - they become chemically resistant
D. Reaction of dilute nitric acid with metals
HNO3(diff) + Au, Pt, Ir, Rh, Ta ≠ no reaction
Very passive metals (Au, Pt) can be dissolved with aqua regia - a mixture of one volume of concentrated nitric acid with three volumes of concentrated of hydrochloric acid. The oxidizing agent in it is atomic chlorine, which splits off from nitrosyl chloride, which is formed as a result of the reaction: HNO3 + 3HCl → 2H2O + NOCl + Cl2
HNO3(dil) + alkali/alkaline earth metal → NH3(NH4NO3) + metal nitrate at maximum oxidation state + H2O
NH3 is converted to NH4NO3 in excess of nitric acid
4Ca + 10HNO3(diff) → 4Ca(NO3)2 + NH4NO3 + 3H2O
HNO3(razb) + metal in the voltage series up to H2 → NO/N2O/N2/NH3 (depending on conditions) + metal nitrate in the maximum oxidation state + H2O
With the rest of the metals, standing in a series of voltages up to hydrogen and non-metals, HNO3 (dil) forms salt, water and, mainly NO, but, depending on the conditions, both N2O, and N2, and NH3 / NH4NO3 (the more dilute the acid is , the lower the degree of nitrogen oxidation in the evolved gaseous product)
3Zn + 8HNO3(razb) → 3Zn(NO3)2 + 2NO + 4H2O
4Zn + 10HNO3(diff) → 4Zn(NO3)2 + N2O + 5H2O
5Zn + 12HNO3(diff) → 5Zn(NO3)2 + N2 + 6H2O
4Zn + 10HNO3(very dilute) → 4Zn(NO3)2 + NH4NO3 + 3H2O
HNO3(razb) + metal after H2 → NO + metal nitrate at maximum oxidation state + H2O
With low-active metals standing after H2, HNO3razb forms salt, water and NO
3Cu + 8HNO3(diff) → 3Cu(NO3)2 + 2NO + 4H2O
8. Decomposition of acids at temperature
acid (t) → oxide + water
H2CO3 (t) → CO2 + H2O
H2SO3 (t) → SO2 + H2O
H2SiO3 (t) → SiO2 + H2O
2H3PO4 (t) → H4P2O7 + H2O
H4P2O7 (t) → 2HPO3 + H2O
4HNO3 (t) → 4NO2 + O2 + 2H2O
3HNO2 (t) → HNO3 + 2NO + H2O
2HNO2 (t) → NO2 + NO + H2O
3HCl (t) → 2HCl + HClO3
4H3PO3 (t) → 3H3PO4 + PH3
9. Interaction of acid with non-metals (redox reaction). In this case, the non-metal is oxidized to the corresponding acid, and the acid is reduced to a gaseous oxide: H2SO4 (conc) - to SO2; HNO3(conc) - to NO2; HNO3(razb) - up to NO.
S + 2HNO3(dec) → H2SO4 + 2NO
S + 6HNO3(conc) → H2SO4 + 6NO2 + 2H2O
S + 2H2SO4(conc) → 3SO2 + CO2 + 2H2O
C + 2H2SO4(conc) → 2SO2 + CO2 + 2H2O
C + 4HNO3(conc) → 4NO2 + CO2 + 2H2O
P + 5HNO3(diff) + 2H2O → 3H3PO4 + 5NO
P + 5HNO3(conc) → HPO3 + 5NO2 + 2H2O
H2S + Г2 → 2НГ + S↓ (except F2)
H2SO3 + G2 + H2O → 2HG + H2SO4 (except F2)
2H2S(aq) + O2 → 2H2O + 2S↓
2H2S + 3O2 → 2H2O + 2SO2 (combustion)
2H2S + O2(deficient) → 2H2O + 2S↓
More active halogens displace less active NGs from acids (exception: F2 reacts with water, not acid)
2HBr + Cl2 → 2HCl + Br2↓
2HI + Cl2 → 2HCl + I2↓
2HI + Br2 → 2HBr + I2↓
10. Redox reactions between acids
H2SO4(conc) 2HBr → Br2↓ + SO2 + 2H2O
H2SO4(conc) + 8HI → 4I2↓ + H2S + 4H2O
H2SO4(conc) + HCl ≠
H2SO4(conc) + H2S → S↓ + SO2 + 2H2O
3H2SO4(conc) + H2S → 4SO2 + 4H2O
H2SO3 + 2H2S → 3S↓ + 3H2O
2HNO3(conc) + H2S → S↓ + 2NO2 + 2H2O
2HNO3(conc) + SO2 → H2SO4 + 2NO2
6HNO3(conc) + HI → HIO3 + 6NO2 + 3H2O
2HNO3(conc) + 6HCl → 3Cl2 + 2NO + 4H2O
Chemical properties of amphoteric hydroxides
1. Interaction with the basic oxide
amphoteric hydroxide + basic oxide → salt + water
2Al(OH)3 +Na2O (t)→ 2NaAlO2 + 3H2O
2. Interaction with amphoteric or acidic oxide
amphoteric hydroxide + amphoteric/acidic oxide ≠ no reaction
Some amphoteric oxides (Be (OH) 2, Zn (OH) 2, Pb (OH) 2) react with acidic CO2 oxide to form precipitation of basic salts and water
2Be(OH)2 + CO2 → (BeOH)2CO3↓ + H2O
3. Interaction with alkali
amphoteric hydroxide + alkali → salt + water
Zn(OH)2 + 2KOH(solid) (t) → K2ZnO2 + 2H2O
Zn(OH)2 + 2KOH → K2
4. Do not interact with insoluble bases or amphoteric hydroxides
amphoteric hydroxide + insoluble base/amphoteric hydroxide ≠ no reaction
5. Interaction with acids
amphoteric hydroxide + acid → salt + water
Al(OH)3 + 3HCl → AlCl3 + 3H2O
6. Do not react with salts
amphoteric hydroxide + salt ≠ no reaction
7. Do not react with metals / non-metals (simple substances)
amphoteric hydroxide + metal/non-metal ≠ no reaction
8. Thermal decomposition
amphoteric hydroxide (t) → amphoteric oxide + water
2Al(OH)3 (t) → Al2O3 + 3H2O
Zn(OH)2 (t) → ZnO + H2O
General information about salts
Imagine that we have an acid and a base, we will carry out a neutralization reaction between them and we will get an acid and a salt.
NaOH + HCl → NaCl (sodium chloride) + H2O
It turns out that the salt consists of a metal cation and an anion of an acid residue.
Salts are:
1. Acidic (with one or two hydrogen cations (that is, they have an acidic (or slightly acidic) environment) - KHCO3, NaHSO3).
2. Medium (I have a metal cation and an anion of an acid residue, the medium must be determined using a pH meter - BaSO4, AgNO3).
3. Basic (have a hydroxide ion, that is, an alkaline (or weakly alkaline) medium - Cu (OH) Cl, Ca (OH) Br).
There are also double salts that form cations of two metals (K) upon dissociation.
Salts, with few exceptions, are crystalline solids with high melting points. Most salts are white (KNO3, NaCl, BaSO4, etc.). Some salts are colored (K2Cr2O7 - orange color, K2CrO4 - yellow, NiSO4 - green, CoCl3 - pink, CuS - black). By solubility, they can be divided into soluble, slightly soluble and practically insoluble. Acid salts, as a rule, are better soluble in water than the corresponding medium salts, and basic salts are worse.
Chemical properties of salts
1. Salt + water
When many salts are dissolved in water, their partial or complete decomposition- hydrolysis. Some salts form crystalline hydrates. When dissolved in water, medium salts containing an amphoteric metal in the anion form complex salts.
NaCl + H2O → NaOH + HCl
Na2ZnO2 + 2H2O = Na2
2. Salt + Basic oxide ≠ no reaction
3. Salt + amphoteric oxide → (t) acid volatile oxide + salt
Amphoteric oxides displace volatile acidic oxides from their salts during fusion.
Al2O3 +K2CO3 → KAlO2 + CO2
Fe2O3 + Na2CO3 → 2NaFeO2 + CO2
4. Salt + acidic non-volatile oxide → acidic volatile oxide + salt
Non-volatile acid oxides displace volatile acid oxides from their salts during fusion.
SiO2 + CaCO3 → (t) CaSiO3 + CO2
P2O5 + Na2CO3 → (t) 2Na3PO4 + 3CO2
3SiO2 + Ca3(PO4)2 → (t) 3CaSiO3 + P2O5
5. Salt + base → base + salt
Reactions between salts and bases are ion exchange reactions. Therefore, under normal conditions, they proceed only in solutions (both the salt and the base must be soluble) and only on condition that a precipitate or a weak electrolyte (H2O / NH4OH) is formed as a result of the exchange; gaseous products are not formed in these reactions.
A. Soluble base + soluble acid salt → medium salt + water
If the salt and base are formed by different cations, then two middle salts are formed; in the case of acidic ammonium salts, an excess of alkali leads to the formation of ammonium hydroxide.
Ba(OH)2 + Ba(HCO3) → 2BaCO3 + 2H2O
2KOH + 2NaHCO3 → Na2CO3 + K2CO3 + 2H2O
2NaOH + 2NH4HS → Na2S + (NH4)2S + 2H2O
2NaOH(ex) + NH4Hs → Na2S + NH4OH + H2O
B. Soluble base + soluble medium/basic salt → insoluble salt↓ + base
Soluble base + soluble medium/basic salt → salt + insoluble base↓
Soluble base + soluble medium/basic salt → salt + weak electrolyte NH4OH
Soluble base + soluble medium/basic salt → no reaction
A reaction between soluble bases and a medium/basic salt occurs only if the exchange of ions produces an insoluble salt, or an insoluble base, or a weak electrolyte NH4OH.
Ba(OH)2 + Na2SO4 → BaSO4↓ + 2NaOH
2NH4OH + CuCl2 → 2NH4Cl + Cu(OH)2↓
Ba(OH)2 + NH4Cl → BaCl2 + NH4OH
NaOH + KCl ≠
If the original salt is formed by a polyacid base, with a lack of alkali, a basic salt is formed.
NaOH(deficient) + AlCl3 → Al(OH)Cl2 + NaCl
Under the action of alkalis on salts of silver and mercury (II), not AgOH and Hg (OH) 2 are released, which decompose at room temperature, but insoluble oxides Ag2O and HgO.
2AgNO3 + 2NaOH → Ag2O↓ 2NaNO3 + H2O
Hg(NO3)2 + 2KOH → HgO↓ + 2KNO3 + H2O
6. Salt + amphoteric hydroxide → no reaction
7. Salt + acid → acid + salt
Mostly. reactions of acids with salts are ion exchange reactions, therefore they proceed in solutions and only if an acid-insoluble salt or a weaker and volatile acid is formed.
HCl + AgNO3 → AgCl↓ + HNO3
2HBr + K2SiO3 → 2KBr +H2SiO3↓
2HNO3 + Na2CO3 → 2NaNO3 + H2O + CO2
A. Acid1 + salt of more volatile/weak acid2 → salt of acid1 + more volatile/ weak acid 2
Acids interact with solutions of salts of weaker or volatile acids. Regardless of the composition of the salt (medium, acidic, basic), as a rule, a medium salt and a weaker volatile acid are formed.
2CH3COOH + Na2S → 2CH3COONa + H2S
HCl + NaHS → NaCl + H2S
B. Strong acid + strong/medium acid salt → insoluble salt↓ + acid
Strong acids react with solutions of salts of other strong acids if an insoluble salt is formed. The non-volatile H3PO4 (medium strength acid) displaces the strong but volatile hydrochloric HCl and nitric acid HNO3 from their salts, provided an insoluble salt is formed.
H2SO4 + Ca(NO3)2 → CaSO4↓ + 2HNO3
2H3PO4 + 3CaCl2 → Ca3(PO4)2↓ + 6HCl
H3PO4 + 3AgNO3 → Ag3PO4↓ + 3HNO3
B. Acid1 + basic salt of acid1 → medium salt + water
When an acid reacts with a base salt of the same acid, a middle salt and water are formed.
HCl + Mg(OH)Cl → MgCl2 + H2O
D. Polybasic acid1 + medium/acid salt of acid1 → acid salt of acid1
When a polybasic acid acts on the average salt of the same acid, an acid salt is formed, and when an acid salt is acted upon, an acid salt containing a larger number of hydrogen atoms is formed.
H3PO4 + Ca3(PO4) → 3CaHPO4
H3PO4 + CaHPO4 → Ca(H2PO4)2
CO2 + H2O + CaCO3 → Ca(HCO3)2
E. Acid H2S + salt Ag, Cu, Pb, Cd, Hg → Ag2S/CuS/PbS/CdS/HgS↓ + acid
Weak and flighty hydrosulfide acid H2S displaces even strong acids from solutions of Ag, Cu, Pb, Cd, and Hg salts, forming sulfide precipitates with them, which are insoluble not only in water, but also in the resulting acid.
H2S + CuSO4 → CuS↓ + H2SO4
E. Acid + medium/complex salt with amphoteric Me in the anion → medium salt + amphoteric hydroxide↓
→ medium salt + medium salt + H2O
When an acid acts on an average or complex salt with an amphoteric metal in the anion, the salt is destroyed and formed:
a) in case of acid deficiency - middle salt and amphoteric hydroxide
b) in case of an excess of acid - two medium salts and water
2HCl(week) + Na2ZnO2 → 2NaCl + Zn(OH)2↓
2HCl(week) + Na2 → 2NaCl + Zn(OH)2↓ + 2H2O
4HCl(ex) + Na2ZnO2 → 2NaCl + ZnCl2 + 2H2O
4HCl(ex) + Na2 → 2NaCl + ZnCl2 + 4H2O
It should be borne in mind that in some cases, OVR or complex formation reactions occur between acids and salts. So, the OVR enter:
H2SO4 conc. and I‾/Br‾ (products H2S and I2/SO2 and Br2)
H2SO4 conc. and Fe²+ (SO2 and Fe³ products + )
HNO3 diluted/conc. and Fe² + (products NO/NO2 and Fe 3 + )
HNO3 diluted/conc. and SO3²‾/S²‾ (NO/NO2 products and sulphate/sulphur or sulphate)
HCl conc. and KMnO4/K2Cr2O7/KClO3 (products chlorine (gas) and Mn²+ /Cr³ + /Cl‾.
G. The reaction proceeds without a solvent
Sulfuric acid conc. + salt (tv.) → salt sour/medium + sour
Non-volatile acids can displace volatiles from their dry salts. Most often, the interaction of concentrated sulfuric acid with dry salts of strong and weak acids is used, in this case an acid and an acidic or medium salt are formed.
H2SO4(conc) + NaCl(solid) → NaHSO4 + HCl
H2SO4(conc) + 2NaCl(solid) → Na2SO4 + 2HCl
H2SO4(conc) + KNO3(s) → KHSO4 + HNO3
H2SO4(conc) + CaCO3(s) → CaSO4 + CO2 + H2O
8. Soluble salt + soluble salt → insoluble salt↓ + salt
Reactions between salts are exchange reactions. Therefore, under normal conditions, they proceed only if:
a) both salts are soluble in water and are taken as solutions
b) as a result of the reaction, a precipitate or a weak electrolyte is formed (the latter is very rare).
AgNO3 + NaCl → AgCl↓ + NaNO3
If one of the initial salts is insoluble, the reaction proceeds only when, as a result of it, an even more insoluble salt is formed. The criterion of "insolubility" is the value of PR (solubility product), however, since its study is beyond the scope of the school course, cases where one of the reagent salts is insoluble are not considered further.
If a salt is formed in the exchange reaction, which is completely decomposed as a result of hydrolysis (there are dashes in the solubility table in place of such salts), then the products of the hydrolysis of this salt become the products of the reaction.
Al2(SO4)3 + K2S ≠ Al2S3↓ + K2SO4
Al2(SO4)3 + K2S + 6H2O → 2Al(OH)3↓ + 3H2S + K2SO4
FeCl3 + 6KCN → K3 + 3KCl
AgI + 2KCN → K + KI
AgBr + 2Na2S2O3 → Na3 + NaBr
Fe2(SO4)3 + 2KI → 2FeSO4 + I2 + K2SO4
NaCl + NaHSO4 → (t) Na2SO4 + HCl
Middle salts sometimes interact with each other to form complex salts. OVR is possible between salts. Some salts interact when fused.
9. Salt of less active metal + more active metal → less active metal↓ + salt
A more active metal displaces a less active metal (to the right in the voltage series) from its salt solution, while a new salt is formed, and a less active metal is released in a free form (settles on a plate of active metal). The exception is alkaline and alkaline earth metals react with water in solution.
Salts with oxidizing properties enter into solution with metals and other redox reactions.
FeSO4 + Zn → Fe↓ + ZnSO4
ZnSO4 + Fe ≠
Hg(NO3)2 + Cu → Hg↓ + Cu(NO3)2
2FeCl3 + Fe → 3FeCl2
FeCl3 + Cu → FeCl2 + CuCl2
HgCl2 + Hg → Hg2Cl2
2CrCl3 + Zn → 2CrCl2 + ZnCl2
Metals can also displace each other from molten salts (the reaction is carried out without air access). In doing so, it must be remembered that:
a) when melted, many salts decompose
b) the voltage series of metals determines the relative activity of metals only in aqueous solutions (for example, Al in aqueous solutions is less active than alkaline earth metals, and in melts it is more active)
K + AlCl3(melt) →(t) 3KCl + Al
Mg + BeF2(melt) → (t) MgF2 + Be
2Al + 3CaCl2(melt) → (t) 2AlCl3 + 3Ca
10. Salt + non-metal
Reactions of salts with non-metals are few. These are redox reactions.
5KClO3 + 6P →(t) 5KCl + 3P2O5
2KClO3 + 3S →(t) 2KCl + 2SO2
2KClO3 + 3C →(t) 2KCl + 3CO2
More active halogens displace less active ones from solutions of salts of hydrohalic acids. An exception is molecular fluorine, which reacts in solutions not with salt, but with water.
2FeCl2 + Cl2 →(t) 2FeCl3
2NaNO2 + O2 → 2NaNO3
Na2SO3 + S →(t) Na2S2O3
BaSO4 + 2C →(t) BaS + 2CO2
2KClO3 + Br2 →(t) 2KBrO3 + Cl2 (the same reaction is typical for iodine)
2KI + Br2 → 2KBr + I2↓
2KBr + Cl2 → 2KCl + Br2↓
2NaI + Cl2 → 2NaCl + I2↓
11. Decomposition of salts.
Salt →(t) products thermal decomposition
1. Salts of nitric acid
Products of thermal decomposition of nitrates depend on the position of the metal cation in the series of metal stresses.
MeNO3 → (t) (for Me, to the left of Mg (excluding Li)) MeNO2 + O2
MeNO3 → (t) (for Me from Mg to Cu and also Li) MeO + NO2 + O2
MeNO3 → (t) (for Me, Cu is to the right) Me + NO2 + O2
(Thermal decomposition of iron(II)/chromium(II) nitrate produces iron(III)/chromium(III) oxide.
2. Ammonium salts
All ammonium salts decompose upon calcination. Most often, ammonia NH3 and acid or its decomposition products are released.
NH4Cl →(t) NH3 + HCl (=NH4Br, NH4I, (NH4)2S)
(NH4)3PO4 →(t)3NH3 + H3PO4
(NH4)2HPO4 →(t) 2NH3 + H3PO4
NH4H2PO4 →(t) NH3 + H3PO4
(NH4)2CO3 →(t) 2NH3 + CO2 + H2O
NH4HCO3 →(t) NH3 + CO2 + H2O
Sometimes ammonium salts containing anions - oxidizing agents decompose when heated with the release of N2, NO or N2O.
(NH4)Cr2O7 →(t) N2 + Cr2O3 + 4H2O
NH4NO3 →(t) N2O + 2H2O
2NH4NO3 →(t) N2 + 2NO + 4H2O
NH4NO2 →(t) N2 + 2H2O
2NH4MnO4 →(t) N2 + 2MnO2 + 4H2O
3. Salts of carbonic acid
Almost all carbonates decompose to metal oxide and CO2. Alkali metal carbonates, except for lithium, do not decompose when heated. Silver and mercury carbonates decompose to free metal.
MeCO3 →(t) MeO + CO2
2Ag2CO3 →(t) 4Ag + 2CO2 + O2
All bicarbonates decompose to the corresponding carbonate.
MeHCO3 →(t) MeCO3 + CO2 + H2O
4. Sulfuric acid salts
Sulfites disproportionate when heated, forming sulfide and sulfate. The sulfide (NH4)2S formed during the decomposition of (NH4)2SO3 immediately decomposes into NH3 and H2S.
MeSO3 →(t) MeS + MeSO4
(NH4)2SO3 →(t) 2NH3 + H2S + 3(NH4)2SO4
Hydrosulfites decompose to sulfites, SO2 and H2O.
MeHSO3 →(t) MeSO3 + SO2 +H2O
5. Salts of sulfuric acid
Many sulfates at t > 700-800 C decompose to metal oxide and SO3, which at this temperature decomposes to SO2 and O2. Alkali metal sulfates are heat resistant. Silver and mercury sulfates decompose to free metal. Hydrosulfates decompose first to disulfates and then to sulfates.
2CaSO4 →(t) 2CaO + 2SO2 + O2
2Fe2(SO4)3 →(t) 2Fe2O3 + 6SO2 + 3O2
2FeSO4 →(t) Fe2O3 + SO3 + SO2
Ag2SO4 →(t) 2Ag + SO2 + O2
MeHSO4 →(t) MeS2O7 + H2O
MeS2O7 →(t) MeSO4 + SO3
6. Complex salts
Hydroxocomplexes amphoteric metals decompose mainly into medium salt and water.
K →(t) KAlO2 + 2H2O
Na2 →(t) ZnO + 2NaOH + H2O
7. Basic salts
Many basic salts decompose when heated. Basic salts of anoxic acids decompose into water and oxosalts
Al(OH)2Br →(t) AlOBr + H2O
2AlOHCl2 →(t) Al2OCl4 + H2O
2MgOHCl →(t) Mg2OCl2 + H2O
Basic salts of oxygen-containing acids decompose into metal oxide and thermal decomposition products of the corresponding acid.
2AlOH(NO3)2 →(t) Al2O3 + NO2 + 3O2 + H2O
(CuOH)2CO3 →(t) 2CuO + H2O + CO2
8. Examples of thermal decomposition of other salts
4K2Cr2O7 →(t) 4K2CrO4 + 2Cr2O3 + 3O2
2KMnO4 →(t) K2MnO4 + MnO2 + O2
KClO4 →(t) KCl + O2
4KClO3 →(t) KCl + 3KClO4
2KClO3 →(t) 2KCl +3O2
2NaHS →(t) Na2S + H2S
2CaHPO4 →(t) Ca2P2O7 + H2O
Ca(H2PO4)2 →(t) Ca(PO3)2 +2H2O
2AgBr →(hν) 2Ag + Br2 (=AgI)
Most of the material presented is taken from the manual Deryabina N.E. "Chemistry. Main classes inorganic substances". IPO "At the Nikitsky Gate" Moscow 2011.
Chemistry preparation for ZNO and DPA
Comprehensive edition
PART AND
GENERAL CHEMISTRY
CHEMISTRY OF THE ELEMENTS
HALOGENS
Simple substances
Chemical properties of Fluorine
Fluorine is the strongest oxidizing agent in nature. Directly it does not react only with helium, neon and argon.
During the reaction with metals, fluorides are formed, ionic type compounds:
Fluorine reacts vigorously with many non-metals, even with some inert gases:
Chemical properties of chlorine. Interaction with complex substances
Chlorine is a stronger oxidizing agent than bromine or iodine, so chlorine displaces heavy halogens from their salts:
Dissolving in water, chlorine partially reacts with it, resulting in the formation of two acids: chloride and hypochlorite. In this case, one chlorine atom increases the degree of oxidation, and the other atom reduces it. Such reactions are called disproportionation reactions. Disproportionation reactions are self-healing-self-oxidation reactions, i.e. reactions in which one element exhibits the properties of both an oxide and a reducing agent. With disproportionation, compounds are simultaneously formed in which the element is in a more oxidized and reduced state compared to the primitive one. The oxidation state of the Chlorine atom in the hypochlorite acid molecule is +1:
The interaction of chlorine with alkali solutions proceeds similarly. In this case, two salts are formed: chloride and hypochlorite.
Chlorine interacts with various oxides:
Chlorine oxidizes some salts in which the metal is not in the maximum oxidation state:
Molecular chlorine reacts with many organic compounds. In the presence of ferrum(III) chloride as a catalyst, chlorine reacts with benzene to form chlorobenzene, and when irradiated with light, the same reaction produces hexachlorocyclohexane:
Chemical properties of bromine and iodine
Both substances react with hydrogen, fluorine and alkalis:
Iodine is oxidized by various strong oxidizing agents:
Mining methods simple substances
Extraction of fluorine
Since fluorine is the strongest chemical oxide, it is impossible to isolate it from compounds in a free form using chemical reactions, and therefore fluorine is mined by the physicochemical method - electrolysis.
To extract fluorine, potassium fluoride melt and nickel electrodes are used. Nickel is used due to the fact that the surface of the metal is passivated by fluorine due to the formation of insoluble
NiF2, therefore, the electrodes themselves are not destroyed by the action of the substance that is released on them:Extraction of chlorine
Chlorine is commercially produced by electrolysis of sodium chloride solution. As a result of this process, sodium hydroxide is also extracted:
In small quantities, chlorine is obtained by oxidizing a solution of hydrogen chloride by various methods:
Chlorine is a very important product of the chemical industry.
Its world production is millions of tons.
Extraction of bromine and iodine
For industrial use, bromine and iodine are obtained from the oxidation of bromides and iodides, respectively. For oxidation, molecular chlorine, concentrated sulfate acid or manganese dioxide are most often used:
Fluorine and some of its compounds are used as an oxidizing agent for rocket fuel. Large amounts of fluorine are used to produce various refrigerants (freons) and some polymers that are characterized by chemical and thermal resistance (Teflon and some others). Fluorine is used in nuclear technology to separate uranium isotopes.
Most of the chlorine is used to produce hydrochloric acid, and also as an oxidizing agent for the extraction of other halogens. In industry, it is used to bleach fabrics and paper. In larger quantities than fluorine, it is used for the production of polymers (PVC and others) and refrigerants. Disinfect with chlorine drinking water. It is also needed to extract some solvents such as chloroform, methylene chloride, carbon tetrachloride. And it is also used to produce many substances, such as potassium chlorate (bertolet salt), bleach and many other compounds containing chlorine atoms.
Bromine and iodine are not used in industry on the same scale as chlorine or fluorine, but the use of these substances is increasing every year. Bromine is used in the manufacture of various sedative medicines. Iodine is used in the manufacture of antiseptic preparations. Bromine and Iodine compounds are widely used in the quantitative analysis of substances. With the help of iodine, some metals are purified (this process is called iodine refining), such as titanium, vanadium and others.
Chemical properties of the main classes of inorganic compounds
Acid oxides
- Acid oxide + water \u003d acid (exception - SiO 2)
SO 3 + H 2 O \u003d H 2 SO 4
Cl 2 O 7 + H 2 O \u003d 2HClO 4 - Acid oxide + alkali \u003d salt + water
SO 2 + 2NaOH \u003d Na 2 SO 3 + H 2 O
P 2 O 5 + 6KOH \u003d 2K 3 PO 4 + 3H 2 O - Acid oxide + basic oxide = salt
CO 2 + BaO = BaCO 3
SiO 2 + K 2 O \u003d K 2 SiO 3Basic oxides
- Basic oxide + water \u003d alkali (oxides of alkali and alkaline earth metals react)
CaO + H 2 O \u003d Ca (OH) 2
Na 2 O + H 2 O \u003d 2NaOH - Basic oxide + acid = salt + water
CuO + 2HCl \u003d CuCl 2 + H 2 O
3K 2 O + 2H 3 PO 4 = 2K 3 PO 4 + 3H 2 O - Basic oxide + acid oxide = salt
MgO + CO 2 \u003d MgCO 3
Na 2 O + N 2 O 5 \u003d 2NaNO 3Amphoteric oxides
- Amphoteric oxide + acid = salt + water
Al 2 O 3 + 6HCl \u003d 2AlCl 3 + 3H 2 O
ZnO + H 2 SO 4 \u003d ZnSO 4 + H 2 O - Amphoteric oxide + alkali \u003d salt (+ water)
ZnO + 2KOH \u003d K 2 ZnO 2 + H 2 O (More correct: ZnO + 2KOH + H 2 O \u003d K 2)
Al 2 O 3 + 2NaOH = 2NaAlO 2 + H 2 O (More correct: Al 2 O 3 + 2NaOH + 3H 2 O = 2Na) - Amphoteric oxide + acid oxide = salt
ZnO + CO 2 = ZnCO 3 - Amphoteric oxide + basic oxide = salt (when fused)
ZnO + Na 2 O \u003d Na 2 ZnO 2
Al 2 O 3 + K 2 O \u003d 2KAlO 2
Cr 2 O 3 + CaO \u003d Ca (CrO 2) 2acids
- Acid + basic oxide = salt + water
2HNO 3 + CuO \u003d Cu (NO 3) 2 + H 2 O
3H 2 SO 4 + Fe 2 O 3 \u003d Fe 2 (SO 4) 3 + 3H 2 O - Acid + Amphoteric Oxide = Salt + Water
3H 2 SO 4 + Cr 2 O 3 \u003d Cr 2 (SO 4) 3 + 3H 2 O
2HBr + ZnO = ZnBr 2 + H 2 O - Acid + base = salt + water
H 2 SiO 3 + 2KOH \u003d K 2 SiO 3 + 2H 2 O
2HBr + Ni(OH) 2 = NiBr 2 + 2H 2 O - Acid + Amphoteric Hydroxide = Salt + Water
3HCl + Cr(OH) 3 = CrCl 3 + 3H 2 O
2HNO 3 + Zn(OH) 2 = Zn(NO 3) 2 + 2H 2 O - Strong acid + salt of a weak acid = weak acid + salt of a strong acid
2HBr + CaCO 3 \u003d CaBr 2 + H 2 O + CO 2
H 2 S + K 2 SiO 3 \u003d K 2 S + H 2 SiO 3 - Acid + metal (located to the left of hydrogen in the voltage series) \u003d salt + hydrogen
2HCl + Zn \u003d ZnCl 2 + H 2
H 2 SO 4 (razb.) + Fe \u003d FeSO 4 + H 2
Important: oxidizing acids (HNO 3 , conc. H 2 SO 4) react differently with metals.
Amphoteric hydroxides
- Amphoteric Hydroxide + Acid = Salt + Water
2Al(OH) 3 + 3H 2 SO 4 = Al 2 (SO 4) 3 + 6H 2 O
Be(OH) 2 + 2HCl = BeCl 2 + 2H 2 O - Amphoteric hydroxide + alkali \u003d salt + water (when fused)
Zn(OH) 2 + 2NaOH = Na 2 ZnO 2 + 2H 2 O
Al(OH) 3 + NaOH = NaAlO 2 + 2H 2 O - Amphoteric hydroxide + alkali = salt (in aqueous solution)
Zn(OH) 2 + 2NaOH \u003d Na 2
Sn (OH) 2 + 2NaOH \u003d Na 2
Be(OH) 2 + 2NaOH = Na 2
Al(OH) 3 + NaOH = Na
Cr(OH) 3 + 3NaOH = Na 3alkalis
- Alkali + acid oxide \u003d salt + water
Ba (OH) 2 + N 2 O 5 \u003d Ba (NO 3) 2 + H 2 O
2NaOH + CO 2 \u003d Na 2 CO 3 + H 2 O - Alkali + acid \u003d salt + water
3KOH + H 3 PO 4 = K 3 PO 4 + 3H 2 O
Ba(OH) 2 + 2HNO 3 = Ba(NO 3) 2 + 2H 2 O - Alkali + amphoteric oxide \u003d salt + water
2NaOH + ZnO = Na 2 ZnO 2 + H 2 O (More correct: 2NaOH + ZnO + H 2 O = Na 2) - Alkali + amphoteric hydroxide = salt (in aqueous solution)
2NaOH + Zn(OH) 2 = Na 2
NaOH + Al(OH) 3 = Na - Alkali + soluble salt = insoluble base + salt
Ca(OH) 2 + Cu(NO 3) 2 = Cu(OH) 2 + Ca(NO 3) 2
3KOH + FeCl 3 \u003d Fe (OH) 3 + 3KCl - Alkali + metal (Al, Zn) + water = salt + hydrogen
2NaOH + Zn + 2H 2 O \u003d Na 2 + H 2
2KOH + 2Al + 6H 2 O = 2K + 3H 2salt
- Salt of a Weak Acid + Strong Acid = Salt of a Strong Acid + Weak Acid
Na 2 SiO 3 + 2HNO 3 \u003d 2NaNO 3 + H 2 SiO 3
BaCO 3 + 2HCl \u003d BaCl 2 + H 2 O + CO 2 (H 2 CO 3) - Soluble salt + soluble salt = insoluble salt + salt
Pb(NO 3) 2 + K 2 S = PbS + 2KNO 3
CaCl 2 + Na 2 CO 3 \u003d CaCO 3 + 2NaCl - Soluble salt + alkali \u003d salt + insoluble base
Cu(NO 3) 2 + 2NaOH = 2NaNO 3 + Cu(OH) 2
2FeCl 3 + 3Ba(OH) 2 = 3BaCl 2 + 2Fe(OH) 3 - Soluble metal salt (*) + metal (**) = metal salt (**) + metal (*)
Zn + CuSO 4 \u003d ZnSO 4 + Cu
Cu + 2AgNO 3 \u003d Cu (NO 3) 2 + 2Ag
Important: 1) metal (**) must be in the voltage series to the left of metal (*), 2) metal (**) must NOT react with water.You may also be interested in other sections of the Chemistry Handbook:
- Salt of a Weak Acid + Strong Acid = Salt of a Strong Acid + Weak Acid
- Alkali + acid oxide \u003d salt + water
- Acid + basic oxide = salt + water
- Amphoteric oxide + acid = salt + water
- Basic oxide + water \u003d alkali (oxides of alkali and alkaline earth metals react)
elementary particles physical matter on our planet are atoms. In a free form, they can exist only at very high temperatures. Under normal conditions elementary particles tend to combine with each other using chemical bonds: ionic, metallic, covalent polar or non-polar. In this way, substances are formed, examples of which we will consider in our article.
Simple substances
The processes of interaction between atoms of the same chemical element end with the formation of chemicals called simple. So, coal is formed only by carbon atoms, hydrogen gas is formed by hydrogen atoms, and liquid mercury consists of mercury particles. The concept of a simple substance should not be identified with the concept of a chemical element. For example, carbon dioxide does not consist of simple substances of carbon and oxygen, but of the elements carbon and oxygen. Conventionally, compounds consisting of atoms of the same element can be divided into metals and non-metals. Consider some examples of the chemical properties of such simple substances.
Metals
Based on the position of the metallic element in the periodic system, the following groups can be distinguished: active metals, elements of the main subgroups of the third - eighth groups, metals of secondary subgroups of the fourth - seventh groups, as well as lanthanides and actinides. Metals are simple substances, examples of which we will give below, have the following general properties: thermal and electrical conductivity, metallic luster, plasticity and malleability. Such characteristics are inherent in iron, aluminum, copper and others. With an increase in the serial number in periods, the boiling and melting temperatures, as well as the hardness of metal elements, increase. This is due to the compression of their atoms, that is, a decrease in radius, as well as the accumulation of electrons. All parameters of metals are due internal structure crystal lattice of these compounds. Below we consider chemical reactions, and also give examples of the properties of substances related to metals.
Features of chemical reactions
All metals with an oxidation state of 0 exhibit only the properties of reducing agents. Alkaline and alkaline earth elements interact with water to form chemically aggressive bases - alkalis:
- 2Na+2H 2 0=2NaOH+H 2
A typical reaction of metals is oxidation. As a result of connection with oxygen atoms, substances of the class of oxides arise:
- Zn + O 2 \u003d ZnO
These are binary compounds related to complex substances. Examples of basic oxides are oxides of sodium Na 2 O, copper CuO, calcium CaO. They are capable of interacting with acids, as a result, salt and water are found in the products:
- MgO + 2HCl \u003d MgCl 2 + H 2 O
Substances of classes of acids, bases, salts are complex compounds and exhibit a variety of chemical properties. For example, between hydroxides and acids, a neutralization reaction occurs, leading to the appearance of salt and water. The composition of the salts will depend on the concentration of the reagents: for example, with an excess of acid in the reacting mixture, acid salts, for example, NaHCO 3 - sodium bicarbonate, and a high concentration of alkali causes the formation of basic salts, such as Al (OH) 2 Cl - aluminum dihydroxochloride.
non-metals
The most important non-metallic elements are found in the nitrogen, carbon, and halogen and chalcogen groups. periodic system. Let us give examples of substances related to non-metals: these are sulfur, oxygen, nitrogen, chlorine. All their physical features are opposite to the properties of metals. They don't spend electricity, poorly transmit heat rays, have low hardness. Interacting with oxygen, non-metals form complex compounds - acid oxides. The latter, reacting with acids, give acids:
- H 2 O + CO 2 → H 2 CO 3
A typical reaction characteristic of acidic oxides is the interaction with alkalis, leading to the appearance of salt and water.
The chemical activity of non-metals in the period increases, this is due to an increase in the ability of their atoms to attract electrons from others chemical elements. In groups, we observe the opposite phenomenon: non-metallic properties weaken due to the inflation of the volume of the atom due to the addition of new energy levels.
So, we examined the types of chemicals, examples illustrating their properties, position in the periodic system.
General properties of metals.
The presence of valence electrons weakly bound to the nucleus determines the general chemical properties of metals. AT chemical reactions they always act as a reducing agent, simple metal substances never exhibit oxidizing properties.
Getting metals:
- recovery from oxides with carbon (C), carbon monoxide(CO), hydrogen (H2) or more active metal(Al, Ca, Mg);
- recovery from salt solutions with a more active metal;
- electrolysis of solutions or melts of metal compounds - recovery of the most active metals (alkali, alkaline earth metals and aluminum) using electric current.
In nature, metals are found mainly in the form of compounds, only low-active metals are found in the form of simple substances (native metals).
Chemical properties of metals.
1. Interaction with simple substances non-metals:
Most metals can be oxidized with non-metals such as halogens, oxygen, sulfur, nitrogen. But most of these reactions require preheating to start. In the future, the reaction can proceed with the release of a large amount of heat, which leads to the ignition of the metal.
At room temperature, reactions are possible only between the most active metals (alkali and alkaline earth) and the most active non-metals (halogens, oxygen). Alkali metals (Na, K) react with oxygen to form peroxides and superoxides (Na2O2, KO2).
a) interaction of metals with water.
At room temperature, alkali and alkaline earth metals interact with water. As a result of the substitution reaction, an alkali (soluble base) and hydrogen are formed: Metal + H2O \u003d Me (OH) + H2
When heated, other metals interact with water, standing in the activity series to the left of hydrogen. Magnesium reacts with boiling water, aluminum - after a special surface treatment, resulting in the formation of insoluble bases - magnesium hydroxide or aluminum hydroxide - and hydrogen is released. Metals in the activity range from zinc (inclusive) to lead (inclusive) interact with water vapor (i.e. above 100 C), while oxides of the corresponding metals and hydrogen are formed.
Metals to the right of hydrogen in the activity series do not interact with water.
b) interaction with oxides:
active metals interact in a substitution reaction with oxides of other metals or non-metals, reducing them to simple substances.
c) interaction with acids:
Metals located to the left of hydrogen in the activity series react with acids to release hydrogen and form the corresponding salt. Metals to the right of hydrogen in the activity series do not interact with acid solutions.
A special place is occupied by the reactions of metals with nitric and concentrated sulfuric acids. All metals except noble ones (gold, platinum) can be oxidized by these oxidizing acids. As a result of these reactions, the corresponding salts will always be formed, water and the product of nitrogen or sulfur reduction, respectively.
d) with alkalis
Metals that form amphoteric compounds (aluminum, beryllium, zinc) are capable of reacting with melts (with the formation of medium salts of aluminates, beryllates or zincates) or alkali solutions (with the formation of the corresponding complex salts). All reactions will produce hydrogen.
e) In accordance with the position of the metal in the activity series, reactions of reduction (displacement) of a less active metal from a solution of its salt by another more active metal are possible. As a result of the reaction, a salt of a more active and simple substance is formed - a less active metal.
General properties of nonmetals.
There are much fewer non-metals than metals (22 elements). However, the chemistry of non-metals is much more complicated due to the greater filling of the external energy level of their atoms.
The physical properties of non-metals are more diverse: among them are gaseous (fluorine, chlorine, oxygen, nitrogen, hydrogen), liquids (bromine) and solids, which are very different from each other in melting point. Most non-metals do not conduct electricity, but silicon, graphite, germanium have semiconductor properties.
Gaseous, liquid and some solid non-metals (iodine) have a molecular structure crystal lattice, the remaining non-metals have an atomic crystal lattice.
Fluorine, chlorine, bromine, iodine, oxygen, nitrogen and hydrogen under normal conditions exist in the form of diatomic molecules.
Many non-metal elements form several allotropic modifications of simple substances. So oxygen has two allotropic modifications - oxygen O2 and ozone O3, sulfur has three allotropic modifications - rhombic, plastic and monoclinic sulfur, phosphorus has three allotropic modifications - red, white and black phosphorus, carbon - six allotropic modifications - soot, graphite, diamond , carbine, fullerene, graphene.
Unlike metals, which exhibit only reducing properties, non-metals in reactions with simple and complex substances can act both as a reducing agent and as an oxidizing agent. According to their activity, non-metals occupy a certain place in the series of electronegativity. Fluorine is considered the most active non-metal. He only shows oxidizing properties. Oxygen is in second place in terms of activity, nitrogen is in third, then halogens and other non-metals. Hydrogen has the lowest electronegativity among non-metals.
Chemical properties of non-metals.
1. Interaction with simple substances:
Nonmetals interact with metals. In such a reaction, metals act as a reducing agent, non-metals as an oxidizing agent. As a result of the reaction of the compound, binary compounds are formed - oxides, peroxides, nitrides, hydrides, salts of oxygen-free acids.
In the reactions of non-metals with each other, a more electronegative non-metal exhibits the properties of an oxidizing agent, a less electronegative one - the properties of a reducing agent. As a result of the compound reaction, binary compounds are formed. It must be remembered that non-metals can exhibit variable oxidation states in their compounds.
2. Interaction with complex substances:
a) with water:
Under normal conditions, only halogens interact with water.
b) with oxides of metals and non-metals:
Many non-metals can react at high temperatures with oxides of other non-metals, reducing them to simple substances. Non-metals to the left of sulfur in the electronegativity series can also interact with metal oxides, reducing metals to simple substances.
c) with acids:
Some non-metals can be oxidized with concentrated sulfuric or nitric acids.
d) with alkalis:
Under the action of alkalis, some non-metals can undergo dismutation, being both an oxidizing agent and a reducing agent.
For example, in the reaction of halogens with alkali solutions without heating: Cl2 + 2NaOH = NaCl + NaClO + H2O or when heated: 3Cl2 + 6NaOH = 5NaCl + NaClO3 + 3H2O.
e) with salts:
When interacting, being strong oxidizing agents, they exhibit reducing properties.
Halogens (except fluorine) enter into substitution reactions with solutions of salts of hydrohalic acids: a more active halogen displaces a less active halogen from a salt solution.
