Sodium ferrate properties and contraindications. Development of an effective method for obtaining ferrates (vi) of alkali metals

d- element of group VIII; serial number - 26; atomic mass - 56; (26p11; 30n01), 26ē

1s22s22p63s23p63d64s2

Medium activity metal, reducing agent.

Main oxidation states - +2, +3

Iron and its compounds

Chemical properties

1) In air, iron is easily oxidized in the presence of moisture (rusting):

4Fe + 3O2 + 6H2 O ® 4Fe(OH)3

A heated iron wire burns in oxygen, forming scale - iron oxide (II, III):

3Fe + 2O2 ® Fe3O4

2) At high temperatures (700–900°C), iron reacts with water vapor:

3Fe + 4H2O –t°® Fe3O4 + 4H2

3) Iron reacts with non-metals when heated:

Fe + S –t°® FeS

4) Iron dissolves easily in hydrochloric and dilute sulfuric acids:

Fe + 2HCl ® FeCl2 + H2

Fe + H2SO4(dec.) ® FeSO4 + H2

In concentrated oxidizing acids, iron dissolves only when heated.

2Fe + 6H2SO4(conc.) –t°® Fe2(SO4)3 + 3SO2 + 6H2O

Fe + 6HNO3(conc.) –t°® Fe(NO3)3 + 3NO2 + 3H2O

(in the cold, concentrated nitric and sulfuric acids passivate iron).

5) Iron displaces metals to the right of it in a series of stresses from solutions of their salts.

Fe + CuSO4 ® FeSO4 + Cu¯

Ferrous compounds

Iron(II) hydroxide

It is formed by the action of alkali solutions on iron (II) salts without air access:

FeCl + 2KOH ® 2KCl + Fe(OH)2¯

Fe(OH)2 is a weak base, soluble in strong acids:

Fe(OH)2 + H2SO4 ® FeSO4 + 2H2O

Fe(OH)2 + 2H+ ® Fe2+ + 2H2O

When Fe (OH) 2 is calcined without air access, iron oxide (II) FeO is formed:

Fe(OH)2 –t°® FeO + H2O

In the presence of atmospheric oxygen, a white precipitate of Fe (OH) 2, oxidizing, turns brown - forming iron (III) hydroxide Fe (OH) 3:

4Fe(OH)2 + O2 + 2H2O ® 4Fe(OH)3

Iron (II) compounds have reducing properties, they are easily converted into iron (III) compounds under the action of oxidizing agents:

10FeSO4 + 2KMnO4 + 8H2SO4 ® 5Fe2(SO4)3 + K2SO4 + 2MnSO4 + 8H2O

6FeSO4 + 2HNO3 + 3H2SO4 ® 3Fe2(SO4)3 + 2NO + 4H2O

Iron compounds are prone to complex formation (coordination number = 6):

FeCl2 + 6NH3 ® Cl2

Fe(CN)2 + 4KCN ® K4(yellow blood salt)

Qualitative reaction for Fe2+

Under the action of potassium hexacyanoferrate (III) K3 (red blood salt) on solutions of ferrous salts, a blue precipitate (turnbull blue) is formed:

3FeSO4 + 2K3 ® Fe32¯ + 3K2SO4

3Fe2+ + 3SO42- +6K+ + 23- ® Fe32¯ + 6K+ + 3SO42-

3Fe2+ + 23- ®Fe32¯

Ferric compounds

Iron(III) oxide

It is formed during the combustion of iron sulfides, for example, during the firing of pyrite:

4FeS2 + 11O2 ® 2Fe2O3 + 8SO2

or when calcining iron salts:

2FeSO4 –t°® Fe2O3 + SO2 + SO3

Fe2O3 - basic oxide, showing amphoteric properties to a small extent

Fe2O3 + 6HCl –t°® 2FeCl3 + 3H2O

Fe2O3 + 6H+ –t°® 2Fe3+ + 3H2O

Fe2O3 + 2NaOH + 3H2O –t°® 2Na

Fe2O3 + 2OH- + 3H2O ® 2-

Iron(III) hydroxide

It is formed by the action of alkali solutions on ferric iron salts: it precipitates as a red-brown precipitate

Fe(NO3)3 + 3KOH ® Fe(OH)3¯ + 3KNO3

Fe3+ + 3OH- ® Fe(OH)3¯

Fe(OH)3 is a weaker base than iron(II) hydroxide.

This is explained by the fact that Fe2+ has a smaller ion charge and a larger radius than Fe3+, and therefore, Fe2+ retains hydroxide ions weaker, i.e. Fe(OH)2 dissociates more easily.

In this regard, iron (II) salts are hydrolyzed slightly, and iron (III) salts are very strongly hydrolyzed. For a better understanding of the materials in this section, it is recommended to watch the video clip (only available on CDROM). The color of solutions of Fe(III) salts is also explained by hydrolysis: despite the fact that the Fe3+ ion is almost colorless, the solutions containing it are colored yellow-brown, which is explained by the presence of iron hydroxoions or Fe(OH)3 molecules, which are formed due to hydrolysis:

Fe3+ + H2O « 2+ + H+

2+ + H2O « + + H+

H2O « Fe(OH)3 + H+

When heated, the color darkens, and when acids are added, it becomes lighter due to the suppression of hydrolysis. Fe (OH) 3 has a weakly pronounced amphotericity: it dissolves in dilute acids and in concentrated alkali solutions

Fe(OH)3 + 3HCl ® FeCl3 + 3H2O

Fe(OH)3 + 3H+ ® Fe3+ + 3H2O

Fe(OH)3 + NaOH ® Na

Fe(OH)3 + OH- ® -

Iron (III) compounds are weak oxidizing agents, they react with strong reducing agents:

2Fe+3Cl3 + H2S-2 ® S0 + 2Fe+2Cl2 + 2HCl

Qualitative reactions for Fe3+

1) Under the action of potassium hexacyanoferrate (II) K4 (yellow blood salt) on solutions of ferric salts, a blue precipitate (Prussian blue) is formed:

4FeCl3 +3K4 ® Fe43¯ + 12KCl

4Fe3+ + 12Cl- + 12K+ + 34- ® Fe43¯ + 12K+ + 12Cl-

4Fe3+ + 3 4- ®Fe43¯

2) When potassium or ammonium thiocyanate is added to a solution containing Fe3+ ions, an intense blood-red color of iron(III) thiocyanate appears: FeCl3 + 3NH4CNS « 3NH4Cl + Fe(CNS)3 (when interacting with Fe2+ thiocyanates, the solution remains almost colorless) .

Ferrates are salts containing the ferrate anion FeO42-. Correspond to iron acid H2FeO4, which does not exist in free form. As a rule, they have a purple color. Contents

Properties

Ferrates are the strongest oxidizing agents. The redox potential of the ferrate ion

FeO42- + 8H+ + 3e- = Fe3+ + 4H2O E0 = +2.2V

FeO42- + 4H2O + 3e- = Fe(OH)3 + 5OH- E0 = +0.72V

IN acidic environment decompose with the release of oxygen:

4FeO42- + 20H+ = 4Fe3+ + 3O2 + 10H2O

Ferrates also decompose slowly in a neutral environment:

4FeO42- + 10H2O = 4Fe(OH)3 + 3O2 + 8OH-

The solubility of ferrates is close to that of sulfates. So, potassium ferrate is soluble quite well, and barium ferrate is insoluble.

Application

Being strong oxidizing agents, ferrates easily oxidize organic pollutants and have an antiseptic effect. However, unlike chlorine, they do not form toxic products. Therefore, ferrates are increasingly being used in water treatment and water treatment.

Receipt

There are several ways to synthesize ferrates.

The first method is the oxidation of iron (III) hydroxide with chlorine or hypochlorite in a strongly alkaline medium:

2Fe(OH)3 + 3Cl2 + 10OH- = 2FeO42- + 6Cl- + 8H2O

2Fe(OH)3 + 3ClO- + 4OH- = 2FeO42- + 3Cl- + 5H2O

The second method is the electrolysis of an alkali solution on an iron anode:

Fe + 2KOH + 2H2O = K2FeO4 + 3H2

Question #72
Cobalt and nickel are elements of group VIIIB (9, 10). Electronic configurations

valence levels: Co - 3 d 74s 2, Ni - 3 d 84s 2. Cobalt and nickel are characterized by degrees

+2 and +3 oxidation, and in aqueous solutions the most stable degree is

oxidation +2.

Simple substances Co and Ni in powder form show a rather high

activity towards acids. As a result of their interaction with acids,

salts with an oxidation state of +2. Cobalt salts are pink in color due to

formation of an aqua complex 2–, and aqueous solutions of Ni salts are colored green

due to the presence of the ion 2–.

E + 2HCl = CoCl2 + H2

E + H2SO4 = CoSO4 + H2

3E + 8HNO3(razb.) = 3Co(NO3)2 + 2NO + 4H2O

Cold concentrated nitric acid passivates Co and Ni. Protective when heated

the film is destroyed and both metals react with concentrated nitric acid:

E + 4HNO3 (conc.) = E (NO3) 2 + 2NO2 + 2H2O

With oxygen, cobalt and nickel form EO oxides with basic properties. These oxides are

dissolve in water, do not interact with alkalis, but easily react with acids,

forming E(II) salts.

Co(II) or Ni(II) salts are most often used for the synthesis of the corresponding hydroxides,

for example:

ECl2 + NaOH = E(OH)2↓ + NaCl

Upon receipt of cobalt (II) hydroxide from salts, a blue precipitate is first formed

sparingly soluble basic salts Co(OH)nX2-n⋅ x H2O and then pink hydroxide Co(OH)2.

The appearance of a blue color can also be explained by the formation of cobalt hydroxide.

composition 3Co(OH)2⋅2H2O, which is formed together with basic salts. With further

adding alkali as a result of dehydration and aging, it changes color from blue to

Cobalt(II) hydroxide shows slight signs of amphotericity with predominantly

basic properties. It readily dissolves in acids (with the formation of Co(II) salts), and

dissolution in alkali goes with great difficulty. However, the presence of acidic properties of Co(OH)2

confirmed by the existence of the hydroxocomplex 2-.

Co(II) hydroxide is very slowly oxidized by atmospheric oxygen and turns into hydroxide

Co(III) colored brown:

4Co(OH)2 + O2 + 2H2O = 4Co(OH)3

In the presence of stronger oxidizing agents, such as hydrogen peroxide, the oxidation process

Co(II) goes much faster:

2Co(OH)2 + H2O2 = 2Co(OH)3

A qualitative reaction to the Co(II) ion is the reaction of the formation of its nitro complex

yellow color.

CoCl2 + 7KNO2 + 2CH3COOH = K3 ↓ + NO + 2CH3COOK + 2KCl

The oxidation state (III) is unstable for cobalt, so Co(III) hydroxide

exhibits oxidizing properties, even under the influence of such a weak reducing agent as

2Co(OH)3 + 6HCl = 2CoCl2 + Cl2 + 6H2O

Nickel (II) hydroxide, green in color, similar in acid-base properties to

hydroxide Co(II). It is easily soluble in acids and practically insoluble in alkalis.

Prolonged exposure of alkalis to the Ni(OH)2 precipitate leads to the formation of a hydroxo complex

indefinite composition with the conditional formula 2−.

Nickel(II) hydroxide is not oxidized to Ni(OH)3 by either air oxygen or peroxide

hydrogen. To oxidize it, a stronger oxidizing agent is needed, for example, bromine:

2Ni(OH)2 + 2NaOH + Br2 = 2Ni(OH)3 + 2NaBr

Nickel and cobalt in oxidation states +2 and +3 form a large number of complex

connections. Their most stable cationic complexes are aqua complexes and

ammonia, as well as complexes where the ligands are polydentate organic

molecules such as dimethylglyoximate. Formation of an insoluble complex

bright red dimethylglyoximate is qualitative reaction for nickel (II):

H3C O− …HO CH3

C=N N=C

 Ni2+ 

C=N N=C

H3COH…..O− CH3

NiCl2 + 2NH4OH + 2(CH3CNOH)2 = (CH3CNO)4H2N+ 2NH4Cl + 2H2O (Chugaev reaction)

Cobalt and nickel form a large number of insoluble salts, many of which,

for example, phosphates can be synthesized using exchange reactions in aqueous solutions:

3MeCl2 + 4Na2HPO4 = 2Me3(PO)4 + 8NaCl + HCl

Medium Co(II) or Ni(II) carbonates by adding an alkali metal carbonate to their solutions

salt is not available. Due to increased hydrolysis in the presence of carbonate ions,

there are processes with the formation of poorly soluble basic, and not medium carbonates:

2CoCl2 + Na2CO3 + 2H2O → (CoOH)2CO3↓ + 2NaCl + 2HCl

2NiCl2 + Na2CO3 + 2H2O → (NiOH)2CO3↓ + 2NaCl + 2HCl

Question 73. general characteristics elements of the platinum family.

The platinum subgroup includes 6 transition metals. According to the number of electrons in 4d3s orbitals (Ru, Rh, Pd) and 5d6s orbitals (Os, Ir, Pt) and by analogy of physicochemical properties, all elements of group 8B are divided into three subgroups:

1) Ru- Os 2) Rh- Ir 3) Pd- Pt The atomic radius of all 6 elements varies in a small range: 134pm (Ru) – 139pm (Pt). This causes the proximity of the properties of all 6 elements.

In the electrochemical series, all platinum metals come after hydrogen. In terms of electronegativity, all elements of the group are closer to non-metals than to metals. Therefore, the compounds of these elements exhibit amphotericity, expressed to varying degrees. The hydroxides of not all of these elements are soluble in both acid and alkali. Nevertheless, elements of the platinum family form not only cationic, but also anionic complexes. Stable valence compounds for elements of the platinum family are as follows: ruthenium-4,6,8, Rhodium - 3.4; palladium - 2.4, osmium-4.6, iridium-3.4, platinum -2.4.

Hydroxides of elements in the tetravalent state exist in the format MO 2 *nH 2 O where n=2 (for platinum 2 and 3). The water content depends on the temperature, the higher the temperature, the less water.

Hydroxides of rhenium, palladium and platinum dissolve in acids and alkalis

PtO 2 * 3 H 2 O + 2NaOH → Na 2 + H 2 O

PtO 2 * 3 H 2 O + 6HCl → H 2 + 5H 2 O

Under normal conditions, platinum metals do not interact with such strong oxidizing agents as F2, Cl2, O2. The low reactivity of elementary substances is determined by the high binding energy in the crystal lattice. The same reason determines the high melting point and high density values.

Only platinum reacts without heating with an oxidizing mixture of acids.

3Pt + 18HCl + 4HNO 3 \u003d 3H 2 + 4NO + 8H 2 O

Or with hydrochloric acid in the presence of oxygen

Pt + 6HCl + O 2 \u003d H 2 + 2H 2 O

All metals of the platinum family, except for iridium, pass into the 4-valence state when fused with alkaline oxidizing mixtures, for example:

Ru + 2KOH + 3KNO 3 → K 2 RuO 4 + 3KNO 2 + H 2 O

Iridium goes into a 3-valent form.

When heated, platinum metals react with NaCl or HCl in a stream of chlorine, which leads to the formation of a complex

Ir + 2NaCl +2 Cl 2 →Na 2

Platinum can form a cyanide complex when heated.

Pt+ 6KCN+4H 2 O→K 2 + 4KOH +2H 2

Elements of the platinum family form complex compounds with coordination numbers of 4 and 6. The most studied are cyanide, halide, and ammonia complexes. Complex compounds can be cationic, anionic or neutral.

Cation complex

2NH3 =

Neutral

Cl 2 +2HCl \u003d + 2NH 4 Cl

Cationic anionic

Cl 2 + K 2 \u003d [Pt (NH 3) 4] + 2KCl

Question 74.d - elements of the first group. General characteristics of the group. Physical and chemical properties of simple substances.

In the 1B group (copper group) there are transition metals Cu, Ag, Au, which have a similar distribution of electrons, determined by the phenomenon of "breakthrough" or "failure" of electrons.

The phenomenon of "leakage" is a symbolic transfer of one of the two valence s-electrons to the d-sublevel, which reflects the non-uniform retention of external electrons by the nucleus.

The transition of one s-electron to the outer level leads to the stabilization of the d-sublevel. Therefore, depending on the degree of excitation, atoms of the 1B group can donate from 1 to 3 electrons for the formation of a chemical bond. As a result, elements of group 1B can form compounds with oxidation states +1, +2, +3. However, there are differences: for mei, the most stable oxidation states are +1 and +2, for silver +1, and for gold +1 and +3. The most characteristic coordination numbers in this group are 2,3,4.
Elements of group 1B are relatively inert. In the electrochemical series, they stand after hydrogen, which is manifested in their weak reducing ability. Therefore, in nature, they are found in native form.

Restorative and basic properties decrease from copper to gold, molar mass in this direction increases, the density increases, the ionization energy increases in the order of silver-copper-gold.

Gold does not corrode

Chemical properties:

Silver oxide Ag 2 O is obtained by heating silver with oxygen or by treating AgNO 3 solutions with alkalis

2 AgNO 3 + 2KOH → Ag 2 O + 2KNO 3 + H 2 O

Ag 2 O dissolves slightly in water, however, due to hydrolysis, the solutions have an alkaline reaction

Ag 2 O + H 2 O → 2Ag + +2 OH -

In cyanide solutions it turns into a complex

Ag 2 O+4KCN +H 2 O→2K + 2KOH

Ag 2 O is an energetic oxidizing agent. Oxidizing properties are determined by the use of its suspension as an antiseptic.

In the electrochemical series of normal redox potentials, silver comes only after hydrogen. Therefore, metallic silver reacts only with oxidizing concentrated nitric and sulfuric acids.

2Ag + 2H 2 SO 4 → Ag 2 SO 4 + SO 2 + 2H 2 O
most silver salts are poorly or slightly soluble. Practically insoluble halides, phosphates. Silver sulfate and silver carbonate are poorly soluble. Solutions of silver halides decompose under the action of ultraviolet and X-rays:

2AgCl→2Ag +Cl 2

Insoluble silver chloride and silver bromide dissolve in ammonia to form ammonia:

AgCl + 2NH 3 →Cl

Ag blackens Ag + H 2 S + O 2 \u003d Ag 2 S ↓ + H 2 O

To lighten: dip in ammonia solution.

AgNO 3 + NaOH \u003d AgOH (AgO + H 2 O) + NaNO 3

Ag 2 O + H 2 O 2 \u003d 2 Ag + H 2 O + O 2

Copper(1) forms insoluble halides. These salts dissolve in ammonia and form complexes

CuCl + 2NH 3 \u003d Cl

Copper (2) oxide and hydroxide are insoluble in water, they have a basic character and dissolve in acids

Cu(OH) 2 + 2HCl +4H 2 O →Cl 2

Copper (2) hydroxide dissolves in ammonia, forms a complex that turns the solution blue:

Cu (OH) 2 + 2HCl + 4H 2 O \u003d (OH) 2 this reaction is used as a qualitative reaction for copper (2) ions

Copper in humid air reacts as follows:

Cu + H 2 O + CO 2 + O 2 \u003d (CuOH) 2 CO 3

Copper dissolves in sulfuric acid and nitric

Salts of copper, silver and gold interact with alkali metal sulfides and hydrogen sulfides to form water-insoluble precipitates Ag 2 S, Cu 2 S, CuS, Au 2 S 3

The most common gold(3) compound is AuCl 3 chloride, which is readily soluble in water.

Gold oxide and hydroxide are amphoteric compounds with more pronounced acidic properties. Gold hydroxide does not dissolve in water, but dissolves in alkalis with the formation of a hydroxo complex:

AuO (AuOH) + NaOH + H 2 O \u003d Na

Reacts with acids to form an acid complex

AuO(AuOH)+2H 2 SO 4 →H + 2H 2 O

Gold dissolves in aqua regia

Au + 4HCl + HNO 3 \u003d H + NO + 2H 2 O

Gold salts with an oxidation state of +3 are completely hydrolyzed

Physical properties

Copper- a golden-pink ductile metal, in air it quickly becomes covered with an oxide film, which gives it a characteristic intense yellowish-red tint. Thin films of copper in the light have a greenish-blue color.

Copper has a high thermal and electrical conductivity (it ranks second in electrical conductivity among metals after silver). Electrical conductivity at 20 °C 55.5-58 MS/m. Copper has a relatively large temperature coefficient of resistance: 0.4%/°C and, over a wide temperature range, is weakly dependent on temperature.

Pure silver- rather heavy (lighter than lead, but heavier than copper), unusually ductile silver-white metal (light reflection coefficient is close to 100%). Thin silver foil is purple in transmitted light. Over time, the metal tarnishes, reacting with traces of hydrogen sulfide in the air and forming a deposit of sulfide, whose thin film then gives the metal its characteristic pinkish color. It has high thermal conductivity. At room temperature, it has the highest electrical conductivity among all known metals (electrical resistivity 1.59 10 −8 Ohm m at 20 °C). Melting point 962°C.

Pure gold- soft yellow metal. The reddish hue of some gold products, such as coins, is given by impurities of other metals, in particular copper. Gold has exceptionally high thermal conductivity and low electrical resistance.

Gold is a very heavy metal: the density of pure gold is 19321 kg/m³ (a ball of pure gold with a diameter of 46 mm has a mass of 1 kg). Among metals, it ranks sixth in density: after osmium, iridium, rhenium, platinum and plutonium. The high density of gold makes it easier to mine. The simplest technological processes, such as, for example, flushing at locks, can provide a very high degree of recovery of gold from the washed rock.

Gold is also highly ductile: it can be forged into sheets up to ~0.1 µm thick (gold leaf); with such a thickness, gold is translucent and in reflected light it has a yellow color, in transmitted light it is colored bluish-greenish in addition to yellow. Gold can be drawn into wire with a linear density of up to 500 m/g.

The melting point of gold is 1064 °C. The density of liquid gold is less than that of solid gold, and is 17 g/cm3 at the melting point. Liquid gold is quite volatile, and actively evaporates long before the boiling point.

Compounds of copper (I) and copper (II), their KO and OV characteristics, ability to complex formation. Complex compounds of copper (II) with ammonia, amino acids, polyhydric alcohols. The complex nature of copper-containing enzymes and the chemistry of their action in metabolic reactions. Chemical basis for the use of copper compounds in medicine and pharmacy.

Copper compounds

Copper (I) oxide Cu2O3 and copper oxide (I) Cu2O, like other copper (I) compounds, are less stable than copper (II) compounds. Copper (I) oxide, or cuprous oxide Cu2O, occurs naturally in the form of the mineral cuprite. In addition, it can be obtained as a precipitate of red copper (I) oxide by heating a solution of copper (II) salt and alkali in the presence of a strong reducing agent.

Copper (II) oxide, or copper oxide, CuO is a black substance found in nature (for example, in the form of the mineral tenerite). It is obtained by calcining copper(II) hydroxocarbonate (CuOH)2CO3 or copper(II) nitrate Cu(NO2)2. Copper (II) oxide is a good oxidizing agent.

Copper (II) hydroxide Cu (OH) 2 is precipitated from solutions of copper (II) salts under the action of alkalis in the form of a blue gelatinous mass. Already at low heating, even under water, it decomposes, turning into black oxide of copper (II).

Copper(II) hydroxide is a very weak base. Therefore, solutions of copper (II) salts in most cases are acidic, and with weak acids, copper forms basic salts.

Copper (II) sulfate CuSO4 in the anhydrous state is a white powder that turns blue when water is absorbed. Therefore, it is used to detect traces of moisture in organic liquids. An aqueous solution of copper sulfate has a characteristic blue-blue color. This color is characteristic of hydrated 2+ ions; therefore, all dilute solutions of copper (II) salts have the same color, unless they contain any colored anions. From aqueous solutions, copper sulfate crystallizes with five molecules of water, forming transparent blue crystals of copper sulphate.

Copper sulphate is used for the electrolytic coating of metals with copper, for the preparation of mineral paints, and also as a starting material in the preparation of other copper compounds. IN agriculture A dilute solution of copper sulphate is used to spray plants and dress grains before sowing to kill spores of harmful fungi.

Copper (II) chloride CuCl2. 2H2O. Forms dark green crystals, easily soluble in water. Very concentrated solutions of copper chloride (II) are green, dilute - blue-blue.

Copper(II) nitrate Cu(NO3)2.3H2O. Obtained by dissolving copper in nitric acid. When heated, blue crystals of copper nitrate first lose water, and then easily decompose with the release of oxygen and brown nitrogen dioxide, turning into copper (II) oxide.

Copper(II) hydroxocarbonate (CuOH)2CO3. It occurs naturally in the form of the mineral malachite, which has a beautiful emerald green color. It is artificially prepared by the action of Na2CO3 on solutions of copper (II) salts.

2CuSO4 + 2Na2CO3 + H2O = (CuOH)2CO3↓ + 2Na2SO4 + CO2

It is used to obtain copper chloride (II), for the preparation of blue and green mineral paints, as well as in pyrotechnics.

Copper(II) acetate Cu(CH3COO)2.H2O. Obtained by treating metallic copper or copper (II) oxide with acetic acid. Usually it is a mixture of basic salts of various composition and color (green and blue-green). Under the name verdigris, it is used for the preparation of oil paint.

Copper complex compounds are formed as a result of the combination of doubly charged copper ions with ammonia molecules.

Copper is an essential element for all higher plants and animals. In the blood stream, copper is transported primarily by the protein ceruloplasmin. After absorption of copper by the intestines, it is transported to the liver with the help of albumin. Copper is found in a wide variety of enzymes, such as cytochrome c oxidase, the copper-zinc enzyme superoxide dismutase, and the oxygen-carrying protein hemocyanin. In the blood of most molluscs and arthropods, copper is used instead of iron to transport oxygen.

It is assumed that copper and zinc compete with each other during absorption in the digestive tract, so an excess of one of these elements in the diet can cause a deficiency of the other element. A healthy adult needs 0.9 mg of copper per day.

Toxicity

Some copper compounds can be toxic if the MPC is exceeded in food and water. The content of copper in drinking water should not exceed 2 mg / l, however, a lack of copper in drinking water is also undesirable. The World Health Organization formulated this rule in 1998 as follows: “The risks to human health from a lack of copper in the body are many times higher than the risks from its excess.”

In 2003, as a result of intensive research, WHO revised previous assessments of copper toxicity. It has been recognized that copper is not the cause of digestive disorders.

There were fears that hepatocerebral dystrophy is accompanied by the accumulation of copper in the body, since it is not excreted by the liver into bile. This disease causes damage to the brain and liver. However, a causal relationship between the onset of the disease and copper intake was not confirmed. Only an increased sensitivity of persons diagnosed with this disease to an increased content of copper in food and water has been established.

Bactericidal

The bactericidal properties of copper and its alloys have been known to man for a long time. In 2008, after lengthy research, the US Federal Environmental Protection Agency officially granted copper and several copper alloys the status of substances with a germicidal surface. Particularly pronounced bactericidal action copper surfaces are shown against a methicillin-resistant strain of Staphylococcus aureus, known as the "supermicrobe" MRSA. In the summer of 2009, the role of copper and copper alloys in the inactivation of influenza A/H1N1 virus was established.

Organoleptic properties

Copper ions impart a distinct "metallic taste" to excess copper in water. In different people, the threshold for the organoleptic determination of copper in water is approximately 2-10 mg / l. The natural ability to detect high levels of copper in water in this way is a natural defense mechanism against the ingestion of water with excessive copper content.

DEVELOPMENT OF AN EFFICIENT METHOD FOR OBTAINING FERRATES (VI) ALKALI METALS

Kaztaev Alimzhan

Tokbanov Bauyrzhan

class 10, AOSCTLIOYU,G.Aktobe, Kazakhstan

Agisheva Almagul Abilkairovna

scientific director,cand. chem. Sciences, Senior Lecturer, Aktobe State pedagogical institute, Kazakhstan

Based on the analysis of literature data, taking into account the shortcomings existing methods For the production of ferrates, an electrolytic method for the anodic oxidation of iron was proposed and the operating parameters of the corresponding electrolytic unit were refined. The aim of the work was to create a simple method for obtaining compounds of hexavalent iron, reproducible in any industrial and scientific laboratory.

According to literature data, ammonium, copper, iron ferrates are extremely unstable, potassium (VI) ferrate is more stable than sodium (VI) ferrate. Dry potassium ferrate is quite stable - no signs of decomposition are observed after two days of standing in air. The solubility of potassium ferrate is relatively low; when the solution is cooled, the precipitate can be separated and dried. Sodium ferrate is highly soluble. Dry salt is hygroscopic. It is quite difficult to isolate it from the solution. Potassium salt during drying can be washed with alcohol. The sodium salt under these conditions oxidizes the alcohol.

Thus, the preparation of sodium ferrate is associated with certain difficulties. On the other hand, the low solubility of potassium ferrate in cold solutions limits its use in wastewater treatment plants. Therefore, in accordance with the purpose of the work, we chose the sodium salt of hexavalent iron as the electrolysis product. It is known that gaseous chlorine in alkaline solutions of an iron (III) compound oxidizes it to iron (VI) compounds - that is, to ferrates. Therefore, for a more complete oxidation of iron, sodium chloride was added to the reaction vessel, during the electrolysis of which chlorine was formed.

According to the method proposed by us, 1 liter of distilled water, 150 grams of sodium hydroxide NaOH of the pure grade, 2 iron electrodes (anode, cathode), a lead-acid battery and a volt-ammeter (multimeter) were taken. In subsequent experiments, sodium chloride of chemically pure grade weighing 20 grams was added to the electrolyte.

Capacitor, accumulator, transformer plates, metal scrap served as electrodes.

The electrolysis time varied from half an hour to three hours.

The current strength was 2-15 Amperes.

The current density is 0.2-0.4 A/cm 2 and above.

The qualitative determination of sodium ferrate (VI) was carried out by the presence or absence of a characteristic purple color of the solution, by the reaction of displacement of chlorine from a solution of concentrated hydrochloric acid, by the exchange reaction with barium chloride.

The quantitative determination of the product yield consisted in carrying out the exchange reaction of 10 ml of anolyte with an excess of 0.1 M barium chloride, filtering the formed dark red precipitate of barium ferrate, washing it with alcohol, drying it on a Buchner funnel with a water jet pump, and keeping it for a day in a vacuum desiccator with subsequent weighing.

In the process of electrolysis, after the current is applied from the anode, purple stains begin to go, then a stable light purple color appears. After five minutes, the color is already dark purple in the case of separation of the anode and cathode spaces. In this case, the evolution of gas bubbles can be observed in the cathode space.

Anode process: Fe 0 + 8 OH - - 6e - = FeO 4 2- + 4H 2 O

Cathodic process: 2H + + 2 e - = H 2

In total, the reaction looks like this: Fe 0 + 2NaOH + 2H 2 O → Na 2 FeO 4 + 2H 2

After the termination of electrolysis, the anolyte was poured into an open container. When standing, a slow release of gas bubbles was observed - sodium ferrate is unstable. There is a hydrolysis reaction, accompanied by the reduction of iron and the release of oxygen.

4Na 2 FeO 4 + 10 H 2 O \u003d 8NaOH + 4Fe (OH) 3 + 3O 2

When the anolyte is diluted with water, decomposition is more noticeable, since flakes of Fe (OH) 3 appear, the violet color of the ferrate turns into a rusty color, and after 4 hours a precipitate of iron hydroxide (III) collects on the bottom of the dish. Dilution with alkali stabilizes the ferrate. Dry sodium ferrate could not be isolated due to the rapid decomposition and high solubility of the salt.

When concentrated hydrochloric acid is added to the anolyte, chlorine is released, which can be identified by a sharp characteristic odor.

2Na 2 FeO 4 + 16HCl = 3Cl 2 + 4NaCl + 2FeCl 3 + 8H 2 O

The exchange reaction of K 2 FeO 4 with barium chloride gave a pink-brownish precipitate of insoluble barium ferrate. It is more stable and was isolated by filtration and subsequent drying.

To prevent alkali from absorbing CO 2 from the air with the transition to carbonate and accelerating hydrolysis, the anolyte was subsequently kept in a flask with a ground-in lid to restrict air access.

Based on the results of several experiments, the masses of the resulting sodium ferrate were calculated.

1) m(BaFeO4) = 0.8013 g

Na 2 FeO 4 - BaFeO 4

166 g/mol - 257 g/mol

X g - 0.8013

X= 0.5175g

m(Na 2 FeO 4) anolyte = 31.05 g

2) m(BaFeO4) = 0.9763 g

Na 2 FeO 4 - BaFeO 4

166 g/mol - 257 g/mol

X g - 0.9763

X= 0.6300g

m(Na 2 FeO 4) anolyte = 37.8 g

3) m(BaFeO4) = 0.8905 g

Na 2 FeO 4 - BaFeO 4

166 g/mol - 257 g/mol

X g - 0.8905

X= 0.5750 g

m(Na 2 FeO 4) anolyte = 34.5 g

For the exchange reaction, the anolyte was taken in a volume of 10 ml, the volume of the anolyte was only 600 ml. Given this ratio, the mass of the synthesized sodium ferrate was calculated. The results show a certain convergence (31.05 g; 37.8 g; 34.5 g).

According to Faraday's law, the mass of sodium ferrate is:

m= 0.0002867x10800x15=46.450 g

The calculated value differs from the calculated one analytical method. This is due to the instability of the product and a high degree its hydrolysis, in addition, unaccounted for electrode processes can occur.

Thus, the resulting alkaline solution of sodium ferrate is not subject to long-term storage and must be used within four hours after preparation. Precipitation of barium ferrate, planted 1 hour after preparation, is approximately four times less by weight than precipitation planted immediately after the termination of electrolysis. As a result of the high oxidizing power, sodium ferrate in an alkaline solution decomposes by 25% already in 1 hour.

As a result of the work carried out, a new anodic method for obtaining sodium ferrate was developed. The method is characterized by the use of inexpensive reagents, scrap metal, improvised current sources. It is planned to continue this work in order to obtain reproducible results, improve the design of the electrolytic cell and then propose the scope of its application, depending on the requests of potential users.

Bibliography:

1. Armoes P., Henze M., Lyakuryansen J., Arvan E. Wastewater treatment // M.: Mir, 2004. - P. 20-60.

2. Kaztaev A.E., Duisen A.B., Rakhmetova G.T., Agisheva A.A. Prospects for obtaining compounds of hexavalent iron by the method of anodic oxidation // Mat. II M-nar. internet conf. " Contemporary Issues natural and mathematical education”, Aktobe, 2012. - S. 368-372.

3. Kokarovtseva I.G. Oxygen compounds of iron (VI, V, IV) // Advances in Chemistry. - 1972. - T. 41. - S. 1978-1993.

4. Kulikov L.A., Yurchenko A.Yu., Perfilyev Yu.D. Preparation of cesium ferrate (VI) from metallic iron // Vestn. Moscow University Ser. 2 chemistry. - 1999. Volume 40. - No. 2. - S. 137-138.

5. Rylov Yu.B., Dvoretsky S.I. Development of an energy-saving process and hardware-technological design of the production of a regenerative product with potassium ferrate (VI) // Vestnik TSTU. 2012. Volume 18. No. 3. - S. 656-663.

6. Stupin D.Yu. Removal of Ni(II) from aqueous solutions in the presence of EDTA with sodium ferrate (VI) // Journal of Applied Chemistry. - 2004. - No. 8. - S. 1327-1330.

18. Iron, cobalt, nickel

Element Properties

VIII B group.

Properties
Atomic mass
Electronic configuration*


Ionization energy
Relative electronegativity
Possible oxidation states
clarke, at.% (prevalence in nature)		1×10-3	3×10-3
State of aggregation	S E R D E S E S T V A
	silver white	steel gray	silver white


Density
Standard electrode potential

*External configurations shown electronic levels atoms of the corresponding elements. The configurations of the remaining electronic levels coincide with those for noble gases ending the previous period and indicated in brackets.

General characteristics.
The elements iron, cobalt and nickel form the iron triad, or iron family. The atoms of the elements of the iron triad have 2 electrons at the external energy level, which they donate to chemical reactions. However, in education chemical bonds 3d electrons are also involved -orbitals of the second outside level. In their stable compounds, these elements exhibit an oxidation state+2,+3. Form oxides of the composition RO and R 2 O 3 . They correspond to hydroxides of the composition R(OH) 2 and R(OH) 3 .

The elements of the triad (family) of iron are characterized by the property of attaching neutral molecules, for example, carbon monoxide (

II). Carbonyls Ni(CO) 4 , Fe(CO) 5 (liquids at t = 20 ¸60°C) and Co(CO) 8 (crystals with t pl >200 ° Cinsoluble in water and poisonous ) used to produce ultrapure metals.

Cobalt and nickel are less reactive than iron. At ordinary temperatures, they are resistant to corrosion in air, in water and in various solutions. Dilute hydrochloric and sulfuric acids easily dissolve iron and cobalt, and nickel only when heated. Concentrated nitric acid passivates all three metals.

Metals of the iron family, when heated, interact with oxygen, water vapor, halogens, sulfur, phosphorus, silicon, coal and boron . Iron compounds are the most stable

(III) cobalt (II) and nickel (II) - almost all salts are known for them.

Iron, cobalt and nickel are located before hydrogen in the series of standard electrode potentials. Therefore, they are common in nature in the form of compounds (oxides, sulfides, sulfates, carbonates), in a free state they are rare - in the form of iron meteorites. By prevalence in nature

iron is followed by nickel and then cobalt. The compounds of the elements of the iron family in the +2 oxidation state are similar to each other. In the state of higher oxidation states, they exhibit oxidizing properties.

Iron, cobalt, nickel and their alloys are very important materials of modern technology. But iron is the most important.

Iron. Finding in nature. Iron, after aluminum, is the most abundant metal in nature. Its general content earth's crust is 5.1%. Iron is a constituent of many minerals. The most important iron ores are: 1) magnetic iron ore

Fe 3 O 4 , large deposits of this high quality ore are located in the Urals - the mountains High, Blagodat, Magnitnaya; 2) hematite Fe 2 O 3 - the most powerful field - Krivorozhskoe; 3) brown iron ore Fe2O3 ×H2O ; a large deposit - Kerch. Russia is rich in iron ores. Large deposits of them have been discovered in the area of the Kursk magnetic anomaly, on the Kola Peninsula, in Siberia and the Far East.

Often found in nature in large quantities sulfur pyrite (pyrite)

FeS2. It serves as a feedstock for the production of sulfuric acid.

physical properties. Iron is a lustrous, silvery-white metal with a density of 7.87 g/cm

3 , T . sq. 1539 ° C. It has good plasticity. Iron is easily magnetized and demagnetized, and therefore is used as the core of dynamos and electric motors.

Iron consists of four stable isotopes with mass numbers 54.56 (basic), 57 and 58. Radioactive isotopes are used

55 26 Fe and 59 26 F e.

Chemical properties. At the last level, iron atoms have 2 electrons, at the penultimate level - 14, including 6 superoctet ones.

Iron, giving up two outer electrons, exhibits an oxidation state of +2; giving up three electrons (two external and one superoctet from the penultimate energy level), exhibits an oxidation state of +3:

Other oxidation states for iron are not typical.

In air, iron oxidizes easily, especially in the presence of moisture. (rusting):

Interacting with halogens when heated, iron always forms iron (III) halides, for example:

In dilute hydrochloric and sulfuric acids, iron dissolves, i.e., is oxidized by hydrogen ions:

Iron also dissolves in dilute nitric acid, forming an iron salt

(III) water and nitric acid reduction product NH 3 or N 2 O and N 2.

Concentrated acids - oxidizing agents (HNO

3, H2SO4) passivate iron in the cold, but dissolve it when heated:

At high temperature (700-900 ° C), iron reacts with steam

a mi water:

An incandescent iron wire burns brightly in oxygen, forming scale - iron oxide (II,

III):

When heated slightly, iron reacts with chlorine and

sulfur, and at high temperature - with coal, silicon and phosphorus. Iron carbide Fe 3 C is called cementite. It is a gray solid, very brittle and refractory.

With metals and non-metals, iron forms alloys of exceptional importance in the national economy.

For iron, two series of compounds are most characteristic: compounds of iron (II) and iron

(III). iron oxide (
II).Iron(II) oxide FeO -black easily oxidized powder. It turns outreduction of iron oxide(III) carbon monoxide(II) at 500°C:

FeO exhibits the properties of a basic oxide: it readily dissolves in acids, forming iron (II) salts.

iron oxide (

III). Iron oxide (III) Fe 2 O 3- the most stable natural oxygen-containing iron compound. Dissolves in acids to form iron salts(III). Iron oxide (II-III).Iron oxide (II, III) Fe 3 O 4occurs naturally as the mineral magnetite. It is a good conductor of current, therefore it is used for the manufacture of electrodes.

Oxides correspond to iron hydroxides.

Iron hydroxide (

II). Iron(II) hydroxide Fe(OH)2 formed by the action of alkalis on iron (II) salts without air access:

A white precipitate falls out. In the presence of air, the color becomes greenish, and then brown. Iron(II) cations

Fe2+ very easily oxidized by atmospheric oxygen or other oxidizing agents to iron cations(III) Fe 3+ . Therefore, in solutions of iron compounds(II ) there are always iron (III) cations. For the same reason, white iron(II) hydroxide Fe(OH)2 in air, it first becomes greenish, and then brown, turning into iron (III) hydroxide F e (OH) 3: exhibits basic properties, dissolves well in mineral acids, forming salts.

Iron hydroxide (

III). iron hydroxide(III) Fe (OH) 3 formed as a reddish-brown precipitate by the action of alkalis on iron salts(III): e(OH) 3 is a weaker base than iron(II) hydroxide. This is explained by the fact that F e 2+ smaller ion charge and larger radius than Fe 3+ and hence Fe 2+ weaker retains hydroxide ions, i.e. F e (OH) 2 dissociates more easily. Therefore, iron salts(II) slightly hydrolyzed, and iron salts(III) - very strong:

iron hydroxide
(III) has a weakly expressed amphoteric: it dissolves in dilute acids and in concentrated alkali solutions:

Iron salts (
II ) and (III).Of the iron salts, the most widely used are: 1) iron (II) sulfate heptahydrate (ferrous sulfate) FeSO4 × 7 H 2 O for plant pest control, preparation of mineral paints, etc., 2) ferric chloride(III) F eC l 3 as a coagulant in water purification, as well as a mordant in dyeing fabrics; 3) ferrous sulfate nonahydrate(III) Fe 2 (SO 4) 3 ×9 H 2 O as a coagulant, as well as for etching metals;4) iron nitrate nonahydrate ( III) F e (N O 3) 3 ×9 H 2 O as a mordant in the dyeing of cotton fabrics and as a weighting agent for silk.

Qualitative reactions to iron ions (

II ) and (III ). Complex compounds of iron. Iron(III) cation ) is easily detected using a colorless ammonium thiocyanate solution NH4NCS or potassium thiocyanate KNCS, more precisely, the thiocyanate ion NCS - . Under the action of NCS - in a solution of iron salt(III) a blood-red compound is formed -iron thiocyanate(III) F e(NCS) 3:

thiocyanate ion

NCS" serves as a reagent for the iron cation(III) Fe3*.

For iron cation detection

(III) Fe3+ it is convenient to use a complex (complex) iron compound hexacyanoferrate (II) potassium, the so-called. yellow blood salt, K4. In solution, this salt dissociates into ions:

When interacting hexacyanoferrate (

II )-ions 4 - with iron cations(III) Fe3+ a dark blue precipitate is formed - iron hexacyanoferrate (II)(III) ( Prussian blue ):

Another complex iron compound is hexacyanoferrate

(III) potassium (red blood salt) dissociates in solution:

and in the interaction of hexacyanoferrate (

III )-ions 3 - with iron cations ( II) Fe2+ a dark blue precipitate of hexacyanoferrate is also formed(III) iron (II) (turnbull blue):

So the connections

K4 and K3 are important reactants respectively for the iron cation(III) Fe3+ and iron(II) cation Fe2+. Iron ferrates (
VI).A fairly small number of iron (VI) compounds are known - ferrates, for example, potassium ferrate K 2 FeO 4 , barium ferrate BaFeO 4 and calcium ferrate CaFeO if you need fast, detailed and inexpensive
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17. d - elements. Iron, general characteristics, properties. Oxides and hydroxides, CO and OM characteristics, biorole, ability to complex formation.

1. General characteristics.

Iron - d-element of the secondary subgroup of the eighth group of the fourth period of PSCE with atomic number 26.

One of the most common metals in the earth's crust (second place after aluminum).

A simple substance iron is a malleable silver-white metal with a high chemical reactivity: iron quickly corrodes at high temperatures or high humidity in the air.

4Fe + 3O2 + 6H2O = 4Fe(OH)3

In pure oxygen, iron burns, and in a finely dispersed state, it ignites spontaneously in air.

3Fe + 2O2 = FeO + Fe2O3

3Fe + 4H2O = FeO*Fe2O3

FeO*Fe2O3 = Fe3O4 (iron scale)

Actually, iron is usually called its alloys with a low content of impurities (up to 0.8%), which retain the softness and ductility of pure metal. But in practice, alloys of iron with carbon are more often used: steel (up to 2.14 wt.% carbon) and cast iron (more than 2.14 wt.% carbon), as well as stainless (alloyed) steel with the addition of alloying metals (chromium, manganese, nickel, etc.). The combination of the specific properties of iron and its alloys make it "metal No. 1" in importance to humans.

In nature, iron is rarely found in its pure form, most often it occurs as part of iron-nickel meteorites. The prevalence of iron in the earth's crust is 4.65% (4th place after O, Si, Al). It is also believed that iron makes up most of the earth's core.

2.Properties

1.Physical St. Iron is a typical metal, in the free state it is silvery-white in color with a grayish tinge. Pure metal is ductile, various impurities (in particular, carbon) increase its hardness and brittleness. It has pronounced magnetic properties. The so-called "iron triad" is often distinguished - a group of three metals (iron Fe, cobalt Co, nickel Ni) that have similar physical properties, atomic radii and electronegativity values.

2.Chemical St. Islands.

Oxidation state	Oxide	Hydroxide	Character	Notes
			Weakly basic
			Very weak base, sometimes amphoteric
	Not received	*	Acid	Strong oxidizing agent

For iron, the oxidation states of iron are characteristic - +2 and +3.

The oxidation state +2 corresponds to black oxide FeO and green hydroxide Fe(OH) 2 . They are basic. In salts, Fe(+2) is present as a cation. Fe(+2) is a weak reducing agent.

+3 oxidation states correspond to red-brown Fe 2 O 3 oxide and brown Fe(OH) 3 hydroxide. They are amphoteric in nature, although their acidic and basic properties are weakly expressed. So, Fe 3+ ions are completely hydrolyzed even in an acidic environment. Fe (OH) 3 dissolves (and even then not completely), only in concentrated alkalis. Fe 2 O 3 reacts with alkalis only when fused, giving ferrites(formal salts of an acid that does not exist in the free form of acid HFeO 2):

Iron (+3) most often exhibits weak oxidizing properties.

The +2 and +3 oxidation states easily transition between themselves when the redox conditions change.

In addition, there is Fe 3 O 4 oxide, the formal oxidation state of iron in which is +8/3. However, this oxide can also be considered as iron (II) ferrite Fe +2 (Fe +3 O 2) 2 .

There is also an oxidation state of +6. The corresponding oxide and hydroxide do not exist in free form, but salts - ferrates (for example, K 2 FeO 4) have been obtained. Iron (+6) is in them in the form of an anion. Ferrates are strong oxidizing agents.

Pure metallic iron is stable in water and in dilute solutions. alkalis. Iron does not dissolve in cold concentrated sulfuric and nitric acids due to the passivation of the metal surface with a strong oxide film. Hot concentrated sulfuric acid, being a stronger oxidizing agent, interacts with iron.

FROM hydrochloric and diluted (about 20%) sulfuric acids iron reacts to form iron(II) salts:

When iron reacts with approximately 70% sulfuric acid when heated, the reaction proceeds with the formation iron(III) sulfate:

3. Oxides and hydroxides, CO and OM char-ka ...

Iron(II) compounds

Iron oxide (II) FeO has basic properties, it corresponds to the base Fe (OH) 2. Salts of iron (II) have a light green color. When stored, especially in moist air, they turn brown due to oxidation to iron (III). The same process occurs during storage of aqueous solutions of iron(II) salts:

Of iron(II) salts in aqueous solutions, stable mora salt- double ammonium and iron (II) sulfate (NH 4) 2 Fe (SO 4) 2 6H 2 O.

The reagent for Fe 2+ ions in solution can be potassium hexacyanoferrate(III) K 3 (red blood salt). When Fe 2+ and 3− ions interact, a precipitate turnbull blue:

For the quantitative determination of iron (II) in solution, use phenanthroline, which forms a red FePhen 3 complex with iron (II) in a wide pH range (4-9)

Iron(III) compounds

Iron(III) oxide Fe 2 O 3 weakly amphoterene, it corresponds to an even weaker than Fe (OH) 2, base Fe (OH) 3, which reacts with acids:

Fe 3+ salts tend to form crystalline hydrates. In them, the Fe 3+ ion is usually surrounded by six water molecules. Such salts are pink or purple in color. The Fe 3+ ion is completely hydrolyzed even in an acidic environment. At pH>4, this ion is almost completely precipitated in the form of Fe (OH) 3:

With partial hydrolysis of the Fe 3+ ion, polynuclear oxo- and hydroxocations are formed, due to which the solutions become brown. The main properties of iron (III) hydroxide Fe (OH) 3 are very weakly expressed. It is able to react only with concentrated alkali solutions:

The resulting iron(III) hydroxocomplexes are stable only in strongly alkaline solutions. When solutions are diluted with water, they are destroyed, and Fe (OH) 3 precipitates.

When fused with alkalis and oxides of other metals, Fe 2 O 3 forms a variety of ferrites:

Iron(III) compounds in solutions are reduced by metallic iron:

Iron(III) is capable of forming double sulfates with singly charged cations type alum, for example, KFe (SO 4) 2 - potassium iron alum, (NH 4) Fe (SO 4) 2 - iron ammonium alum, etc.

For qualitative detection of iron(III) compounds in a solution, a qualitative reaction of Fe 3+ ions with thiocyanate ions is used SCN − . When Fe 3+ ions interact with SCN − anions, a mixture of bright red iron thiocyanate complexes 2+ , + , Fe(SCN) 3 , - is formed. The composition of the mixture (and hence the intensity of its color) depends on various factors, so this method is not applicable for an accurate qualitative determination of iron.

Another high-quality reagent for Fe 3+ ions is potassium hexacyanoferrate(II) K 4 (yellow blood salt). When Fe 3+ and 4− ions interact, a bright blue precipitate is formed prussian blue:

Iron(VI) compounds

ferrates- salts of iron acid H 2 FeO 4 that do not exist in free form. These are violet-colored compounds, reminiscent of permanganates in oxidizing properties, and sulfates in solubility. Ferrates are obtained by the action of gaseous chlorine or ozone on a suspension of Fe (OH) 3 in alkali , for example, potassium ferrate (VI) K 2 FeO 4 . Ferrates are colored purple.

Ferrates can also be obtained electrolysis 30% alkali solution on an iron anode:

Ferrates are strong oxidizing agents. In an acidic environment, they decompose with the release of oxygen:

The oxidizing properties of ferrates are used to water disinfection.

4.Biorol

1) In living organisms, iron is an important trace element that catalyzes the processes of oxygen exchange (respiration).

2) Iron is usually included in enzymes in the form of a complex. In particular, this complex is present in hemoglobin, the most important protein that provides oxygen transport with blood to all organs of humans and animals. And it is he who stains the blood in a characteristic red color.

4) An excessive dose of iron (200 mg and above) can have a toxic effect. An overdose of iron depresses the antioxidant system of the body, so it is not recommended for healthy people to use iron preparations.

DEVELOPMENT OF AN EFFICIENT METHOD FOR OBTAINING FERRATES (VI) ALKALI METALS

Kaztaev Alimzhan

Tokbanov Bauyrzhan

class 10, AOSCTLIOYU,G.Aktobe, Kazakhstan

Agisheva Almagul Abilkairovna

scientific director,cand. chem. Sciences, Senior Lecturer, Aktobe State Pedagogical Institute, Kazakhstan

Based on the analysis of literature data, taking into account the shortcomings of existing methods for producing ferrates, an electrolytic method for the anodic oxidation of iron was proposed and the operating parameters of the corresponding electrolyzer were refined. The aim of the work was to create a simple method for obtaining compounds of hexavalent iron, reproducible in any industrial and scientific laboratory.

Capacitor, accumulator, transformer plates, metal scrap served as electrodes.

The electrolysis time varied from half an hour to three hours.

The current strength was 2-15 Amperes.

The current density is 0.2-0.4 A/cm 2 and above.

Anode process: Fe 0 + 8 OH - - 6e - = FeO 4 2- + 4H 2 O

Cathodic process: 2H + + 2 e - = H 2

In total, the reaction looks like this: Fe 0 + 2NaOH + 2H 2 O → Na 2 FeO 4 + 2H 2

4Na 2 FeO 4 + 10 H 2 O \u003d 8NaOH + 4Fe (OH) 3 + 3O 2

When concentrated hydrochloric acid is added to the anolyte, chlorine is released, which can be identified by a sharp characteristic odor.

2Na 2 FeO 4 + 16HCl = 3Cl 2 + 4NaCl + 2FeCl 3 + 8H 2 O

The exchange reaction of K 2 FeO 4 with barium chloride gave a pink-brownish precipitate of insoluble barium ferrate. It is more stable and was isolated by filtration and subsequent drying.

Based on the results of several experiments, the masses of the resulting sodium ferrate were calculated.

1) m(BaFeO4) = 0.8013 g

Na 2 FeO 4 - BaFeO 4

166 g/mol - 257 g/mol

X g - 0.8013

X= 0.5175g

m(Na 2 FeO 4) anolyte = 31.05 g

2) m(BaFeO4) = 0.9763 g

Na 2 FeO 4 - BaFeO 4

166 g/mol - 257 g/mol

X g - 0.9763

X= 0.6300g

m(Na 2 FeO 4) anolyte = 37.8 g

3) m(BaFeO4) = 0.8905 g

Na 2 FeO 4 - BaFeO 4

166 g/mol - 257 g/mol

X g - 0.8905

X= 0.5750 g

m(Na 2 FeO 4) anolyte = 34.5 g

According to Faraday's law, the mass of sodium ferrate is:

m= 0.0002867x10800x15=46.450 g

The calculated value differs from that calculated by the analytical method. This is due to the instability of the product and the high degree of its hydrolysis; in addition, unaccounted for electrode processes can occur.

Bibliography:

1. Armoes P., Henze M., Lyakuryansen J., Arvan E. Wastewater treatment // M.: Mir, 2004. - P. 20-60.

2. Kaztaev A.E., Duisen A.B., Rakhmetova G.T., Agisheva A.A. Prospects for obtaining compounds of hexavalent iron by the method of anodic oxidation // Mat. II M-nar. internet conf. "Modern problems of natural and mathematical education", Aktobe, 2012. - P. 368-372.

3. Kokarovtseva I.G. Oxygen compounds of iron (VI, V, IV) // Advances in Chemistry. - 1972. - T. 41. - S. 1978-1993.

4. Kulikov L.A., Yurchenko A.Yu., Perfilyev Yu.D. Preparation of cesium ferrate (VI) from metallic iron // Vestn. Moscow University Ser. 2 chemistry. - 1999. Volume 40. - No. 2. - S. 137-138.

6. Stupin D.Yu. Removal of Ni(II) from aqueous solutions in the presence of EDTA with sodium ferrate (VI) // Journal of Applied Chemistry. - 2004. - No. 8. - S. 1327-1330.

Sodium ferrate properties and contraindications. Development of an effective method for obtaining ferrates (vi) of alkali metals

18. Iron, cobalt, nickel

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