General patterns of changes in the properties of simple substances. Diagonal similarity of elements
Chemistry is a must! how do the oxidizing properties change in the series of elements S---Se---Te---Po? explain the answer. and got the best answer
Answer from Pna Aleksandrovna Tkachenko[active]
In the oxygen subgroup, with increasing atomic number, the radius of atoms increases, and the ionization energy, which characterizes the metallic properties of elements, decreases. Therefore, in the 0--S-Se-Te-Po series, the properties of the elements change from non-metallic to metallic. Under normal conditions, oxygen is a typical non-metal (gas), while polonium is a metal similar to lead.
With an increase in the atomic number of the elements, the value of the electronegativity of the elements in the subgroup decreases. The negative oxidation state is becoming less and less characteristic. The oxidative oxidation state becomes less and less characteristic. The oxidizing activity of simple substances in the series 02--S-Se-Te decreases. So, if sulfur is much weaker, selenium directly interacts with hydrogen, then tellurium does not react with it.
In terms of electronegativity, oxygen is second only to fluorine, therefore, in reactions with all other elements, it exhibits exclusively oxidizing properties. Sulfur, selenium and tellurium in their properties. belong to the group of oxidizing-reducing agents. In reactions with strong reducing agents, they exhibit oxidizing properties, and under the action of strong oxidizing agents. they are oxidized, that is, they exhibit reducing properties.
Possible valencies and oxidation states of the elements of the sixth group of the main subgroup in terms of the structure of the atom.
Oxygen, sulfur, selenium, tellurium and polonium make up the main subgroup of group VI. The outer energy level of the atoms of the elements of this subgroup contains 6 electrons each, which have the s2p4 configuration and are distributed over the cells as follows:
Answer from 2 answers[guru]
Hello! Here is a selection of topics with answers to your question: chemistry, it is very necessary! how do the oxidizing properties change in the series of elements S---Se---Te---Po? explain the answer.
in a series of elements O-S-Se with an increase in the ordinal number of a chemical element, electronegativity 1) increases. 2) smart.
O-S-Se - decreases
C-N-O-F - increases
Fluorine is the most electronegative element.
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AKHMETOV M. A. LESSON 3. ANSWERS TO TASKS.
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Periodic law and the periodic system of chemical elements. Radii of atoms, their periodic changes in the system of chemical elements. Patterns of changes in the chemical properties of elements and their compounds by periods and groups.
1. Arrange the following chemical elements N, Al, Si, C in order of increasing their atomic radii.
ANSWER:
NandClocated in the same period. To the right is locatedN. So nitrogen is less than carbon.
C andSilocated in the same group. But C is higher. So C is less thanSi.
SiandAllocated in one third period, but to the right isSi, meansSiless thanAl
The order of increasing the size of atoms will be as follows:N, C, Si, Al
2. Which of the chemical elements phosphorus or oxygen exhibits more pronounced non-metallic properties? Why?
ANSWER:
Oxygen exhibits more pronounced non-metallic properties, since it is located above and to the right in the periodic table of elements.
3. How do the properties of group IV hydroxides of the main subgroup change when moving from top to bottom?
ANSWER:
The properties of hydroxides change from acidic to basic. SoH2 CO3 - carbonic acid, as its name implies, exhibits acidic properties, andPb(Oh)2 is the base.
ANSWERS TO TESTS
A1. The strength of oxygen-free acids of non-metals of group VIIA, according to the increase in the charge of the nucleus of atoms of elements
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increases |
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decreases |
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does not change |
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changes periodically |
ANSWER: 1
It's about acids.HF, HCl, HBr, HI. In a rowF, Cl, Br, Ian increase in the size of the atoms. Therefore, the internuclear distance increasesH– F, H– Cl, H– Br, H– I. And if so, it means that the bond energy is weakening. And the proton is more easily split off in aqueous solutions
A2. The element has the same valence value in the hydrogen compound and the higher oxide
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germanium |
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ANSWER: 2
Of course, we are talking about an element of the 4th group (see period. c-th elements)
A3. In which order are simple substances arranged in order of increasing metallic properties?
ANSWER: 1
The metallic properties in a group of elements are known to increase from top to bottom.
A4. In the series Na ® Mg ® Al ®Si
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the number of energy levels in atoms increases |
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the metallic properties of the elements are enhanced |
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the highest oxidation state of elements decreases |
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weaken the metallic properties of the elements |
ANSWER: 4
In the period from left to right, non-metallic properties are enhanced, and metallic properties are weakened.
A5. For elements of the carbon subgroup, with increasing atomic number, the
ANSWER: 4.
Electronegativity is the ability to move electrons towards itself when a chemical bond is formed. Electronegativity is almost directly related to non-metallic properties. The non-metallic properties decrease, and the electronegativity also decreases.
A6. In the series of elements: nitrogen - oxygen - fluorine
increases
ANSWER: 3
The number of outer electrons is equal to the group number
A7. Among the chemical elements:
boron - carbon - nitrogen
increases
ANSWER:2
The number of electrons in the outer layer is equal to the highest oxidation state except for (F, O)
A8. Which element has more pronounced non-metallic properties than silicon?
ANSWER: 1
Carbon is in the same group as silicon, only higher.
A9. The chemical elements are arranged in ascending order of their atomic radius in the following order:
ANSWER: 2
In groups of chemical elements, the atomic radius increases from top to bottom.
A10. The metallic properties of the atom are most pronounced:
1) lithium 2) sodium
3) potassium 4) calcium
ANSWER: 3
Among these elements, potassium is located below and to the left.
A11. The most pronounced acidic properties:
Answer: 4 (see answer to A1)
A12. Acid properties of oxides in the series SiO2 ® P2O5 ®SO3
1) weaken
2) intensify
3) do not change
4) change periodically
ANSWER: 2
The acid properties of oxides, as well as non-metallic properties, increase in periods from left to right
A13. With an increase in the charge of the nucleus of atoms, the acidic properties of oxides in the series
N2O5 ® P2O5 ®As2O5 ® Sb2O5
1) weaken
2) intensify
3) do not change
4) change periodically
ANSWER: 1
In groups from top to bottom, acidic properties, like non-metallic ones, weaken
A14. Acidic Properties of Hydrogen Compounds of Group VIA Elements with Increasing Ordinal Number
1) amplify
2) weaken
3) remain unchanged
4) change periodically
ANSWER: 3
The acidic properties of hydrogen compounds are related to the binding energyH- El. This energy from top to bottom weakens, which means that the acidic properties are enhanced.
A15. The ability to donate electrons in the series Na ® K ® Rb ®Cs
1) is weakening
2) amplifies
3) does not change
4) changes periodically
ANSWER: 2
In this series, the number of electron layers and the distance of electrons from the nucleus increase, therefore, the ability to donate an external electron increases.
A16. In the series Al ®Si ®P ®S
1) the number of electron layers in atoms increases
2) non-metallic properties are enhanced
3) the number of protons in the nuclei of atoms decreases
4) the radii of atoms increase
ANSWER: 2
In the period with an increase in the charge of the nucleus, non-metallic properties are enhanced
A17. In the main subgroups of the periodic system, the reducing ability of atoms of chemical elements increases c
ANSWER: 1
With an increase in the number of electronic levels, the remoteness and screening of the outer electrons from the nucleus increases. Consequently, the ability to return them increases (restorative properties)
A18. According to modern ideas, the properties of chemical elements are periodically dependent on
ANSWER: 3
A19. Atoms of chemical elements that have the same number of valence electrons are located
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diagonally |
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in one group |
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in one subgroup |
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in one period |
ANSWER: 2
A20. The element with serial number 114 must have properties similar to
ANSWER: 3. This element will be located in the cell corresponding to the one occupied by lead inVIgroup
A21. In periods, the reducing properties of chemical elements from right to left
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increase |
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decrease |
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do not change |
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change periodically |
ANSWER: 1
The nuclear charge decreases.
A22. Electronegativity and ionization energy in the О–S–Se–Te series, respectively
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increases, increases |
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increases, decreases |
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decreases, decreases |
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decreasing, increasing |
ANSWER: 3
The electronegativity decreases as the number of filled electron layers increases. Ionization energy is the energy required to remove an electron from an atom. She also shrinks
A23. In which order are the signs of chemical elements arranged in order of increasing atomic radii?
Problem 773.
What explains the difference in the properties of the elements of the 2nd period from the properties of their electronic counterparts in subsequent periods?
Solution:
The difference between the properties of the elements of the 2nd period and the properties of their electronic counterparts in subsequent periods is explained
the fact that the atoms of the elements of the 2nd period in the outer electron layer do not contain a d-sublevel. For example, the elements of the main subgroup of group VI: O, S, Se, Te, Po are electronic analogues, since their atoms contain six electrons on the outer electronic layer, two on the s- and four on the p-sublevel. The electronic configuration of their valence layer is: ns2np4. The oxygen atom differs from the atoms of other elements of the subgroup by the absence of a d-sublevel in the outer electron layer:
Such an electronic structure of the oxygen atom does not allow the atom to unpair electrons, therefore covalence oxygen is usually equal to 2 (the number of unpaired valence electrons). Here, an increase in the number of unpaired electrons is possible only by transferring the electron to the next energy level, which, of course, is associated with a large expenditure of energy. Atoms of elements of subsequent periods +16S, +34Se, +52Te and +84Po on the valence electron layer have free d-orbitals:
Such an electronic structure of atoms allows the atoms of these elements to pair electrons, therefore, in an excited state, the number of unpaired electrons increases due to the transfer of s- and p-electrons to free d-orbitals. As a result, these elements are covalence equal not only to 2, but also to 4 and 6:
a) ( covalence – 4) 
b) ( covalence – 4) 
Therefore, unlike the oxygen atom, the atoms of sulfur, selenium and tellurium can participate in the formation of not only two, but also four or six covalent bonds. Atoms of other periods, which also have unoccupied d-orbitals, behave similarly, can go into an excited state and form an additional number of unpaired electrons.
Diagonal similarity of elements
Problem 774.
What is manifested diagonal element similarity? What reasons cause it? Compare the properties of beryllium, magnesium and aluminum.
Solution:
Diagonal similarity - the similarity between the elements located in the Periodic system of elements diagonally from each other, as well as their corresponding simple substances and compounds. The diagonal from the upper left corner to the lower right corner unites somewhat similar elements. This is due to approximately the same increase in non-metallic properties in periods and metallic properties in groups. Diagonal analogy can manifest itself in two forms: the similarity of the general chemical nature of the elements, which is manifested in all compounds of the same type. Diagonal analogy in a broad sense is due to the proximity of the energy and dimensional characteristics of analog elements. In turn, this is determined by a nonmonotonic change, for example, in the electronegativity and orbital radii of the elements horizontally (in a period) and vertically (in a group). Due to this nonmonotonicity, it turns out that the situation is possible when the difference between the characteristics of elements along the diagonal turns out to be smaller than along the horizontal and vertical lines, which leads to a greater chemical similarity of the diagonally located elements in neighboring groups compared to the group analogy. The moeno similarity can be explained by the close charge/radius ratios of the ion.
diagonal similarity observed in pairs of elements Li and Mg, Be and Al, B and Si, etc. This pattern is due to the tendency to change properties vertically (in groups) and their change horizontally (in periods).
It is associated with an increase in non-metallic properties in periods from left to right and in groups from bottom to top. Therefore, lithium is similar to magnesium, beryllium to aluminum, boron to silicon, carbon to phosphorus. Thus, lithium and magnesium form many alkyl and aryl compounds, which are often used in organic chemistry. Beryllium and aluminum have similar redox potentials. Boron and silicon form volatile, highly reactive molecular hydrides.
The chemical properties of beryllium are in many ways similar to those of magnesium (Mg) and, especially, aluminum (Al). The closeness of the properties of beryllium and aluminum is explained by the almost identical ratio of the cation charge to its radius for the Be 2+ and Al 3+ ions. Ve - exhibits, like aluminum, amphoteric properties.
For beryllium and aluminum, the ratio of ion radius to charge, 1/nm, respectively, is 45.4 and 41.7, much higher than for magnesium - 24.4. Magnesium hydroxide has a medium base, while beryllium and aluminum have amphoteric bases. In magnesium, the crystal lattice of chloride is ionic, while in beryllium and aluminum it is molecular (anhydrous); ionic (crystalline hydrate). Magnesium hydride is an ionic compound, and beryllium and aluminum hydrides are polymers.
Physical and chemical properties of simple substances of elements of the main subgroups
Problem 775.
What are the general patterns of change in the physical and chemical properties of simple substances formed by elements of the main subgroups of the periodic system of elements: a) in a period; b) in a group?
Solution:
a) in a period.
In periods (from left to right) - the nuclear charge increases, the number of electronic levels does not change and is equal to the period number, the number of electrons in the outer layer increases, the radius of the atom decreases, the reducing properties decrease, the oxidizing properties increase, the highest oxidation state increases from +1 to +7 , the lowest oxidation state increases from -4 to +1, the metallic properties of substances weaken, non-metallic properties increase. This is due to an increase in the number of electrons in the last layer. In periods from left to right, for higher oxides and their hydrates, the basic properties decrease, while the acid ones increase.
b) in a group.
In the main subgroups (from top to bottom) - the nuclear charge increases, the number of electronic levels increases, the number of electrons on the outer layer does not change and is equal to the group number, the radius of the atom increases, the reducing properties increase, the oxidizing properties decrease, the highest oxidation state is constant and equal to the number groups, the lowest oxidation state does not change and is equal to (- No. of the group), the metallic properties of substances are enhanced, non-metallic properties are weakened .. The formulas of higher oxides (and their hydrates) are common to the elements of the main and secondary subgroups. In higher oxides and their hydrates of elements of groups I - III (except boron), basic properties predominate, from IV to VIII - acidic. In each main subgroup (except VIII), the basic character of oxides and hydroxides increases from top to bottom, while the acidic properties weaken.
This is due to an increase in the number of electron layers and, consequently, to a decrease in the forces of attraction of the electrons of the last layer to the nucleus.
Acid-base properties of oxides and hydroxides of elements
Problem 776.
How do the acid-base and redox properties of higher oxides and hydroxides of elements change with an increase in the charge of their nuclei: a) within a period; b) within the group?
Solution:
a) Within a period, with an increase in the charge of the nuclei of atoms of elements, the acid-base properties of their higher oxides change as follows, the ability to form acids decreases. The change in acid-base properties over the period can be well seen on the example of the following compounds of elements of the third period:
The redox properties over periods with an increase in the charges of the atoms of the elements change as follows, the reducing properties weaken and the oxidizing properties of the elements increase. For example, in the third period, the reducing ability decreases in the sequence: Na 2 O, MgO, Al 2 O 3, SiO 2, P 2 O 5, and the oxidizing ability increases in the sequence: NaOH, Mg (OH) 2, Al (OH ) 3 , H 3 PO 4 , H 2 SO 4 , HClO 4 . The acid-reducing properties of elements depend on the number of oxidation states they exhibit. According to the period, the number of oxidation states shown by the elements naturally increases: Na shows two degrees of oxidation (0 and +1), Cl - seven (0, -1, +1, +3, +4, +5, +6, +7).
b) In groups, with an increase in the charges of the nuclei of the atoms of the elements, the acid-base properties of the oxides and hydroxides of the elements change as follows, the basic properties increase and the acid ones weaken. For example, in groups of electropositive elements, the base strength increases: Be (OH) 2 is an amphoteric compound, and Ba (OH) 2 is a strong base. In groups, with an increase in the charges of atoms of elements, the reducing ability of higher oxides and hydroxides of elements increases, and the oxidizing ability decreases, for example, in the elements of the VIIth group (HClO 4, HBrO 4, HIO 4), the strongest reducing agent is HClO 4, and the weakest is HIO four . In group II (BeO, MgO, CaO, SrO, BaO), the strongest reducing agent is BaO, and the weakest is BeO.
Atoms have 6 electrons in s p orbitals of the outer level. In the O-S-Se-Te-Po series, the ionization energy and electronegativity decrease, the size of atoms and ions increases, the reducing properties increase, and non-metallic features weaken. According to EOTI, oxygen is second only to fluorine. Other elements (-1), (-2) with metals, with non-metals (+4), (+6) In living organisms - O S Se (-2)
Chem. sv.
Oxygen.
4K + O2 > 2K2O
2Sr + O2 > 2SrO
2NO + O2 > 2NO2
CH3CH2OH + 3O2 > 2CO2 + 3H2O
2Na + O2 > Na2O2
2BaO + O2 > 2BaO2
H2 + O2 > H2O2
Na2O2 + O2 > 2NaO2
Selenium is an analogue of sulfur. Just like sulfur, it can be burned in the air. It burns with a blue flame, turning into SeO2 dioxide. Only SeO2 is not a gas, but a crystalline substance, highly soluble in water. Obtaining selenous acid (SeO2 + H2O > H2SeO3) is no more difficult than sulfurous acid. And acting on it with a strong oxidizing agent (for example, HClO3), they get selenic acid H2SeO4, almost as strong as sulfuric acid. Tellurium is less active chemically than sulfur. It is soluble in alkalis, amenable to the action of nitric and sulfuric acids, but slightly soluble in dilute hydrochloric acid. Metallic tellurium begins to react with water at 100°C, and in the form of a powder it oxidizes in air even at room temperature, forming Te02 oxide. When heated in air, tellurium burns to form Te02. This strong compound is less volatile than tellurium itself. Therefore, to purify tellurium from oxides, they are reduced by running hydrogen at 500-600 °C. In the molten state, tellurium is rather inert; therefore, graphite and quartz are used as container materials for its melting.
Polonium metal oxidizes rapidly in air. Polonium dioxide (PoO2)x and polonium monoxide PoO are known. Forms tetrahalides with halogens. Under the action of acids, it goes into solution with the formation of pink Po2 + cations:
Po + 2HCl > PoCl2 + H2^.
When polonium is dissolved in hydrochloric acid in the presence of magnesium, hydrogen polonium is formed:
Po + Mg + 2HCl > MgCl2 + H2Po,
9. Oxygen- the most common element on Earth, its share (in the composition of various compounds, mainly silicates), accounts for about 47.4% of the mass of the solid earth's crust. Sea and fresh waters contain a huge amount of bound oxygen - 88.8% (by mass), in the atmosphere the content of free oxygen is 20.95% by volume and 23.12% by mass. More than 1500 compounds of the earth's crust contain oxygen in their composition. Oxygen is a constituent of many organic substances and is present in all living cells. By the number of atoms in living cells, it is about 25%, by mass fraction - about 65%. Oxygen is a chemically active non-metal, it is the lightest element from the chalcogen group. The simple substance oxygen (CAS number: 7782-44-7) under normal conditions is a colorless, tasteless and odorless gas, the molecule of which consists of two oxygen atoms (formula O2), and therefore it is also called dioxygen. Liquid oxygen has a light blue color. At present, in industry, oxygen is obtained from the air. In laboratories, industrial oxygen is used, supplied in steel cylinders under a pressure of about 15 MPa. The most important laboratory method for its production is the electrolysis of aqueous solutions of alkalis. Small amounts of oxygen can also be obtained by reacting a potassium permanganate solution with an acidified hydrogen peroxide solution. Oxygen plants based on membrane and nitrogen technologies are also well known and successfully used in industry. When heated, potassium permanganate KMnO4 decomposes to potassium manganate K2MnO4 and manganese dioxide MnO2 with the simultaneous release of gaseous oxygen O2:
2KMnO4 > K2MnO4 + MnO2 + O2^
Under laboratory conditions, it is also obtained by catalytic decomposition of hydrogen peroxide H2O2:
2H2O2 > 2H2O + O2^
The catalyst is manganese dioxide (MnO2) or a piece of raw vegetables (they contain enzymes that accelerate the decomposition of hydrogen peroxide). Oxygen can also be obtained by catalytic decomposition of potassium chlorate (bertolet salt) KClO3:
2KClO3 > 2KCl + 3O2^
MnO2 also acts as a catalyst
Physical properties of oxygen
Under normal conditions, oxygen is a colorless, tasteless and odorless gas. 1 liter of it weighs 1.429 g. A little heavier than air. Slightly soluble in water (4.9 ml/100g at 0°C, 2.09 ml/100g at 50°C) and alcohol (2.78 ml/100g). It dissolves well in molten silver (22 volumes of O2 in 1 volume of Ag at 961 °C). It is paramagnetic. When gaseous oxygen is heated, its reversible dissociation into atoms occurs: at 2000 °C - 0.03%, at 2600 °C - 1%, 4000 °C - 59%, 6000 °C - 99.5%. Liquid oxygen (boiling point? 182.98 °C) is a pale blue liquid. Solid oxygen (melting point? 218.79 ° C) - blue crystals
Chem. saints
A strong oxidizing agent, interacts with almost all elements, forming oxides. Oxidation state?2. As a rule, the oxidation reaction proceeds with the release of heat and accelerates with increasing temperature. An example of reactions occurring at room temperature:
4K + O2 > 2K2O
Oxidizes compounds that contain elements with a non-maximum oxidation state:
2NO + O2 > 2NO2
Oxidizes most organic compounds:
CH3CH2OH + 3O2 > 2CO2 + 3H2O
Under certain conditions, it is possible to carry out a mild oxidation of an organic compound:
CH3CH2OH + O2 > CH3COOH + H2O
Oxygen does not oxidize Au and Pt, halogens and inert gases.
Oxygen forms peroxides with an oxidation state of ?1. For example, peroxides are obtained by burning alkali metals in oxygen:
2Na + O2 > Na2O2
Some oxides absorb oxygen:
2BaO + O2 > 2BaO2
According to the combustion theory developed by A. N. Bach and K. O. Engler, oxidation occurs in two stages with the formation of an intermediate peroxide compound. This intermediate compound can be isolated, for example, when the flame of burning hydrogen is cooled with ice, along with water, hydrogen peroxide is formed:
H2 + O2 > H2O2
Superoxides have an oxidation state of ?1/2, that is, one electron per two oxygen atoms (O2 - ion). Obtained by the interaction of peroxides with oxygen at elevated pressures and temperatures:
Na2O2 + O2 > 2NaO2
KOH(solid) + O3 > KO3 + KOH + O2
The dioxygenyl O2+ ion has an oxidation state of +1/2. Obtained by the reaction: PtF6 + O2 > O2PtF6
Oxygen fluorides
Oxygen difluoride, OF2 oxidation state +2, is obtained by passing fluorine through an alkali solution:
2F2 + 2NaOH > OF2 + 2NaF + H2O
Oxygen monofluoride (Dioxydifluoride), O2F2, is unstable, oxidation state is +1. Obtained from a mixture of fluorine and oxygen in a glow discharge at a temperature of? 196 ° C. Passing a glow discharge through a mixture of fluorine with oxygen at a certain pressure and temperature, mixtures of higher oxygen fluorides O3F2, O4F2, O5F2 and O6F2 are obtained. Oxygen supports the processes of respiration, combustion, and decay. In its free form, the element exists in two allotropic modifications: O2 and O3 (ozone). Ozone is formed in many processes accompanied by the release of atomic oxygen, for example, during the decomposition of peroxides, the oxidation of phosphorus, etc. In industry, it is obtained from air or oxygen in ozonizers the action of an electrical discharge. O3 liquefies more easily than O2 and is therefore easy to separate. Ozone for ozone therapy in medicine is obtained only from pure oxygen. When air is irradiated with hard ultraviolet radiation, ozone is formed. The same process takes place in the upper layers of the atmosphere, where the ozone layer is formed and maintained under the influence of solar radiation.
Physical properties of ozone
Molecular weight - 47.998 amu
The density of the gas under normal conditions is 1.1445 kg/m3. Relative density of gas for oxygen 1.5; by air - 1.62 (1.658).
Liquid density at -183 °C - 1.71 kg/m3
Boiling point -111.9 °C. Liquid ozone is dark blue.
Melting point -251.4 °C. In the solid state - black-blue.
Solubility in water at 0oC - 0.394 kg/m3 (0.494 l/kg), it is 10 times higher compared to oxygen.
In the gaseous state, ozone is diamagnetic; in the liquid state, it is weakly paramagnetic.
The smell is sharp, specific "metallic" (according to Mendeleev - "the smell of crayfish").
Chemical St. Ozone.
Ozone is a powerful oxidizing agent, much more reactive than diatomic oxygen. Oxidizes almost all metals (with the exception of gold, platinum and iridium) to their highest oxidation states. Oxidizes many non-metals.
2 Cu2+(aq) + 2 H3O+(aq) + O3(g) > 2 Cu3+(aq) + 3 H2O(l) + O2(g)
Ozone increases the oxidation state of oxides:
NO + O3 > NO2 + O2
The formation of ozone proceeds by a reversible reaction:
3O2 + 68 kcal (285 kJ)<>2O3.
salt-forming oxides:
basic oxides (for example, sodium oxide Na2O, copper (II) oxide CuO): metal oxides, the oxidation state of which is I-II;
acidic oxides (for example, sulfur(VI) oxide SO3, nitric oxide(IV) NO2): metal oxides with oxidation state V-VII and non-metal oxides;
amphoteric oxides (for example, zinc oxide ZnO, aluminum oxide Al2O3): metal oxides with oxidation states III-IV and exceptions (ZnO, BeO, SnO, PbO);
Non-salt-forming oxides: carbon monoxide (II) CO, nitric oxide (I) N2O, nitric oxide (II) NO, silicon oxide (II) SiO.
Chem. sv-va osn oks
1. Basic oxide + acid \u003d salt + water
CuO + H2SO4 = CuSO4 + H2O (orthophosphoric or strong acid).
2. Strong base oxide + water = alkali
CaO + H2O = Ca(OH)2
3. Strong base oxide + acid oxide = salt
CaO + Mn2O7 = Ca(MnO4)2
Na2O + CO2 = Na2CO3
4. Basic oxide + hydrogen = metal + water
CuO + H2 = Cu + H2O (Note: the metal is less active than aluminium).
Chem. St. sour ox
1. Acid oxide + water = acid
SO3 + H2O = H2SO4
Some oxides, such as SiO2, do not react with water, so their acids are obtained indirectly.
2. Acid oxide + basic oxide = salt
CO2 + CaO = CaCO3
3. Acid oxide + base = salt + water
SO2 + 2NaOH = Na2SO3 + H2O
If the acid oxide is an anhydride of a polybasic acid, the formation of acid or medium salts is possible:
Ca(OH)2 + CO2 = CaCO3v + H2O
CaCO3 + CO2 + H2O = Ca(HCO3)2
4. Non-volatile oxide + salt1 = salt2 + volatile oxide
SiO2 + Na2CO3 = Na2SiO3 + CO2^
10. Water (hydrogen oxide)- a transparent liquid, colorless (in a small volume) and odor. Chemical formula: H2O. In the solid state it is called ice or snow, and in the gaseous state it is called water vapor. About 71% of the Earth's surface is covered with water (oceans, seas, lakes, rivers, ice at the poles). It is a good highly polar solvent. Under natural conditions, it always contains dissolved substances (salts, gases). Water is of key importance in the creation and maintenance of life on Earth, in the chemical structure of living organisms, in the formation of climate and weather. Water has a number of unusual features: When ice melts, its density increases (from 0.9 to 1 g/cm?). For almost all other substances, the density decreases when melted. When heated from 0 °C to 4 °C (more precisely, 3.98 °C), water contracts. Thanks to this, fish can live in freezing water bodies: when the temperature drops below 4 ° C, colder water, as less dense, remains on the surface and freezes, and a positive temperature remains under the ice. High temperature and specific heat of fusion (0 °C and 333.55 kJ/kg), boiling point (100 °C) and specific heat of vaporization (2250 kJ/kg), compared to hydrogen compounds of similar molecular weight. High heat capacity of liquid water. High viscosity. High surface tension. Negative electric potential of the water surface. According to the state, they distinguish:
Solid - ice
Liquid - water
Gaseous - water vapor. Both oxygen and hydrogen have natural and artificial isotopes. Depending on the type of isotopes included in the molecule, the following types of water are distinguished: Light water (just water), Heavy water (deuterium) and Super heavy water (tritium). Water is the most common solvent on Earth, largely determining the nature of terrestrial chemistry as a science. Most of chemistry, at its inception as a science, began precisely as the chemistry of aqueous solutions of substances. It is sometimes considered as an ampholyte - both an acid and a base at the same time (cation H + anion OH-). In the absence of foreign substances in water, the concentrations of hydroxide ions and hydrogen ions (or hydronium ions) are the same, pKa ? OK. 16. Water itself is relatively inert under normal conditions, but its highly polar molecules solvate ions and molecules, form hydrates and crystalline hydrates. Solvolysis, and in particular hydrolysis, occurs in animate and inanimate nature, and is widely used in the chemical industry. Aquacomplexes, coordination Comm., containing as a ligand-ve one or more. water molecules. The latter is connected to the center, a metal atom, through an oxygen atom. Distinguish A. cationic type (eg, [Co (H2O) 6] C12), anionic (eg, K [Cr (H2O) 2 (OH) 4]) and non-electrolyte complexes (eg, ).A. in many cases are easily formed in aqueous solutions from other coordinates. conn. as a result of intrasphere substitution, hydration of cations, and addition of H2O molecules. In the latter case, the coordination center number. atom can increase, for example. as a result of the addition of two water molecules to the anions [AuC14] - or - two molecules of water. Thus, the time of almost complete isotopic exchange of H2O to 18H2O in [A1(H2O)6]3+, 3+, etc. at 25°C is approx. 1 minute. For stable A., for example. [Cr (H2O) 6] C13, half-life during isotopic exchange - approx. 40 h at 25°C.A. have acid properties, e.g. -5.69, for 4+ -4.00. A hydrogen bond is an intermolecular bond formed due to the partial acceptance of a lone pair of electrons of an atom by a hydrogen atom that is not chemically bonded to it. Autoprotolysis is a reversible process of formation of an equal number of cations and anions from uncharged molecules of a liquid individual substance due to the transfer of a proton from one molecule to another. Due to thermal vibrations, a hydrogen atom forming a hydrogen bond can momentarily occupy an intermediate position between oxygen atoms. From a particle with such a hydrogen atom, both the initial water molecules bound by hydrogen bonds and two ions can be formed with equal probability: a hydroxide ion and an oxonium ion. That is, the reaction 2H2O = H3O + OH proceeds in water.
The reverse process also easily occurs - the formation of two water molecules in the collision of an oxonium ion with a hydroxide ion: H3O + OH \u003d 2H2O.
Both of these reactions proceed in water constantly and at the same rate, therefore, there is an equilibrium in water: 2H2O AH3O + OH. This equilibrium is called the water autoprotolysis equilibrium.
11. Peroxide(previously - peroxide) - a substance containing a peroxo group -O-O- (for example, hydrogen peroxide H2O2, sodium peroxide Na2O2). Peroxide readily releases oxygen. For inorganic substances, it is recommended to use the term peroxide; for organic substances, the term peroxide is often used in Russian today. Peroxides of many organic substances are explosive (acetone peroxide); in particular, they are easily formed photochemically when ethers are illuminated for a long time in the presence of oxygen. Therefore, before distillation, many ethers (diethyl ether, tetrahydrofuran) require testing for the absence of peroxides. Peroxides slow down protein synthesis in the cell.
Hydrogen peroxide
In nature, it is formed as a by-product during the oxidation of many substances with atmospheric oxygen. Traces of it are constantly found in atmospheric precipitation. Hydrogen peroxide is also partially formed in the flame of burning hydrogen, but decomposes when the combustion products cool. In fairly high concentrations (up to several percent), H2O2 can be obtained by the interaction of hydrogen at the time of release with molecular oxygen. Hydrogen peroxide is also partially formed when moist oxygen is heated to 2000 ° C, when a quiet electric discharge passes through a wet mixture of hydrogen and oxygen, and when water is exposed to ultraviolet rays or ozone. Hydrogen peroxide is easiest to obtain from barium peroxide (BaO2), acting on it with dilute sulfuric acid:
BaO2 + H2SO4 = BaSO4 + H2O2.
In this case, along with hydrogen peroxide, water-insoluble barium sulfate is formed, from which the liquid can be separated by filtration. H2O2 is usually sold in the form of a 3% aqueous solution. The main method for producing hydrogen peroxide is the interaction of persulfuric acid (or some of its salts) with water, which easily proceeds according to the scheme:
H2S2O8 + 2 H2O = 2 H2SO4 + H2O2.
Of lesser importance are some new methods (decomposition of organic peroxide compounds, etc.) and the old method of obtaining from BaO2. For storage and transportation of large quantities of hydrogen peroxide, aluminum containers (not lower than 99.6% purity) are most suitable. Pure hydrogen peroxide is a colorless syrupy liquid (with a density of about 1.5 g / ml), distilled under sufficiently reduced pressure without decomposition. The freezing of H2O2 is accompanied by compression (unlike water). White crystals of hydrogen peroxide melt at -0.5 ° C, i.e., almost at the same temperature as ice. The heat of fusion of hydrogen peroxide is 13 kJ/mol, the heat of vaporization is 50 kJ/mol (at 25 °C). Under ordinary pressure, pure H2O2 boils at 152°C with strong decomposition (and the vapors can be explosive). For its critical temperature and pressure, the theoretically calculated values are 458 °C and 214 atm. The density of pure H2O2 is 1.71 g/cm3 in the solid state, 1.47 g/cm3 at 0°C, and 1.44 g/cm3 at 25°C. Liquid hydrogen peroxide, like water, is highly associated. The refractive index of H2O2 (1.41), as well as its viscosity and surface tension, are slightly higher than those of water (at the same temperature). Hydrogen peroxide is a strong oxidizing agent, that is, it easily gives up its extra (compared to the more stable compound - water) oxygen atom. Thus, under the action of anhydrous and even highly concentrated H2O2 on paper, sawdust and other combustible substances, they ignite. The practical use of hydrogen peroxide is based mainly on its oxidizing effect. The annual world production of H2O2 exceeds 100 thousand tons. The oxidative decomposition characteristic of hydrogen peroxide can be schematically depicted as follows:
H2O2 \u003d H2O + O (for oxidation).
An acidic environment is more conducive to this disintegration than an alkaline one. Much less typical for hydrogen peroxide is the reductive decomposition according to the scheme:
H2O2 \u003d O2 + 2 H (for recovery)
An alkaline environment is more conducive to such disintegration than an acidic one. The reductive decomposition of hydrogen peroxide takes place, for example, in the presence of silver oxide:
Ag2O + H2O2 = 2 Ag + H2O + O2.
Similarly, in essence, its interaction with ozone (O3 + H2O2 = 2 H2O + 2 O2) and with potassium permanganate in an acidic medium proceeds:
2 KMnO4 + 5 H2O2 + 3 H2SO4 = K2SO4 + 2 MnSO4 + 5 O2 + 8 H2O.
More than half of all hydrogen peroxide produced is spent on bleaching various materials, usually carried out in very dilute (0.1-1%) aqueous solutions of H2O2. An important advantage of hydrogen peroxide over other oxidizing agents lies in the "softness" of action, due to which the bleached material itself is almost not affected. Related to this is the medical use of very dilute hydrogen peroxide as an antiseptic (for gargling, etc.). Very concentrated (80% and above) aqueous solutions of H2O2 are used as energy sources.
12. Sulfur - highly electronegative element, exhibits non-metallic properties. In hydrogen and oxygen compounds, it is part of various ions, forms many acids and salts. Many sulfur-containing salts are sparingly soluble in water. The most important natural sulfur compounds FeS2 are iron pyrite or pyrite, ZnS is zinc blende or sphalerite (wurtzite), PbS is lead gloss or galena, HgS is cinnabar, Sb2S3 is antimonite. In addition, sulfur is present in oil, natural coal, natural gases and shale. Sulfur is the sixth element in natural waters, occurs mainly in the form of sulfate ion and causes the "permanent" hardness of fresh water. A vital element for higher organisms, an integral part of many proteins, is concentrated in the hair. Sulfur is obtained mainly by smelting native sulfur directly in places where it occurs underground. Sulfur ores are mined in different ways - depending on the conditions of occurrence. Sulfur deposits are almost always accompanied by accumulations of poisonous gases - sulfur compounds. In addition, we must not forget about the possibility of its spontaneous combustion. Ore mining in an open way is as follows. Walking excavators remove layers of rocks under which ore lies. Explosions crush the ore layer, after which the blocks of ore are sent to a sulfur smelter, where sulfur is extracted from the concentrate. Sulfur is quite widespread in nature. In the earth's crust, its content is estimated at 0.05% by weight. Significant deposits of native sulfur are often found in nature (usually near volcanoes); In 1890, Herman Frasch proposed to melt sulfur underground and pump it to the surface through wells, similar to oil wells. The relatively low (113°C) melting point of sulfur confirmed the reality of Frasch's idea. There are several methods for obtaining sulfur from sulfur ores: steam-water, filtration, thermal, centrifugal and extraction. Sulfur is also found in large quantities in natural gas in the gaseous state (in the form of hydrogen sulfide, sulfur dioxide). During extraction, it is deposited on the walls of pipes and equipment, disabling them. Therefore, it is captured from the gas as soon as possible after extraction. The resulting chemically pure fine sulfur is an ideal raw material for the chemical and rubber industries. Sulfur differs significantly from oxygen in its ability to form stable chains and cycles from sulfur atoms. The most stable are cyclic S8 molecules, which have the shape of a crown and form rhombic and monoclinic sulfur. This is crystalline sulfur - a brittle yellow substance. In addition, molecules with closed (S4, S6) chains and open chains are possible. Such a composition has plastic sulfur, a brown substance. The formula for plastic sulfur is most often written simply as S, since, although it has a molecular structure, it is a mixture of simple substances with different molecules. Sulfur is insoluble in water, some of its modifications are dissolved in organic solvents, such as carbon disulfide. Sulfur forms several tens of both crystalline and amorphous modifications. At normal pressure and temperatures up to 98.38 ° C, the a-modification of sulfur is stable (otherwise this modification is called rhombic), which forms lemon-yellow crystals. Above 95.39 ° C, the b-modification of sulfur (the so-called monoclinic sulfur) is stable. With prolonged exposure at temperatures of 20-95 ° C, all sulfur modifications turn into a-sulfur. The melting point of rhombic a-sulfur is 112.8 ° C, and monoclinic b-sulfur 119.3°C. In both cases, an easily mobile yellow liquid is formed, which darkens at a temperature of about 160 ° C; its viscosity increases, and at temperatures above 200 ° C, the molten sulfur becomes dark brown and viscous, like a resin. This is explained by the fact that the S8 ring molecules are first destroyed in the melt. The resulting fragments combine with each other to form long chains S of several hundred thousand atoms. Further heating of the molten sulfur (above a temperature of 250°C) leads to a partial break in the chains, and the liquid again becomes more mobile. At about 190°C, its viscosity is about 9000 times greater than at 160°C. At a temperature of 444.6°C, molten sulfur boils. Sulfur is used for the production of sulfuric acid, rubber vulcanization, as a fungicide in agriculture and as colloidal sulfur - a drug. Also, sulfur in the composition of sulfur-bitumen compositions is used to obtain sulfur asphalt, and as a substitute for Portland cement - to obtain sulfur concrete. Sulfur is practically insoluble in water. Some of its modifications dissolve in organic liquids (toluene, benzene) and especially well in carbon disulfide CS2 and liquid ammonia NH3. At room temperature, sulfur reacts with fluorine and chlorine, exhibiting reducing properties:
Sulfur reacts with concentrated oxidizing acids (HNO3, H2SO4) only during prolonged heating, oxidizing:
S + 6HNO3(conc.) = H2SO4 + 6NO2 ^ + 2H2O
S + 2H2SO4 (conc.) = 3SO2 ^ + 2H2O
In air, sulfur burns, forming sulfur dioxide - a colorless gas with a pungent odor:
Using spectral analysis, it was found that in fact the process of oxidation of sulfur to dioxide is a chain reaction and occurs with the formation of a number of intermediate products: sulfur monoxide S2O2, molecular sulfur S2, free sulfur atoms S and free radicals of sulfur monoxide SO. When interacting with metals, it forms sulfides. 2Na + S = Na2S
When sulfur is added to these sulfides, polysulfides are formed: Na2S + S = Na2S2
When heated, sulfur reacts with carbon, silicon, phosphorus, hydrogen:
C + 2S = CS2 (carbon disulfide)
Sulfur dissolves in alkalis when heated - disproportionation reaction
3S + 6KOH = K2SO3 + 2K2S + 3H2O
Finely ground sulfur is prone to chemical spontaneous combustion in the presence of moisture, in contact with oxidizing agents, and also in a mixture with coal, fats, oils. Sulfur forms explosive mixtures with nitrates, chlorates and perchlorates. It ignites spontaneously on contact with bleach. About half of the produced sulfur is used for the production of sulfuric acid, about 25% is used to produce sulfites, 10-15% is used to control pests of agricultural crops (mainly grapes and cotton) (the most important solution here is copper sulphate CuSO4 5H2O), about 10% is used by the rubber industry for rubber vulcanization. Sulfur is used in the production of dyes and pigments, explosives (it is still part of gunpowder), artificial fibers,
phosphors. Sulfur is used in the manufacture of matches, as it is part of the composition from which the heads of matches are made. Sulfur is still contained in some ointments that treat skin diseases.
13. SO2 (sulphurous anhydride; sulfur dioxide)
Physical Properties
Colorless gas with a pungent odor; very soluble in water (40V SO2 dissolves in 1V H2O at N.O.); t°pl. = -75.5°C; t°boiling = -10°С. Discolors many dyes, kills microorganisms.
Receipt
When burning sulfur in oxygen: S + O2 ® SO2
Sulfide oxidation: 4FeS2 + 11O2 ® 2Fe2O3 + 8SO2
Treatment of salts of sulfurous acid with mineral acids:
Na2SO3 + 2HCl ® 2NaCl + SO2+ H2O
When metals are oxidized with concentrated sulfuric acid:
Cu + 2H2SO4(conc) ® CuSO4 + SO2+ 2H2O
Chemical properties
Sulfur dioxide is an acidic oxide. When dissolved in water, weak and unstable sulfurous acid H2SO3 is formed (exists only in aqueous solution) SO2 + H2O « H2SO3 K1® H+ + HSO3- K2® 2H+ + SO32- H2SO3 forms two series of salts - medium (sulfites) and acidic (bisulfites, hydrosulfites).
Ba(OH)2 + SO2 ® BaSO3?(barium sulfite) + H2OBa(OH)2 + 2SO2 ® Ba(HSO3)2(barium hydrosulfite)
Oxidation reactions (S+4 – 2e ® S+6)SO2 + Br2 + 2H2O ® H2SO4 + 2HBr
5SO2 + 2KMnO4 + 2H2O ® K2SO4 + 2MnSO4 + 2H2SO4
Aqueous solutions of alkali metal sulfites are oxidized in air:
2Na2SO3 + O2 ® 2Na2SO4; 2SO32- + O2 ® 2SO42-
Reduction reactions (S+4 + 4e ® S0)SO2 + С –t°® S + СO2
SO2 + 2H2S ® 3S + 2H2O
Sulfur oxide VI SO3 (sulfuric anhydride)
Physical Properties
Colorless volatile liquid, t°pl. = 17°C; t°boiling = 66°С; "smokes" in air, strongly absorbs moisture (stored in sealed vessels). SO3 + H2O ® H2SO4 Solid SO3 exists in three modifications. SO3 dissolves well in 100% sulfuric acid, this solution is called oleum.
Receipt
1)2SO2 + O2 cat;450°C® 2SO32) Fe2(SO4)3 –t°® Fe2O3 + 3SO3
Chemical properties
Sulfuric anhydride is an acidic oxide. When dissolved in water, it gives a strong dibasic sulfuric acid:
SO3 + H2O ® H2SO4 « H+ + HSO4- « 2H+ + SO42-H2SO4 forms two series of salts - medium (sulfates) and acidic (hydrosulfates): 2NaOH + SO3 ® Na2SO4 + H2O
NaOH + SO3 ® NaHSO4SO3 is a strong oxidizing agent.
H2SO4 is a strong dibasic acid, corresponding to the highest oxidation state of sulfur (+6). Under normal conditions, concentrated sulfuric acid is a heavy oily liquid, colorless and odorless. Sulfuric acid is a fairly strong oxidizing agent, especially when heated and in concentrated form; oxidizes HI and partially HBr to free halogens, carbon to CO2, S to SO2, oxidizes many metals (Cu, Hg, etc.). In this case, sulfuric acid is reduced to SO?, and the strongest reducing agents are reduced to S and H?S. Concentrated H?SO? H? is partially restored. Because of what it cannot be used for drying it. Diluted H?SO? interacts with all metals that are in the electrochemical series of voltages to the left of hydrogen with its release. Oxidizing properties for dilute H?SO? uncharacteristic. Sulfuric acid forms two series of salts: medium - sulfates and acidic - hydrosulfates, as well as esters. Peroxomonosulfuric (or Caro's acid) H2SO5 and peroxodisulfuric H2S2O8 acids are known. H2SO3 is an unstable dibasic acid of medium strength, exists only in dilute aqueous solutions (not isolated in a free state):
SO2 + H2O ? H2SO3? H+ + HSO3- ? 2H+ + SO32-.
Medium strength acid:
H2SO3<=>H+ + HSO3-, KI = 2 10-2
HSO3-<=>H+ + SO32-, KII = 6 10-8
Solutions of H2SO3 always have a sharp specific odor (similar to the smell of a lighted match) due to the presence of SO2 not chemically bound by water. Dibasic acid, forms two series of salts: acidic - hydrosulfites (in the absence of alkali):
H2SO3 + NaOH = NaHSO3 + H2O
and medium - sulfites (in excess of alkali): H2SO3 + 2NaOH = Na2SO3 + 2H2O
Like sulfur dioxide, sulfurous acid and its salts are strong reducing agents:
H2SO3+Br2+H2O=H2SO4+2HBr
When interacting with even stronger reducing agents, it can play the role of an oxidizing agent:
H2SO3+2H2S=3S+3H2O
Qualitative reaction to sulfite ions - discoloration of a solution of potassium permanganate:
5SO3 + 6H+2MnO4=5SO4+2Mn+3H2O
Sulfites are salts of sulfurous acid H2SO3. There are two series of sulfites: medium (normal) of the general formula M2SO3 and acidic (hydrosulfites) of the general formula MHSO3 (M is a monovalent metal). Averages, with the exception of alkali metal and ammonium sulfites, are sparingly soluble in water and dissolve in the presence of SO2. Of the acidic compounds in the free state, only hydrosulfites of alkali metals have been isolated. Sulfites in aqueous solution are characterized by oxidation to sulfates and reduction to thiosulfates M2S2O3. Reactions with an increase in the oxidation state of sulfur from +4 to +6, for example:
Na2SO3 + Cl2 + H2O = Na2SO4 + 2 HCl.
Self-oxidation-self-healing reactions of sulfur are also possible when it interacts with sulfites. So, when boiling a solution with finely ground sulfur, sodium thiosulfate (sometimes called hyposulfite) is formed:
Na2SO3 + S > Na2S2O3.
Thus, sulfurous acid and its salts can exhibit both oxidizing and reducing properties. It is obtained by the interaction of SO2 with hydroxides or carbonates of the corresponding metals in an aqueous medium. Hydrosulfites are mainly used - in the textile industry for dyeing and printing (KHSO3, NaHSO3), in the paper industry for the production of cellulose from wood, in photography, in organic synthesis. Sulphates - sulfate salts, salts of sulfuric acid H2SO4. There are two rows of S. - medium (normal) of the general formula Mg2SO4 and acidic (Hydrosulfates) - MHSO4, where M is a monovalent metal. S. - crystalline substances, colorless (if the cation is colorless), in most cases highly soluble in water. Slightly soluble S. are found in the form of minerals: gypsum CaSO4?2H2O, celestite SrSO4, anglesite PbSO4, and others. Barite BaSO4 and RaSO4 are practically insoluble. Acidic sulfurs have been isolated in the solid state only for the most active metals—Na, K, and others. They are readily soluble in water and readily melt. Normal S. can be obtained by dissolving metals in H2SO4, by the action of H2SO4 on metal oxides, hydroxides, carbonates, etc. Hydrosulfates are obtained by heating normal S. with concentrated H2SO4:
K2SO4 + H2SO4 = 2KHSO4.
The crystalline hydrates of certain heavy metals are called vitriol. Natural sulfates are widely used in many industries.
14. H2S - colorless gas with an unpleasant odor and a sweetish taste. Let's badly dissolve in water, it is good - in ethanol. At high concentrations, it corrodes metal. Explosive mixture with air 4.5 - 45%. Thermally unstable (at temperatures above 400 ° C it decomposes into simple substances - S and H2), poisonous (inhalation of air with its admixture causes dizziness, headache, nausea, and with a significant content leads to coma, convulsions, pulmonary edema and even death exodus), a gas heavier than air with an unpleasant smell of rotten eggs. The hydrogen sulfide molecule has an angular shape, so it is polar (? = 0.34 10-29 C m). Unlike water molecules, hydrogen sulfide molecules do not form strong hydrogen bonds, so H2S is a gas. Saturated aqueous solution (hydrosulfide water) H2S is a very weak hydrosulfide acid. The intrinsic ionization of liquid hydrogen sulfide is negligible. Hydrogen sulfide is sparingly soluble in water, an aqueous solution of H2S is a very weak acid:
Reacts with bases:
H2S + 2NaOH = Na2S + 2H2O (ordinary salt, with an excess of NaOH)
H2S + NaOH = NaHS + H2O (acid salt, at a ratio of 1:1)
Hydrogen sulfide is a strong reducing agent. In air, it burns with a blue flame:
2H2S + 3O2 = 2H2O + 2SO2
with a lack of oxygen: 2H2S + O2 = 2S + 2H2O
(The industrial method for producing sulfur is based on this reaction). Hydrogen sulfide also reacts with many other oxidizing agents; when it is oxidized in solutions, free sulfur or SO42- is formed, for example:
3H2S + 4HClO3 = 3H2SO4 + 4HCl
2H2S + SO2 = 2H2O + 3S
H2S + I2 = 2HI + S
Receipt
Reaction of dilute acids on sulfides: FeS + 2HCl = FeCl2 + H2S
The interaction of aluminum sulfide with water (this reaction is distinguished by the purity of the resulting hydrogen sulfide): Al2SO3 + H2O \u003d 2Al (OH) 3 + H2S
Salts of hydrosulphuric acid are called sulfides. Only sulfides of alkali metals, barium and ammonium are highly soluble in water. Sulfides of other metals are practically insoluble in water; they precipitate when a solution of ammonium sulfide (NH4)2S is introduced into solutions of metal salts. Many sulfides are brightly colored. Hydrosulfides M+HS and M2+(HS)? are also known for alkali and alkaline earth metals. Ca?+ and Sr2+ hydrosulfides are very unstable. Being salts of a weak acid, soluble sulfides undergo hydrolysis. The hydrolysis of sulfides containing metals in high oxidation states (Al?S3, Cr2S3, etc.) is often irreversible. Many natural sulfides in the form of minerals are valuable ores (pyrite, chalcopyrite, cinnabar). Polysulfides - polysulfur compounds of the general formula Me2Sn, for example, ammonium polysulfide (NH4)2Sn. In the structure of these compounds there are chains of atoms -S-S(n)-S. Numerous hydrogen polysulfides are known, with the general formula H2Sn, where n varies from 2 to 23. These are yellow oily liquids, as the sulfur content increases, the color changes from yellow to red. Alkali metal polysulfides are formed by the interaction of elemental sulfur with the corresponding sulfide (when fused or in a concentrated solution):
Na2S + 2 S(diamond) > Na2S3
Na2S + 4S > Na2S5
Na2S + 5S > Na2S6
Na2S + 6S > Na2S7
Na2S + 7S > Na2S8
Usually, the number of sulfur atoms in polysulfide molecules varies from 2 to 8; only one compound with n = 9 is known, this is (NH4)2S9. The most common are polysulfides with two sulfur atoms. These polysulfides can be considered as analogues of the corresponding peroxides. Polysulfides are characterized by oxidizing and reducing properties:
(NH4)2S2 + Sn+2S > (NH4)2Sn+4S3
4FeS2 +11O2 > 2Fe2O3 + 8SO2
When interacting with acids, they decompose with the release of sulfur and H2S. Polysulfides are used in analytical chemistry to separate elements, in the production of some rubbers, etc. A mixture of sodium polysulfides (in the old days it was called "sulphurous liver") has long been used in the leather industry to remove hair.
Introduction
The textbook on the chemistry of chalcogens is the second in a series devoted to the chemistry of the elements of the main subgroups of the periodic system of D.I. Mendeleev. It was written on the basis of a course of lectures on inorganic chemistry delivered at Moscow State University over the past 10 years by Academician Yu.D. Tretyakov and Professor V.P. Zlomanov.
In contrast to previously published methodological developments, the manual presents new factual material (catenation, a variety of chalcogen oxoacids (VI), etc.), a modern explanation of the patterns of changes in the structure and properties of chalcogen compounds using the concepts of quantum chemistry, including the molecular orbital method, relativistic effect, etc. The material of the manual was selected for the purpose of illustrative illustration of the relationship between the theoretical course and practical training in inorganic chemistry.
[previous section] [table of contents]§ one. General characteristics of chalcogens (E).
The elements of the VI main subgroup (or the 16th group according to the new IUPAC nomenclature) of the periodic system of elements of D.I. Mendeleev include oxygen (O), sulfur (S), selenium (Se), tellurium (Te) and polonium (Po). The group name of these elements is chalcogens(term "chalcogen" comes from the Greek words "chalkos" - copper and "genos" - born), that is, "giving birth to copper ores", due to the fact that in nature they occur most often in the form of copper compounds (sulfides, oxides, selenides, etc. ).
In the ground state, chalcogen atoms have the electronic configuration ns 2 np 4 with two unpaired p-electrons. They belong to even elements. Some properties of chalcogen atoms are presented in Table 1.
When moving from oxygen to polonium, the size of atoms and their possible coordination numbers increase, while the ionization energy (E ion) and electronegativity (EO) decrease. By electronegativity (EO), oxygen is second only to the fluorine atom, and the sulfur and selenium atoms are also inferior to nitrogen, chlorine, bromine; oxygen, sulfur and selenium are typical non-metals.
In compounds of sulfur, selenium, tellurium with oxygen and halogens, oxidation states +6, +4 and +2 are realized. With most other elements, they form chalcogenides, where they are in the -2 oxidation state.
Table 1. Properties of atoms of elements of group VI.
|
Properties |
|||||
| atomic number | |||||
| Number of stable isotopes | |||||
| Electronic configuration |
3d 10 4s 2 4p 4 |
4d 10 5s 2 5p 4 |
4f 14 5d 10 6s 2 6p 4 |
||
| Covalent radius, E | |||||
| First ionization energy, E ion, kJ/mol | |||||
| Electronegativity (Pauling) | |||||
| Affinity of an atom to an electron, kJ/mol |
The stability of compounds with the highest oxidation state decreases from tellurium to polonium, for which compounds with oxidation states of 4+ and 2+ are known (for example, PoCl 4 , PoCl 2 , PoO 2). This may be due to an increase in the bond strength of 6s 2 electrons with the nucleus due to relativistic effect. Its essence is to increase the speed of movement and, accordingly, the mass of electrons in elements with a large nuclear charge (Z> 60). The "weighting" of electrons leads to a decrease in the radius and an increase in the binding energy of 6s electrons with the nucleus. This effect is more clearly manifested in compounds of bismuth, an element of group V, and is discussed in more detail in the corresponding manual.
The properties of oxygen, as well as other elements of the 2nd period, differ from the properties of their heavier counterparts. Owing to the high electron density and strong interelectron repulsion, the electron affinity and E-E bond strength of oxygen is less than that of sulfur. Metal-oxygen (M-O) bonds are more ionic than M-S, M-Se, etc. bonds. Due to the smaller radius, the oxygen atom, unlike sulfur, is able to form strong -bonds (p - p) with other atoms - for example, oxygen in the ozone molecule, carbon, nitrogen, phosphorus. When moving from oxygen to sulfur, the strength of a single bond increases due to a decrease in interelectronic repulsion, and the strength of a bond decreases, which is associated with an increase in radius and a decrease in the interaction (overlap) of p-atomic orbitals. Thus, if oxygen is characterized by the formation of multiple (+) bonds, then sulfur and its analogues are characterized by the formation of single chain bonds - E-E-E (see § 2.1).
There are more analogies in the properties of sulfur, selenium and tellurium than with oxygen and polonium. So, in compounds with negative oxidation states, reducing properties increase from sulfur to tellurium, and in compounds with positive oxidation states, oxidizing properties.
Polonium is a radioactive element. The most stable isotope is obtained by bombarding nuclei with neutrons and subsequent -decay:
( 1/2 = 138.4 days).
The decay of polonium is accompanied by the release of a large amount of energy. Therefore, polonium and its compounds decompose solvents and vessels in which they are stored, and the study of Po compounds presents considerable difficulties.
[previous section] [table of contents]§ 2. Physical properties of simple substances.
Table 2. Physical properties of simple substances.
Density |
Temperatures, o C |
Heat of atomization, kJ/mol |
Electrical Resistance (25 ° C), Ohm. cm |
|||
melting |
||||||
| S | ||||||
| Se | hex. | |||||
1.3. 10 5 (liquid, 400 o C) |
||||||
| Those hex. | hex. | |||||
| Ro | ||||||
With an increase in the covalent radius in the O-S-Se-Te-Po series, the interatomic interaction and the corresponding temperatures of phase transitions, as well as atomization energy, that is, the energy of the transition of solid simple substances into the state of a monatomic gas, increases. The change in the properties of chalcogens from typical non-metals to metals is associated with a decrease in the ionization energy (Table 1) and structural features. Oxygen and sulfur are typical dielectrics, that is, substances that do not conduct electricity. Selenium and tellurium - semiconductors[substances whose electrophysical properties are intermediate between the properties of metals and non-metals (dielectrics). The electrical conductivity of metals decreases, and that of semiconductors increases with increasing temperature, which is due to the peculiarities of their electronic structure)], and polonium is a metal.
[previous section] [table of contents] [next section]§ 2.1. Chalcogen catenation. Allotropy and polymorphism.
One of the characteristic properties of chalcogen atoms is their ability to bind to each other in rings or chains. This phenomenon is called catenation. The reason for this is related to the different strengths of single and double bonds. Consider this phenomenon on the example of sulfur (Table 3).
Table 3. Energies of single and double bonds (kJ/mol).
It follows from the given values that the formation of two single -bonds for sulfur instead of one double (+) is associated with a gain in energy (530 - 421 = 109 J / mol). For oxygen, on the contrary, one double bond is energetically preferable (494-292=202 kJ/mol) than two single bonds. The decrease in the strength of the double bond upon the transition from O to S is associated with an increase in the size of the p-orbitals and a decrease in their overlap. Thus, for oxygen, catenation is limited to a small number of unstable compounds: O 3 ozone, O 4 F 2 .
Allotropy and polymorphism of simple substances are associated with catenation. Allotropy is the ability of the same element to exist in different molecular forms. The phenomenon of allotropy is attributed to molecules containing a different number of atoms of the same element, for example, O 2 and O 3, S 2 and S 8, P 2 and P 4, etc. The concept of polymorphism applies only to solids. Polymorphism- the ability of a solid substance with the same composition to have a different spatial structure. Examples of polymorphic modifications are monoclinic sulfur and rhombic sulfur, consisting of the same S 8 cycles, but placed differently in space (see § 2.3). Let us first consider the properties of oxygen and its allotropic form - ozone, and then the polymorphism of sulfur, selenium and tellurium.
