Since d-metal oxides are insoluble in water, their hydroxides are obtained indirectly using exchange reactions between their salts and alkali solutions:

ZnCl 2 + 2NaOH = Zn(OH) 2 + 2NaCl;

MnCl 2 + 2NaOH = Mn(OH) 2 + 2NaCl (in the absence of oxygen);

FeSO 4 + 2KOH = Fe(OH) 2 + K 2 SO 4 (in the absence of oxygen).

Hydroxides of d-elements in lower oxidation states are weak bases; They are insoluble in water, but dissolve well in acids:

Cu(OH) 2 + 2HCl = CuCl 2 + H 2 O

Cu(OH) 2 + H 2 SO 4 = CuSO 4 + H 2 O

Hydroxides of d-elements in intermediate oxidation states and zinc hydroxide dissolve not only in acids, but also in excess alkali solutions with the formation of hydroxo complexes (i.e., they exhibit amphoteric properties), for example:

Zn(OH) 2 + H 2 SO 4 = ZnSO 4 + 2H 2 O;

Zn(OH) 2 + 2NaOH = Na 2;

Cr(OH) 3 + 3HNO 3 = Cr(NO 3) 3 + 3H 2 O;

Cr(OH) 3 + 3KOH = K 3.

In higher oxidation states, transition metals form hydroxides, which exhibit acidic properties or amphoteric properties with a predominance of acidic ones:

With an increase in the degree of oxidation of an element, the basic properties of oxides and hydroxides weaken, and the acidic properties increase.

Therefore, across the period from left to right, there is an increase in the acidic properties of d-metal hydroxides in higher oxidation states up to the Mn subgroup, then the acidic properties weaken:

Sc(OH) 3 - TiO 2 xH 2 O - V 2 O 5 xH 2 O - H 2 CrO 4 - HMnO 4

Strengthening acid properties

Fe(OH) 3 - Co(OH) 2 - Cu(OH) 2 - Zn(OH) 2

Slow weakening of acid properties

Let us consider the change in the properties of d-metal hydroxides in subgroups. From top to bottom in the subgroup, the basic properties of hydroxides of d-elements in higher oxidation states increase, while the acidic properties decrease. For example, for the sixth group of d-metals:

H 2 CrO 4 - sharp - MoO 3 H 2 O - weak - WO 3 H 2 O

Acid properties are reduced

Redox properties of d-element compounds

Connections of d-elements in lower oxidation states they exhibit, mostly, reducing properties, especially in an alkaline environment. Therefore, for example, hydroxides Mn(+2), Cr(+2), Fe(+2) are very unstable and are quickly oxidized by atmospheric oxygen:

2Mn(OH)2 + O2 + 2H2O = 2Mn(OH)4;

4Cr(OH) 2 + O 2 + 2H 2 O = 4Cr(OH) 3

In order to convert cobalt (II) or nickel (II) hydroxide into Co(OH) 3 or Ni(OH) 3, it is necessary to use a stronger oxidizing agent - for example, hydrogen peroxide H 2 O 2 in an alkaline medium or bromine Br 2:

2Co(OH) 2 + H 2 O 2 = 2Co(OH) 3;

2 Ni(OH) 2 + Br 2 +2NaOH = 2 Ni(OH) 3 + 2NaBr

Derivatives of Ti(III), V(III), V(II), Cr (II) are easily oxidized in air, some salts can be oxidized even with water:

2Ti 2 (SO 4) 3 + O 2 + 2H 2 O = 4TiOSO 4 + 2H 2 SO 4;

2CrCl 2 + 2H 2 O = 2Cr(OH) Cl 2 + H 2

Compounds of d-elements in higher oxidation states (from +4 to +7) usually exhibit oxidizing properties. However, Ti(IV) and V(V) compounds are always stable and therefore have relatively weak oxidizing properties:

TiOSO 4 + Zn + H 2 SO 4 = Ti 2 (SO 4) 3 + ZnSO 4 + H 2 O;

Na 3 VO 4 + Zn + H 2 SO 4 = VOSO 4 + ZnSO 4 + H 2 O

Reduction occurs under harsh conditions - with atomic hydrogen at the moment of its release (Zn + 2H + = 2H + Zn 2+).

And chromium compounds in higher oxidation states are strong oxidizing agents, especially in an acidic environment:

K2Cr2O7 + 3SO2 + H2SO4 = Cr2(SO4)3 + K2SO4 + H2O;

2CrO 3 + C 2 H 5 OH = Cr 2 O 3 + CH 3 COH + H 2 O

Mn(VI), Mn(VII) and Fe(VI) compounds exhibit even stronger oxidizing properties:

2KMnO 4 + 6KI + 4H 2 O = 2MnO 2 + 3I 2 + 8KOH;

4K 2 FeO 4 + 10H 2 SO 4 = 2Fe 2 (SO 4) 3 + 3O 2 +10H 2 O+ 4K 2 SO 4

Thus, the oxidizing properties of compounds of d-elements in higher oxidation states increase across the period from left to right.

The oxidizing ability of compounds of d-elements in higher oxidation states in the subgroup from top to bottom weakens. For example, in the chromium subgroup: potassium bichromate K 2 Cr 2 O 7 interacts even with such a weak reducing agent as SO 2 . To reduce molybdate or tungstate ions, a very strong reducing agent is required, for example, a hydrochloric acid solution of tin (II) chloride:

K 2 Cr 2 O 7 + SO 2 + H 2 SO 4 = Cr 2 (SO 4) 3 + K 2 SO 4 + H 2 O

3 (NH 4) 2 MoO 4 + HSnCl 3 + 9HCl = MoO 3 MoO 5 + H 2 SnCl 6 + 4H 2 O + 6NH 4 Cl

The last reaction occurs when heated, and the oxidation state of the d-element decreases very slightly.

Compounds of d-metals in intermediate oxidation states exhibit redox duality. For example, iron (III) compounds, depending on the nature of the partner substance, can exhibit reducing agent properties:

2FeCl3 + Br2 + 16KOH = 2K2FeO4 + 6KBr + 6KCl +8H2O,

and oxidizing properties:

2FeCl 3 + 2KI = 2FeCl 2 + I 2 +2KCl.

Bases, amphoteric hydroxides

Bases are complex substances consisting of metal atoms and one or more hydroxyl groups (-OH). The general formula is Me +y (OH) y, where y is the number of hydroxo groups equal to the oxidation state of the metal Me. The table shows the classification of bases.

Properties of alkalis, hydroxides of alkali and alkaline earth metals

1. Aqueous solutions of alkalis are soapy to the touch, change the color of indicators: litmus - in Blue colour, phenolphthalein - into crimson.

2. Aqueous solutions dissociate:

3. Interact with acids, entering into an exchange reaction:

Polyacid bases can give medium and basic salts:

4. React with acidic oxides, forming medium and acidic salts depending on the basicity of the acid corresponding to this oxide:

5. Interact with amphoteric oxides and hydroxides:

a) fusion:

b) in solutions:

6. Interact with water-soluble salts if a precipitate or gas is formed:

Insoluble bases (Cr(OH) 2, Mn(OH) 2, etc.) interact with acids and decompose when heated:

Amphoteric hydroxides

Amphoteric compounds are compounds that, depending on conditions, can be both donors of hydrogen cations and exhibit acidic properties, and their acceptors, i.e., exhibit basic properties.

Chemical properties of amphoteric compounds

1. Interacting with strong acids, they exhibit basic properties:

Zn(OH) 2 + 2HCl = ZnCl 2 + 2H 2 O

2. Interacting with alkalis - strong bases, they exhibit acidic properties:

Zn(OH) 2 + 2NaOH = Na 2 ( complex salt)

Al(OH) 3 + NaOH = Na ( complex salt)

Complex compounds are those in which at least one covalent bond is formed by a donor-acceptor mechanism.

The general method for preparing bases is based on exchange reactions, with the help of which both insoluble and soluble bases can be obtained.

CuSO 4 + 2KOH = Cu(OH) 2 ↓ + K 2 SO 4

K 2 CO 3 + Ba(OH) 2 = 2 KOH + BaCO 3 ↓

When soluble bases are obtained by this method, an insoluble salt precipitates.

When preparing water-insoluble bases with amphoteric properties, excess alkali should be avoided, since dissolution of the amphoteric base may occur, for example:

AlCl 3 + 4KOH = K[Al(OH) 4 ] + 3KCl

In such cases, ammonium hydroxide is used to obtain hydroxides, in which amphoteric hydroxides do not dissolve:

AlCl 3 + 3NH 3 + ZH 2 O = Al(OH) 3 ↓ + 3NH 4 Cl

Silver and mercury hydroxides decompose so easily that when trying to obtain them by exchange reaction, instead of hydroxides, oxides precipitate:

2AgNO 3 + 2KOH = Ag 2 O↓ + H 2 O + 2KNO 3

In industry, alkalis are usually obtained by electrolysis of aqueous solutions of chlorides.

2NaCl + 2H 2 O → ϟ → 2NaOH + H 2 + Cl 2

Alkalis can also be obtained by reacting alkali and alkaline earth metals or their oxides with water.

2Li + 2H 2 O = 2LiOH + H 2

SrO + H 2 O = Sr(OH) 2

Acids

Acids are complex substances whose molecules consist of hydrogen atoms that can be replaced by metal atoms and acidic residues. Under normal conditions, acids can be solid (phosphoric H 3 PO 4; silicon H 2 SiO 3) and liquid (in pure form the liquid will be sulfuric acid H 2 SO 4).

Gases such as hydrogen chloride HCl, hydrogen bromide HBr, hydrogen sulfide H 2 S form the corresponding acids in aqueous solutions. The number of hydrogen ions formed by each acid molecule during dissociation determines the charge of the acid residue (anion) and the basicity of the acid.

According to protolytic theory of acids and bases, proposed simultaneously by the Danish chemist Brønsted and the English chemist Lowry, an acid is a substance splitting off with this reaction protons, A basis- a substance that can accept protons.

acid → base + H +

Based on such ideas, it is clear basic properties of ammonia, which, due to the presence of a lone electron pair at the nitrogen atom, effectively accepts a proton when interacting with acids, forming an ammonium ion through a donor-acceptor bond.

HNO 3 + NH 3 ⇆ NH 4 + + NO 3 —

acid base acid base

More general definition of acids and bases proposed by the American chemist G. Lewis. He suggested that acid-base interactions are completely do not necessarily occur with the transfer of protones. In the Lewis determination of acids and bases, the main role in chemical reactions is played by electron pairs

Cations, anions, or neutral molecules that can accept one or more pairs of electrons are called Lewis acids.

For example, aluminum fluoride AlF 3 is an acid, since it is able to accept an electron pair when interacting with ammonia.

AlF 3 + :NH 3 ⇆ :

Cations, anions, or neutral molecules capable of donating electron pairs are called Lewis bases (ammonia is a base).

Lewis's definition covers all acid-base processes that were considered by previously proposed theories. The table compares the definitions of acids and bases currently used.

Nomenclature of acids

Since there are different definitions of acids, their classification and nomenclature are rather arbitrary.

According to the number of hydrogen atoms capable of elimination in an aqueous solution, acids are divided into monobasic(e.g. HF, HNO 2), dibasic(H 2 CO 3, H 2 SO 4) and tribasic(H 3 PO 4).

According to the composition of the acid, they are divided into oxygen-free(HCl, H 2 S) and oxygen-containing(HClO 4, HNO 3).

Usually names of oxygen-containing acids are derived from the name of the non-metal with the addition of the endings -kai, -vaya, if the oxidation state of the non-metal is equal to the group number. As the oxidation state decreases, the suffixes change (in order of decreasing oxidation state of the metal): -opaque, rusty, -ovish:

If we consider the polarity of the hydrogen-nonmetal bond within a period, we can easily relate the polarity of this bond to the position of the element in the Periodic Table. From metal atoms, which easily lose valence electrons, hydrogen atoms accept these electrons, forming a stable two-electron shell like the shell of a helium atom, and give ionic metal hydrides.

In hydrogen compounds of elements of groups III-IV of the Periodic Table, boron, aluminum, carbon, and silicon form covalent, weakly polar bonds with hydrogen atoms that are not prone to dissociation. For elements of groups V-VII of the Periodic Table, within a period, the polarity of the nonmetal-hydrogen bond increases with the charge of the atom, but the distribution of charges in the resulting dipole is different than in hydrogen compounds of elements that tend to donate electrons. Non-metal atoms, which require several electrons to complete the electron shell, attract (polarize) a pair of bonding electrons the more strongly, the greater the nuclear charge. Therefore, in the series CH 4 - NH 3 - H 2 O - HF or SiH 4 - PH 3 - H 2 S - HCl, bonds with hydrogen atoms, while remaining covalent, become more polar in nature, and the hydrogen atom in the element-hydrogen bond dipole becomes more electropositive. If polar molecules find themselves in a polar solvent, a process of electrolytic dissociation can occur.

Let us discuss the behavior of oxygen-containing acids in aqueous solutions. These acids have N-O-E connection and, naturally, the polarity of the H-O bond is affected by O-E connection. Therefore, these acids, as a rule, dissociate more easily than water.

H 2 SO 3 + H 2 O ⇆ H 3 O + + HSO 3

HNO 3 + H 2 O ⇆ H 3 O + + NO 3

Let's look at a few examples properties of oxygen-containing acids, formed by elements that are capable of exhibiting different degrees of oxidation. It is known that hypochlorous acid HClO very weak chlorous acid HClO 2 also weak, but stronger than hypochlorous, hypochlorous acid HClO 3 strong. Perchloric acid HClO 4 is one of the strongest inorganic acids.

For acid-type dissociation (with elimination of the H ion), a rupture is required O-N connections. How can we explain the decrease in the strength of this bond in the series HClO - HClO 2 - HClO 3 - HClO 4? In this series, the number of oxygen atoms associated with the central chlorine atom increases. Each time a new oxygen-chlorine bond is formed, electron density is drawn from the chlorine atom, and therefore from the O-Cl single bond. As a result, the electron density partially leaves the O-H bond, which is weakened as a result.

This pattern - strengthening of acidic properties with increasing degree of oxidation of the central atom - characteristic not only of chlorine, but also of other elements. For example, nitric acid HNO 3, in which the oxidation state of nitrogen is +5, is stronger than nitrous acid HNO 2 (the oxidation state of nitrogen is +3); sulfuric acid H 2 SO 4 (S +6) is stronger than sulfurous acid H 2 SO 3 (S +4).

Obtaining acids

1. Oxygen-free acids can be obtained by direct combination of non-metals with hydrogen.

H 2 + Cl 2 → 2HCl,

H 2 + S ⇆ H 2 S

2. Some oxygen-containing acids can be obtained interaction of acid oxides with water.

3. Both oxygen-free and oxygen-containing acids can be obtained by metabolic reactions between salts and other acids.

BaBr 2 + H 2 SO 4 = BaSO 4 ↓ + 2НВr

CuSO 4 + H 2 S = H 2 SO 4 + CuS↓

FeS + H 2 SO 4 (pa zb) = H 2 S + FeSO 4

NaCl (T) + H 2 SO 4 (conc) = HCl + NaHSO 4

AgNO 3 + HCl = AgCl↓ + HNO 3

CaCO 3 + 2HBr = CaBr 2 + CO 2 + H 2 O

4. Some acids can be obtained using redox reactions.

H 2 O 2 + SO 2 = H 2 SO 4

3P + 5HNO 3 + 2H 2 O = ZN 3 PO 4 + 5NO 2

Sour taste, effect on indicators, electrical conductivity, interaction with metals, basic and amphoteric oxides, bases and salts, formation of esters with alcohols - these properties are common to inorganic and organic acids.

can be divided into two types of reactions:

1) are common For acids reactions are associated with the formation of hydronium ion H 3 O + in aqueous solutions;

2) specific(i.e. characteristic) reactions specific acids.

The hydrogen ion can enter into redox reaction, reducing to hydrogen, as well as in a compound reaction with negatively charged or neutral particles having lone pairs of electrons, i.e. in acid-base reactions.

The general properties of acids include reactions of acids with metals in the voltage series up to hydrogen, for example:

Zn + 2Н + = Zn 2+ + Н 2

Acid-base reactions include reactions with basic oxides and bases, as well as with intermediate, basic, and sometimes acidic salts.

2 CO 3 + 4HBr = 2CuBr 2 + CO 2 + 3H 2 O

Mg(HCO 3) 2 + 2HCl = MgCl 2 + 2CO 2 + 2H 2 O

2KHSO 3 + H 2 SO 4 = K 2 SO 4 + 2SO 2 + 2H 2 O

Note that polybasic acids dissociate stepwise, and at each subsequent step the dissociation is more difficult, therefore, with an excess of acid, acidic salts are most often formed, rather than average ones.

Ca 3 (PO 4) 2 + 4H 3 PO 4 = 3Ca (H 2 PO 4) 2

Na 2 S + H 3 PO 4 = Na 2 HPO 4 + H 2 S

NaOH + H 3 PO 4 = NaH 2 PO 4 + H 2 O

KOH + H 2 S = KHS + H 2 O

At first glance, the formation of acid salts may seem surprising monobasic hydrofluoric acid. However, this fact can be explained. Unlike all other hydrohalic acids, hydrofluoric acid in solutions is partially polymerized (due to the formation of hydrogen bonds) and various particles (HF) X may be present in it, namely H 2 F 2, H 3 F 3, etc.

A special case of acid-base equilibrium - reactions of acids and bases with indicators that change their color depending on the acidity of the solution. Indicators are used in qualitative analysis to detect acids and bases in solutions.

The most commonly used indicators are litmus(V neutral environment purple, V sour - red, V alkaline - blue), methyl orange(V sour environment red, V neutral - orange, V alkaline - yellow), phenolphthalein(V highly alkaline environment raspberry red, V neutral and acidic - colorless).

Specific properties different acids can be of two types: firstly, reactions leading to the formation insoluble salts, and secondly, redox transformations. If the reactions associated with the presence of the H + ion are common to all acids (qualitative reactions for detecting acids), specific reactions are used as qualitative reactions for individual acids:

Ag + + Cl - = AgCl (white precipitate)

Ba 2+ + SO 4 2- = BaSO 4 (white precipitate)

3Ag + + PO 4 3 - = Ag 3 PO 4 (yellow precipitate)

Some specific reactions of acids are due to their redox properties.

Anoxic acids in an aqueous solution can only be oxidized.

2KMnO 4 + 16HCl = 5Сl 2 + 2КСl + 2МnСl 2 + 8Н 2 O

H 2 S + Br 2 = S + 2НВг

Oxygen-containing acids can be oxidized only if the central atom in them is in a lower or intermediate oxidation state, as, for example, in sulfurous acid:

H 2 SO 3 + Cl 2 + H 2 O = H 2 SO 4 + 2HCl

Many oxygen-containing acids, in which the central atom has the maximum oxidation state (S +6, N +5, Cr +6), exhibit the properties of strong oxidizing agents. Concentrated H 2 SO 4 is a strong oxidizing agent.

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

Pb + 4HNO 3 = Pb(NO 3) 2 + 2NO 2 + 2H 2 O

C + 2H 2 SO 4 (conc) = CO 2 + 2SO 2 + 2H 2 O

It should be remembered that:

Acid solutions react with metals that are to the left of hydrogen in the electrochemical voltage series, subject to a number of conditions, the most important of which is the formation of a soluble salt as a result of the reaction. The interaction of HNO 3 and H 2 SO 4 (conc.) with metals proceeds differently.

Concentrated sulfuric acid in the cold passivates aluminum, iron, and chromium.

In water, acids dissociate into hydrogen cations and anions of acid residues, for example:

Inorganic and organic acids react with basic and amphoteric oxides, provided that a soluble salt is formed:

Both acids react with bases. Polybasic acids can form both intermediate and acid salts (these are neutralization reactions):

The reaction between acids and salts occurs only if a precipitate or gas is formed:

The interaction of H 3 PO 4 with limestone will stop due to the formation of the last insoluble precipitate of Ca 3 (PO 4) 2 on the surface.

The peculiarities of the properties of nitric HNO 3 and concentrated sulfuric H 2 SO 4 (conc.) acids are due to the fact that when they interact with simple substances (metals and non-metals), the oxidizing agents will not be H + cations, but nitrate and sulfate ions. It is logical to expect that as a result of such reactions, not hydrogen H2 is formed, but other substances are obtained: necessarily salt and water, as well as one of the products of the reduction of nitrate or sulfate ions, depending on the concentration of acids, the position of the metal in the voltage series and reaction conditions (temperature, degree of metal grinding, etc.).

These features of the chemical behavior of HNO 3 and H 2 SO 4 (conc.) clearly illustrate the thesis of the theory of chemical structure about the mutual influence of atoms in the molecules of substances.

The concepts of volatility and stability (stability) are often confused. Volatile acids are acids whose molecules easily pass into a gaseous state, that is, evaporate. For example, hydrochloric acid is a volatile but stable acid. It is impossible to judge the volatility of unstable acids. For example, non-volatile, insoluble silicic acid decomposes into water and SiO 2. Aqueous solutions of hydrochloric, nitric, sulfuric, phosphoric and a number of other acids are colorless. An aqueous solution of chromic acid H 2 CrO 4 is yellow in color, and manganese acid HMnO 4 is crimson.

Reference material for taking the test:

Mendeleev table

Solubility table

Bronze: Cu (70-96%), Sn (everything else).

Constantan: Cu (55%), Ni (44%).

Brass: Cu (54-90%), Zn (everything else).

Neusilver: Cu (50-65%), Ni (8-26%), Zn (everything else).

Application:

Bronze – manufacturing of machine parts. Constantan is an electrical resistance material.

Brass – production of wires, sheets, profiles, fittings. Nickel silver is a material for precision mechanics and metallurgical instruments.

Question No. 21

What factors determine the properties of metal oxides and hydroxides? Explain with specific examples.

The properties of metal oxides and hydroxides depend on the degree of oxidation of the metal. The higher the oxidation state of a metal, the more pronounced its acidic properties. This is clearly seen in the example of chromium oxides.

Chromium(II) oxide and chromium(II) hydroxide exhibit basic properties. When reacting with acids, they form salts.

CrO + 2HCl = CrCl2 + H2O

Cr(OH)2 + 2HCl = CrCl2 + 2H2O

Chromium(III) oxide and hydroxide are amphoteric, they react with both acids and bases:

Сr2 О3 + 6HCl = 2СrСl3 + 3Н2 О Сr2 О3 + 2NaOH + 3Н2 О = 2Na

Cr(OH)3 + 3HCl = CrCl3 + 3H2 O Cr(OH)3 + NaOH = Na

Chromium (VI) oxide is an acidic oxide; when reacted with water, it forms chromic acid H2 CrO4:

CrO3 + H2 O = H2 CrO4

When chromic acid or chromium (VI) oxide reacts with bases, salts are formed - chromates:

CrO2 + 2NaOH = Na2 CrO4 + H2 O

Н2 СrО4, + 2NaOH = Na2 CrO4 + 2Н2 O

Task No. 1

What mass of pure iron can be obtained from 250 tons of ore with a mass fraction of pyrite FeS2 of 0.7, if the yield is 82%?

Iron(III) oxide is first obtained from pyrite:

FeS2 + О2 Fe2 О3 + SO2

To set the odds, we will use the electronic balance method:

2 − 1			3 − 2	4 − 2
FeS2		→ Fe2 O3
−1	− 10e− →
	− 10e− →

		− →
Fe−e		− →
	4e − →		−2
	4e − →

4FeS2 + 11O2 = 2Fe2 O3 + 8SO2

From iron (III) oxide, iron can be obtained using any suitable reducing agent, for example carbon (II) oxide:

Fe2 O3 + 3CO = 2Fe + 3CO2 (2)

Let's calculate the mass of pure pyrite in the ore:

M(FeS2) = w(FeS2) m(ore) = 0.7 250 t = 175 t.

Let's calculate the molar mass of pyrite:

M(FeS2) = 56 + 32 2 = 120 g/mol

Let's calculate the amount of pyrite substance:

According to equation (1), from 4 moles of pyrite you will get 2 moles of iron oxide. According to equation (2), from 1 mole of iron oxide, 2 moles of iron are obtained. In total, this means that from 4 moles of pyrite, 4 moles of iron are obtained. Consequently, from 1.46·106 mol of pyrite with a theoretical 100% yield, 1.46·106 mol of iron can be obtained. Since the iron yield is 82%, or 0.82, it is practically possible to obtain 0.82 1.46 106 ≈ 1.2 106 mol. The molar mass of iron is 56 g/mol, let's calculate the mass of iron:

m(Fe) = ν (Fe) M(Fe) = 1.2 106 mol 56 g/mol = 67.2 106 = 67.2 t.

Answer: you can get 67.2 tons of iron.

Problem No. 2

During the electrolysis of a sodium chloride solution, 7.2 liters of hydrogen (n.e.) were released. Calculate the mass and amount of substance formed in sodium hydroxide in solution.

Let us write down the equations of the processes occurring on the electrodes:

2H2 O + 2e− → 2OH− + H2
2Cl− − 2e− → Cl2

2NaCl + 2H2 O = 2NaOH + Cl2 + H2

Thus, sodium hydroxide is formed in the solution, hydrogen is released at the cathode, and chlorine is released at the anode.

According to the reaction equation, for 1 mole of hydrogen released at the cathode, there are 2 moles of sodium hydroxide formed in the solution. Let the release of 0.32 mol of hydrogen in a solution produce x mol of sodium hydroxide. Let's make a proportion:

1 2 = 0, x 32, x = 0.32 1 2 = 0.64 mol

Let's determine the molar mass of sodium hydroxide:

m(NaOH) = ν (NaOH) M(NaOH) = 0.64 mol 40 g/mol = 25.6 g.

Answer: 0.64 mol (25.6 g) of sodium hydroxide was formed in the solution.

Problem No. 3

23.2 liters of hydrogen sulfide were passed through 1 liter of 18% copper (II) sulfate solution (ρ = 1.12 g/cm3). What substance and how much by mass precipitated?

Copper (II) sulfide precipitates:

CuSO4 + H2 S = CuS↓ + H2 SO4

Let's calculate the amount of hydrogen sulfide:

Let's calculate the mass of copper sulfate solution: m(solution) = ρ · V = 1.12 g/ml · 1000 ml = 1120 g.

Let's calculate the mass of copper sulfate in solution:

m(CuSO4) = c(CuSO4) m(solution) = 0.18 1120 g = 201.6 g

Let's determine the molar mass of copper sulfate:

M(CuSO4) = 64 + 32 + 16 4 = 160 g/mol

Let's calculate the amount of copper sulfate:

According to the reaction equation, 1 mole of hydrogen sulfide reacts with 1 mole of copper sulfate, which means that 1.036 moles of copper sulfate will react with 1.036 moles of hydrogen sulfide, that is, copper sulfate is taken in excess and the calculation is carried out using hydrogen sulfide. According to the reaction equation, 1 mole of copper (II) sulfide is formed from 1 mole of hydrogen sulfide, which means that 1.036 moles of copper (II) sulfide are formed from 1.036 moles of hydrogen sulfide. Let's calculate the molar mass of copper (II) sulfide:

M(CuS) = 64 + 32 = 96 g/mol.

Let's calculate the mass of copper (II) sulfide:

m(CuS) = ν (CuS) M(CuS) = 1.036 mol 96 g/mol ≈ 120.96 g.

121 g of copper (II) sulfide precipitate.

Problem No. 4

When 9 g of a mixture consisting of metallic aluminum and its oxide was exposed to a 40% solution of sodium hydroxide (ρ = 1.4 g/cm3), 3.36 liters of gas (n.e.) were released. Determine the percentage composition of the initial mixture and the volume of NaOH solution that reacted.

Reaction equations:

2Al + 2NaOH + 6H2 O = 2Na + 3H2 (1)

Al2 O3 + 2NaOH + 3H2 O = 2Na (2)

Let's calculate the amount of hydrogen released:

According to the reaction equation, when 2 moles of aluminum react with an alkali solution, 3 moles of hydrogen are released. Let 0.15 mol of hydrogen be released when x mol of aluminum reacts with an alkali solution. Let's make a proportion:

2 3 = 0, x 15, x = 0.15 3 2 = 0.1 mol

The molar mass of aluminum is 27 g/mol, let's calculate the mass of aluminum:

m(Al) = ν (Al) M(Al) = 0.1 mol 27 g/mol = 2.7 g

Let's calculate the mass fraction of aluminum in the mixture:

w(Al) =	m(Al)	100% =	100% = 30%
	m(mixtures)

Let's calculate the mass fraction of aluminum oxide in the mixture: w(Al2 O3) = 100% – w(Al) = 70%.

According to equation (1), 2 mol of aluminum reacts with 2 mol of sodium hydroxide, which means 0.1 mol of aluminum reacts with 0.1 mol of hydroxide

sodium oxide. The mixture contains 9 – 2.7 = 6.3 g of aluminum oxide. Let's calculate the molar mass of aluminum oxide:

M(Al2 O3) = 27 2 + 16 3 = 102 g/mol.

Let's calculate the amount of aluminum oxide substance:

According to reaction equation (2), 1 mol of aluminum oxide reacts with 2 mol of aluminum hydroxide. Let 0.062 mol of aluminum oxide react with x mol of sodium hydroxide. Let's make a proportion:

1 2 = 0.062 x, x = 0.062 1 2 = 0.124 mol

Thus, a total of 0.1 + 0.124 = 0.224 mol of sodium hydroxide is needed. Let's determine the molar mass of sodium hydroxide:

M(NaOH) = 23 + 16 + 1 = 40 g/mol

Let's calculate the mass of sodium hydroxide:

m(NaOH) = ν (NaOH) M(NaOH) = 0.224 mol 40 g/mol = 8.96 g.

Let's calculate the mass of sodium hydroxide solution with concentration

40%, or 0.4.

m(solution) =	m(NaOH)	≈ 22.4 g
	c(NaOH)

Let's calculate the volume of the solution: sodium hydroxide:

V = m ρ = 1.22 4 g.4 / mLg = 16 ml

Answer: the mixture contains 30% aluminum and 70% aluminum oxide; you will need 16 ml of sodium hydroxide solution.

Problem No. 5

The substance obtained by calcining 1.28 g of copper in a stream of oxygen was converted into copper (II) chloride. Calculate what volume (in ml) of 4% hydrochloric acid (ρ = 1.02 g/cm3) was consumed and what is the mass of copper (II) chloride released.

When copper is heated with oxygen, copper(II) oxide is formed:

2Cu + O2 = 2CuO (1)

When copper (II) oxide reacts with hydrochloric acid, copper (II) chloride is formed:

CuO + 2HCl = CuCl2 + H2O

The molar mass of copper is 64 g/mol. Let's calculate the amount of copper substance:

According to the reaction equation (1), from 2 moles of copper 2 moles of copper (II) oxide are formed, which means that from 0.02 moles of copper 0.02 moles of copper (II) oxide are formed. According to equation (2), 1 mol of copper (II) oxide reacts with 2 mol of hydrogen chloride. Let 0.02 mol of copper(II) oxide react with x mol of hydrogen chloride. Let's make a proportion:

1 2 = 0, x 02, x = 0.02 1 2 = 0.04 mol

Let's determine the molar mass of hydrogen chloride:

M(HCl) = 1 +35.5 = 36.5 g/mol.

Let's calculate the mass of hydrogen chloride:

m(HCl) = ν (HCl) M(HCl) = 0.04 mol 36.5 g/mol = 1.46 g.

Let's calculate the mass of 4% hydrochloric acid:

Let's calculate the volume of hydrochloric acid:

According to the reaction equation (2), from 1 mole of copper (II) oxide, 1 mole of copper (II) chloride is formed, which means that from 0.02 moles of copper (II) oxide, 0.02 moles of copper (II) chloride are formed. Let's determine the molar mass of copper (II) chloride:

M(CuCl2) = 64 + 35.5 2 = 135 g/mol.

Let's calculate the mass of copper (II) chloride:

m(CuCl2) = ν (CuCl2) M(CuCl2) = 0.02 mol 135 g/mol = 2.7 g.

Answer: 35.8 ml of hydrochloric acid was consumed; 2.7 g of copper(II) chloride was formed.

Chapter VI. Nonmetals

Problems for §§1-3 (p. 140)

Question No. 1

How does the structure of atoms and simple substances of nonmetals differ from metals?

a) Atoms of most nonmetals have 4 or more electrons in the outer electron shell, while metal atoms have from one to three electrons in the outer shell.

b) Simple substances - metals always form a so-called metallic crystal lattice. Simple substances

– non-metals either form an atomic lattice (for example, carbon, silicon, sulfur, phosphorus) or have a molecular structure (for example, hydrogen, oxygen, nitrogen).

Question No. 2

Based on the periodic table, identify patterns observed when the redox properties of nonmetals change.

IN groups of the periodic table, when moving from top to bottom, the oxidizing properties of non-metals weaken, and accordingly, the reducing properties increase.

IN periods, the oxidizing properties of nonmetals increase from left to right.

Questions No. 4-5

What patterns are observed in the changes in the properties of acid oxides in periods and groups? Given the formulas of acidic

oxides: a) N2 O5, CO2, Cl2 O7 and SO3; b) P2 O5, As2 O5, N2 O5 and Sb2 O5. Ras-

put them in order of increasing acidic properties of the oxides.

The acidic properties of elemental oxides increase in periods from left to right and in groups from bottom to top. Therefore the order will be as follows:

a) CO2, N2 O5, SO3, Cl2 O7

b) Sb2 O5, AS2 O5, P2 O5, N2 O5

Question No. 6

Using the textbook table (p. 134), additionally write two or three equations of chemical reactions of acid oxides, not presented in the table, with bases, basic oxides, and water.

1) Reactions with bases:

SO3 + 2NaOH = Na2 SO4 + H2 O

P2 O5 + 6KOH = 2K3 PO4 + 3H2 O

2NO2 + 2NaOH = NaNO2 + NaNO3 + Н2 O

2) Reactions with basic oxides:

SO2 + CaO = CaSO3

P2 O5 + 3CaO = Ca3 (PO4)2

CO2 + Na2 O = Na2 CO3

3) Reactions with water:

Cl2 O7 + H2 O = 2HClO4 N2 O5 + H2 O = 2HNO3

Question No. 7

State the similar and distinctive chemical properties of sulfuric and nitric acids.

General properties. Concentrated sulfuric and nitric acids are strong oxidizing agents. In nitric acid, the oxidizing agent is nitrogen in the oxidation state +5, in sulfuric acid it is sulfur in the oxidation state +6:

Cu + 4НNO3 = Cu(NO3 )2 + 2NO2 + 2Н2 O

Cu + 2H2 SO4 = CuSO4 + SO2 + 2H2 O

Distinctive properties. Dilute sulfuric acid reacts with metals to release hydrogen, that is, hydrogen in the oxidation state +1 serves as the oxidizing agent.

Mg + H2 SO4 = MgSO4 + H2


Mg−2e		− → Mg

		− → H 2

In dilute nitric acid, the oxidizing agent is still nitrogen in the oxidation state +5. The composition of the reaction products depends on the concentration of the acid and the chemical activity of the metal:

3Zn + 8HNO3 = 3Zn(NO3 )2 + 2NO + 4H2 O 4Mg + 10HNO3 = 4Mg(NO3 )2 + N2 O + 5H2 O

Question No. 8

When concentrated sulfuric acid reacts with iron, the oxidation state of sulfur changes from +6 to +4. Write an equation.

2Fe + 6H2 SO4 = Fe2 (SO4 )3 + 3SO2 + 6H2 O

	2e − →
	2e − →
	−3e − →
	−3e − →

Question No. 9

Why are non-volatile hydrogen compounds so different from volatile hydrogen compounds?

Hydroxides can be thought of as the product of the addition (real or mental) of water to the corresponding oxides. Hydroxides are divided into bases, acids, and amphoteric hydroxides. Bases have the general composition M(OH)x, acids have the general composition HxCo. In molecules of oxygen-containing acids, the replaced hydrogen atoms are connected to the central element through oxygen atoms. In molecules of oxygen-free acids, hydrogen atoms are attached directly to a non-metal atom. Amphoteric hydroxides include primarily hydroxides of aluminum, beryllium and zinc, as well as hydroxides of many transition metals in intermediate oxidation states.
Based on solubility in water, soluble bases are distinguished - alkalis (formed by alkali and alkaline earth metals). The bases formed by other metals do not dissolve in water. Most inorganic acids are soluble in water. To insoluble in water inorganic acids only silicic acid H2SiO3 applies. Amphoteric hydroxides do not dissolve in water.

Chemical properties grounds.

All bases, both soluble and insoluble, have a common characteristic property - to form salts.
Let's consider the chemical properties of soluble bases (alkalis):
1. When dissolved in water, they dissociate to form a metal cation and a hydroxide anion. Change the color of the indicators: violet litmus - to blue, phenolphthalein - to crimson, methyl orange - to yellow, universal indicator paper - to blue.
2. Interaction with acid oxides:
alkali + acid oxide = salt.
3. Interaction with acids:
alkali + acid = salt + water.
The reaction between an acid and alkali is called a neutralization reaction.
4. Interaction with amphoteric hydroxides:
alkali + amphoteric hydroxide = salt (+ water)
5. Interaction with salts (subject to the solubility of the original salt and the formation of a precipitate or gas as a result of the reaction.
Let's consider the chemical properties of insoluble bases:
1. Interaction with acids:
base + acid = salt + water.
Polyacid bases are capable of forming not only intermediate, but also basic salts.
2. Heat decomposition:
base = metal oxide + water.

Chemical properties of acids.

All acids have a common characteristic property - the formation of salts when replacing hydrogen cations with metal/ammonium cations.
Let's consider the chemical properties of water-soluble acids:
1. When dissolved in water, they dissociate to form hydrogen cations and an acid residue anion. Change the color of the indicators to red (pink), with the exception of phenolphthalein (does not react to acids, remains colorless).
2. Interaction with metals in the activity series to the left of hydrogen (subject to the formation of a soluble salt):
acid + metal = salt + hydrogen.
When interacting with metals, the exceptions are oxidizing acids - nitric and concentrated sulfuric acids. Firstly, they also react with some metals that are to the right of hydrogen in the activity series. Secondly, hydrogen is never released in reaction with metals, but a salt of the corresponding acid, water and the reduction products of nitrogen or sulfur are formed, respectively.
3. Interaction with bases/amphoteric hydroxides:
acid + base = salt + water.
4. Interaction with ammonia:
acid + ammonia = ammonium salt
5. Interaction with salts (subject to the formation of gas or sediment):
acid + salt = salt + acid.
Polybasic acids are capable of forming not only intermediate, but also acidic salts.
Insoluble silicic acid does not change the color of indicators (a very weak acid), but is capable of reacting with alkali solutions with slight heating:
1. Interaction of silicic acid with alkali solution:
silicic acid + alkali = salt + water.
2. Decomposition (during long-term storage or heating)
silicic acid = silicon(IV) oxide + water.

Chemical properties of amphoteric hydroxides.

Amphoteric hydroxides are capable of forming two series of salts, since when reacting with alkalis they exhibit the properties of an acid, and when reacting with acids they exhibit the properties of a base.
Let's consider the chemical properties of amphoteric hydroxides:
1. Interaction with alkalis:
amphoteric hydroxide + alkali = salt (+ water).
2. Interaction with acids:
amphoteric hydroxide + acid = salt + water.

Before discussing the chemical properties of bases and amphoteric hydroxides, let's clearly define what they are?

1) Bases or basic hydroxides include metal hydroxides in the oxidation state +1 or +2, i.e. the formulas of which are written either as MeOH or Me(OH) 2. However, there are exceptions. Thus, the hydroxides Zn(OH) 2, Be(OH) 2, Pb(OH) 2, Sn(OH) 2 are not bases.

2) Amphoteric hydroxides include metal hydroxides in the oxidation state +3, +4, as well as, as exceptions, the hydroxides Zn(OH) 2, Be(OH) 2, Pb(OH) 2, Sn(OH) 2. Metal hydroxides in oxidation state +4, in Unified State Exam assignments do not occur, so they will not be considered.

Chemical properties of bases

All grounds are divided into:

Let us remember that beryllium and magnesium are not alkaline earth metals.

In addition to being soluble in water, alkalis also dissociate very well in aqueous solutions, while insoluble bases have a low degree of dissociation.

This difference in solubility and ability to dissociate between alkalis and insoluble hydroxides leads, in turn, to noticeable differences in their chemical properties. So, in particular, alkalis are more chemically active compounds and are often able to enter into reactions that insoluble bases do not.

Interaction of bases with acids

Alkalis react with absolutely all acids, even very weak and insoluble ones. For example:

Insoluble bases react with almost all soluble acids, but do not react with insoluble silicic acid:

It should be noted that both strong and weak bases with the general formula of the form Me(OH) 2 can form basic salts when there is a lack of acid, for example:

Interaction with acid oxides

Alkalis react with all acidic oxides, forming salts and often water:

Insoluble bases are capable of reacting with all higher acid oxides corresponding to stable acids, for example, P 2 O 5, SO 3, N 2 O 5, to form medium salts:

Insoluble bases of the type Me(OH) 2 react in the presence of water with carbon dioxide exclusively to form basic salts. For example:

Cu(OH) 2 + CO 2 = (CuOH) 2 CO 3 + H 2 O

Due to its exceptional inertness, only the strongest bases, alkalis, react with silicon dioxide. In this case, normal salts are formed. The reaction does not occur with insoluble bases. For example:

Interaction of bases with amphoteric oxides and hydroxides

All alkalis react with amphoteric oxides and hydroxides. If the reaction is carried out by fusing an amphoteric oxide or hydroxide with a solid alkali, this reaction leads to the formation of hydrogen-free salts:

If aqueous solutions of alkalis are used, then hydroxo complex salts are formed:

In the case of aluminum, under the action of an excess of concentrated alkali, instead of Na salt, Na 3 salt is formed:

Interaction of bases with salts

Any base reacts with any salt only if two conditions are met simultaneously:

1) solubility of the starting compounds;

2) the presence of precipitate or gas among the reaction products

For example:

Thermal stability of substrates

All alkalis, except Ca(OH) 2, are resistant to heat and melt without decomposition.

All insoluble bases, as well as slightly soluble Ca(OH) 2, decompose when heated. Most heat decomposition of calcium hydroxide – about 1000 o C:

Insoluble hydroxides have much lower decomposition temperatures. For example, copper (II) hydroxide decomposes already at temperatures above 70 o C:

Chemical properties of amphoteric hydroxides

Interaction of amphoteric hydroxides with acids

Amphoteric hydroxides react with strong acids:

Amphoteric metal hydroxides in the oxidation state +3, i.e. type Me(OH) 3, do not react with acids such as H 2 S, H 2 SO 3 and H 2 CO 3 due to the fact that the salts that could be formed as a result of such reactions are subject to irreversible hydrolysis to the original amphoteric hydroxide and corresponding acid:

Interaction of amphoteric hydroxides with acid oxides

Amphoteric hydroxides react with higher oxides, which correspond to stable acids (SO 3, P 2 O 5, N 2 O 5):

Amphoteric metal hydroxides in the oxidation state +3, i.e. type Me(OH) 3, do not react with acidic oxides SO 2 and CO 2.