Electronic configuration of the atom. Arsenic element
Some who died of cholera in the Middle Ages did not die of it. The symptoms of the disease are similar to the manifestations arsenic poisoning.
Realizing this, medieval businessmen began to offer the trioxide of the element as a poison. Substance. A lethal dose is only 60 grams.
They were divided into portions, given over several weeks. As a result, no one suspected that the man did not die of cholera.
Arsenic taste is not felt in small doses, for example, in food or drinks. In modern realities, of course, there is no cholera.
People don't have to fear arsenic. Rather, mice need to be afraid. A toxic substance is a type of poison for rodents.
In their honor, by the way, the element is named. The word "arsenic" is used only in Russian-speaking countries. The official name of the substance is arsenicum.
The designation in - As. The ordinal number is 33. Based on it, we can assume a complete list of the properties of arsenic. But let's not assume. Let's study the question for sure.
Arsenic properties
The Latin name of the element is translated as "strong". Apparently, this refers to the effect of the substance on the body.
With intoxication, vomiting begins, digestion is upset, the stomach twists and work is partially blocked nervous system... not weak.
Poisoning occurs from any of the allotropic forms of the substance. Alltropy is the existence of manifestations of the same element. Arsenic most stable in metal form.
Steel-gray rhombohedral brittle. The units have a characteristic metallic, but from contact with moist air dim.
Arsenic - metal, whose density is almost 6 grams per cubic centimeter. The rest of the element forms have a lower indicator.
In second place is amorphous arsenic. Element characteristic: - almost black.
The density of this shape is 4.7 grams per cubic centimeter. Outwardly, the material resembles.
The usual state of arsenic for ordinary people is yellow. Cubic crystallization is unstable, transforms into amorphous when heated to 280 degrees Celsius, or under the influence of simple light.
Therefore, yellows are soft, like in the dark. Despite the color, the aggregates are transparent.
From a number of modifications of the element, it can be seen that it is only half metal. The obvious answer to the question: - " Arsenic metal, or non-metal", No.
Chemical reactions are evidence. The 33rd element is acid-forming. However, being in the acid itself does not give.
Metals act differently. In the case of arsenic, they do not work even when in contact with one of the strongest.
Salt-like compounds are "born" in the course of reactions of arsenic with active metals.
This refers to oxidants. The 33rd substance interacts only with them. If the partner does not have pronounced oxidative properties, the interaction will not take place.
This even applies to alkalis. That is, arsenic - a chemical element pretty inert. How, then, can you get it if the list of reactions is very limited?
Arsenic mining
Arsenic is mined along the way to other metals. Separate them, the 33rd substance remains.
There are in nature arsenic compounds with other elements... It is from them that the 33rd metal is extracted.
The process is beneficial, since, together with arsenic, they often go,, and.
It is found in granular masses or cubic pewter crystals. Sometimes, a yellow tint is present.
Arsenic compound and metal ferrum has a "brother" in which instead of the 33rd substance is. It is a common gold-colored pyrite.
The aggregates are similar to arseno version, but they cannot serve as arsenic ore, although they also contain impurities.
Arsenic in the usual, by the way, also happens, but, again, as an impurity.
The amount of the element per ton is so small, but even the side extraction does not make sense.
If we evenly distribute the world's arsenic reserves in the earth's crust, we get only 5 grams per ton.
So, the element is not common, in terms of quantity it is comparable to,,.
If you look at the metals with which arsenic forms minerals, then this is not only, but also with cobalt and nickel.
The total number of minerals of the 33rd element reaches 200. There is also a native form of the substance.
Its presence is explained by the chemical inertness of arsenic. Forming next to elements with which reactions are not provided, the hero remains in splendid isolation.
In this case, needle-like or cubic aggregates are often obtained. Usually, they grow together.
Arsenic use
The element arsenic belongs to dual, not only showing the properties of both metal and non-metal.
The perception of the element by humanity is also dual. In Europe, the 33rd substance has always been considered a poison.
In 1733, a decree was even issued prohibiting the sale and purchase of arsenic.
In Asia, the "poison" has been used by physicians for 2000 years in the treatment of psoriasis and syphilis.
Doctors of the modern have proved that the 33rd element attacks proteins that provoke oncology.
In the 20th century, some European doctors also took the side of the Asians. In 1906, for example, Western pharmacists invented the drug salvarsan.
He became the first in official medicine, used against a number of infectious diseases.
True, the drug, like any constant intake of arsenic in small doses, develops immunity.
1-2 courses of the drug are effective. If immunity has developed, people can take a lethal dose of the element and remain alive.
In addition to doctors, metallurgists became interested in the 33rd element, becoming added to the production of shot.
It is done on the basis that is included in heavy metals. Arsenic increases the lead and allows its spherical shape during casting. It is correct, which improves the quality of the shot.
Arsenic can also be found in thermometers, or rather them. It is called Viennese, mixed with the oxide of the 33rd substance.
The compound serves as a clarifier. Arsenic was also used by ancient glassblowers, but as a matting additive.
Glass becomes opaque with an impressive admixture of a toxic element.
Observing proportions, many glassblowers fell ill and died prematurely.
And tanners use sulfides arsenic.
Element the main subgroups The 5th group of the periodic table is part of some paints. In the leather industry, arsenicum helps to remove hair from.
Arsenic price
Pure arsenic is most often offered in metallic form. Prices are set per kilogram or ton.
1000 grams costs about 70 rubles. For metallurgists, they offer ready-made, for example, arsenic with copper.
In this case, they take 1500-1900 rubles per kilo. Arsenous anhydrite is also sold in kilograms.
It is used as a skin medicine. The agent is necrotic, that is, it deadens the affected area, killing not only the causative agent of the disease, but also the cells themselves. The method is radical, but effective.
Test
Write down the electronic formulas of arsenic and vanadium atoms. Indicate on which sublevels the valence electrons are located in the atoms of these elements.
Electronic formulas reflect the distribution of electrons in an atom by energy levels, sublevels (atomic orbitals). Electronic configuration is indicated by groups of symbols nl x, where n Is the principal quantum number, l- orbital quantum number (instead of it indicate the corresponding letter designation - s, p, d, f), x- the number of electrons in a given sublevel (orbital). It should be borne in mind that the electron occupies the energy sublevel at which it has the lowest energy - the smaller amount n+1 (Klechkovsky rule). The sequence of filling energy levels and sublevels is as follows:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → (5d 1) → 4f → 5d → 6p → 7s → (6d 1-2) → 5f → 6d → 7p
Since the number of electrons in an atom of one or another element is equal to its ordinal number in the table of D.I. Mendeleev, then for the elements arsenic (As ordinal number 33) and vanadium (V-ordinal number 23), the electronic formulas are:
V 23 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3
Аs 33 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3
The valence electrons of vanadium - 4s 2 3d 3 - are located at the 4s and 3d sublevels;
The valence electrons of arsenic 4s 2 4p 3 are located at the 4s and 4p sublevels. Thus, these elements are not electronic counterparts and should not be placed in the same subgroup. But the valence orbitals of the atoms of these elements have the same number of electrons - 5. Therefore, both elements are placed in the same group of the Mendeleev's periodic system.
Which element - phosphorus or antimony - has more pronounced oxidizing properties? Give an answer based on a comparison of the electronic structures of the atoms of these elements.
Phosphorus is the 15th element in the Periodic Table of D.I. Mendeleev. Its electronic formula is 1s 2 2s 2 2p 6 3s 2 3p 3
Antimony is the 51st element in the Periodic Table of D.I. Mendeleev. Its electronic formula is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 3
On the outer electronic sublevels of these elements there are 5 electrons each, therefore they belong to the 5th group of the periodic system.
Oxidizing properties associated with the position of the elements in the Periodic Table of D.I. Mendeleev. In each group of the Periodic Table, an element with a higher ordinal number has more pronounced reducing properties in its group, and an element with a lower ordinal number has stronger oxidizing properties.
Phosphorus has more oxidizing properties than antimony. since the radius of the atom is smaller and the valence electrons are more strongly attracted to the nucleus.
Why does nitrogen, oxygen, fluorine, iron, cobalt and nickel have a maximum valence lower than the number of the group in which the indicated elements are located, while their electronic counterparts have the maximum valence corresponding to the group number?
The properties of elements, the shape and properties of compounds of elements are periodically dependent on the magnitude of the charge of the nuclei of their atoms.
The highest oxidation state of an element is determined by the group number of the periodic system of D.I. Mendeleev, in which he is. The lowest oxidation state is determined by the conditional charge that an atom acquires when adding the number of electrons that is necessary to form a stable eight-electron shell (ns 2 nр 6).
Since the elements of the second period lack the d-sublevel, nitrogen, oxygen and fluorine cannot reach a valence equal to the group number. They do not have the ability to steam electrons. Fluorine has a maximum valence of one, oxygen has two, and nitrogen has three. Excitation of a 2s electron can only occur to a level with n = 3, which is energetically extremely disadvantageous. s-electron by 3 d- too big. The interaction of atoms with the formation of a bond between them occurs only in the presence of orbitals with close energies, i.e. orbitals with the same principal quantum number In contrast to nitrogen, oxygen, fluorine, phosphorus, sulfur, chlorine atoms can form, respectively, five, six, seven covalent bonds .. In this case, the participation of 3s electrons in the formation of bonds is possible, since d-AO (3d) have the same principal quantum number.
For most d-elements, the highest valence may differ from the group number. The valence possibilities of the d-element in a particular case are determined by the structure of the electron shell of the atom. d-elements can have a minimum valence above the group number (copper, silver) and below the group number (iron, cobalt, nickel).
Thermochemical equation of the reaction:
CO (g) +2H 2 (d) =CH 3 OH(g) +128 kJ
Calculate at what temperature equilibrium occurs in this system?
In exothermic reactions, the enthalpy of the system decreases and ΔH< 0 (Н 2 < H 1). Тепловые эффекты выражаются через ΔH.
The thermochemical calculations are based on Hess's law (1840): the thermal effect of a reaction depends only on the nature and physical state of the initial substances and final products, but does not depend on the transition path.
In thermochemical calculations, a consequence of Hess's law is used more often: the thermal effect of the reaction (ΔHx.р) is equal to the sum of the enthalpies of formation of the reaction products minus the sum of the enthalpies of formation of the initial substances, taking into account the stoichiometric coefficients.
Entropy S, also enthalpy H is a property of a substance, proportional to its amount. Entropy is a function of state, i.e. its change (ΔS) depends only on the initial (S 1) and final (S 2) states and does not depend on the path of the process:
ΔSх.р = ΣS 0 prod - ΣS 0 ref.
Since the entropy increases with increasing temperature, it can be assumed that
that the measure of disorder is ≈ ТΔS. At Р = const and Т = const, the total driving force of the process, which is denoted by ΔG, can be found from the relation:
ΔG = (H 2 - H 1) - (TS 2 - TS 1); ΔG = ΔH - TΔS.
Chemical equilibrium is a state of the system in which the rate of the forward reaction (V 1) is equal to the rate of the reverse reaction (V 2). At chemical equilibrium the concentration of substances remains unchanged. Chemical equilibrium is dynamic: direct and reverse reactions do not stop at equilibrium
Into a state of equilibrium
ΔG = 0 and ΔH = TΔS.
Find ΔS. for a given system:
S 0 (CO) = 197.55 ∙ 10 -3 kJ / mol K;
S 0 (H 2) = 130.52 · 10 -3 kJ / mol · K;
S 0 (CH 3 OH) = 126.78 · 10 -3 kJ / mol · K;
ΔSх.р = 126.78 · 10 -3 - (197.55 ∙ 10 -3 + 2 · 130.52 · 10 -3) = - 331.81 · 10 -3
From the equilibrium condition
ΔH = TΔS we find T = ΔH / ΔS
Calculate the temperature coefficient of the reaction (γ) if the rate constant of this reaction at 120 degrees C is 5.88 ∙ 10 -4 , and at 170 degrees C 6.7 ∙ 10 -2
The dependence of the rate of a chemical reaction on temperature is determined by the Van't Hoff rule of thumb according to the formula:
,
where v t 1, v t 2 are the reaction rates at the initial (t 1) and final (t 2) temperatures, respectively, and γ is the temperature coefficient of the reaction rate, which shows how many times the reaction rate increases with an increase in the temperature of the reacting substances by 10º.
Hence it follows that
,
Based on the condition of the problem, it follows that:
, whence γ 5 = 113.94;
In which direction will the equilibrium shift in systems with increasing pressure:
2NO + O 2 - 2NO 2
4HCI (G) + O 2 - 2H 2 O (G) + 2CI 2
H 2 + S(To) -H 2 S
Le Chatelier's principle (the principle of displacement of equilibrium), establishes that an external influence, which brings the system out of the state of thermodynamic equilibrium, causes processes in the system that tend to weaken the effect of the influence.
With an increase in pressure, a shift in equilibrium is associated with a decrease in the total volume of the system, and a decrease in pressure is accompanied by physical. or chemical processes leading to an increase in volume.
2NO + O 2 → 2NO 2
2 mol + 1 mol → 2 mol
An increase in pressure leads to a shift in equilibrium towards a reaction leading to the formation of fewer molecules. Consequently, the equilibrium shifts towards the formation of NO 2 V pr> V sample.
4HCI (g) + O 2 → 2H 2 O (g) + 2CI 2
4 moles + 1 mole → 4 moles
An increase in pressure leads to a shift in equilibrium towards a reaction leading to the formation of fewer molecules. Therefore V pr> V arr
H 2 + S (k) → H 2 S
no volume change occurs during the reaction. Consequently, a change in pressure does not in any way affect the displacement of the equilibrium of the reaction.
6.6. Features of the electronic structure of atoms of chromium, copper and some other elements
If you carefully looked at Appendix 4, you probably noticed that the sequence of filling the orbitals with electrons is violated in the atoms of some elements. Sometimes these violations are called "exceptions", but they are not - there are no exceptions to the laws of Nature!
The first element with this violation is chrome. Let's take a closer look at its electronic structure (Fig. 6.16 a). The chromium atom has 4 s-sub-level not two, as one would expect, but only one electron. But at 3 d-sublevel five electrons, but this sublevel is filled after 4 s-sublevel (see Fig. 6.4). To understand why this is happening, let's see what electron clouds are 3 d is the sublevel of this atom.
Each of five 3 d-clouds in this case is formed by one electron. As you already know from § 4 of this chapter, the total electron cloud of these five electrons has a spherical shape, or, as they say, spherically symmetric. By the nature of the electron density distribution in different directions, it is similar to 1 s-EO. The energy of the sublevel, the electrons of which form such a cloud, turns out to be less than in the case of a less symmetric cloud. In this case, the energy of the orbitals is 3 d-sublevel is equal to energy 4 s-orbital. When symmetry is broken, for example, when the sixth electron appears, the energy of the orbitals is 3 d-sublevel again becomes larger than energy 4 s-orbital. Therefore, the manganese atom again has a second electron by 4 s-AO.
The general cloud of any sublevel, filled with electrons, both half and completely, has spherical symmetry. The decrease in energy in these cases is of a general nature and does not depend on whether any sublevel is half or completely filled with electrons. And if so, then we should look for the next violation in the atom, into the electronic shell of which the ninth "comes" last d-electron. Indeed, the copper atom has 3 d-sublayer 10 electrons, and on 4 s- there is only one sublevel (Fig. 6.16 b).
The decrease in the energy of the orbitals of a fully or half-filled sublevel is the cause of a number of important chemical phenomena, some of which you will become familiar with.
6.7. Outer and valence electrons, orbitals and sublevels
In chemistry, the properties of isolated atoms, as a rule, are not studied, since almost all atoms, being part of various substances, form chemical bonds... Chemical bonds are formed when the electronic shells of atoms interact. In all atoms (except hydrogen), not all electrons take part in the formation of chemical bonds: boron has three electrons out of five, carbon has four out of six, and, for example, barium has two out of fifty-six. These "active" electrons are called valence electrons.
Sometimes valence electrons are confused with external electrons, and they are not the same thing.
The electron clouds of the outer electrons have a maximum radius (and maximum principal quantum number).
It is the outer electrons that take part in the formation of bonds in the first place, if only because when the atoms approach each other, the electron clouds formed by these electrons come into contact first of all. But together with them, part of the electrons can take part in the formation of a bond. pre-external(penultimate) layer, but only if they have an energy that does not differ much from the energy of the outer electrons. Both those and other electrons of an atom are valence. (In lanthanides and actinides, even some "pre-external" electrons are valence)
The energy of valence electrons is much higher than the energy of other electrons of the atom, and valence electrons differ significantly less in energy from each other.
External electrons are always valence only if the atom can form chemical bonds at all. So, both electrons of the helium atom are external, but they cannot be called valence, since the helium atom does not form any chemical bonds at all.
Valence electrons occupy valence orbitals, which in turn form valence sublevels.
As an example, consider an iron atom, the electronic configuration of which is shown in Fig. 6.17. Of the electrons of the iron atom, the maximum principal quantum number ( n= 4) have only two 4 s-electron. Therefore, it is they who are the outer electrons of this atom. The outer orbitals of the iron atom are all orbitals with n= 4, and the outer sublevels are all sublevels formed by these orbitals, that is, 4 s-,
4p-, 4d- and 4 f-EPU.
The outer electrons are always valence, therefore, 4 s-electrons of the iron atom - valence electrons. And if so, then 3 d-electrons with slightly higher energy will also be valence. At the outer level of the iron atom, in addition to the filled 4 s-AO there are still free 4 p-, 4d- and 4 f-AO. They are all external, but there are only 4 valence among them. R-AO, since the energy of the remaining orbitals is much higher, and the appearance of electrons in these orbitals is not beneficial for the iron atom.
So, at the iron atom
external electronic level - fourth,
external sublevels - 4 s-, 4p-, 4d- and 4 f-EPU,
outer orbitals - 4 s-, 4p-, 4d- and 4 f-AO,
outer electrons - two 4 s-electron (4 s 2),
outer electron layer - the fourth,
external electronic cloud - 4 s-EO
valence sublevels - 4 s-, 4p-, and 3 d-EPU,
valence orbitals - 4 s-, 4p-, and 3 d-AO,
valence electrons - two 4 s-electron (4 s 2) and six 3 d-electrons (3 d 6).
The valence sublevels can be partially or completely filled with electrons, or they can generally remain free. With an increase in the nuclear charge, the values of the energy of all sublevels decrease, but due to the interaction of electrons with each other, the energy of different sublevels decreases with different "rates". The energy is completely filled d- and f-sublevels decreases so much that they cease to be valence.
As an example, consider the atoms of titanium and arsenic (Fig. 6.18).
In the case of the titanium atom 3 d-The EPU is only partially filled with electrons, and its energy is greater than the energy 4 s-EPU, and 3 d-electrons are valence. Arsenic atom has 3 d-The EPU is completely filled with electrons, and its energy is significantly less than the energy 4 s-EPU, and therefore 3 d-electrons are not valence.
In the examples given, we have analyzed valence electronic configuration
atoms of titanium and arsenic.
The valence electronic configuration of an atom is depicted as valence electronic formula, or in the form energy diagram of valence sublevels.
VALENT ELECTRONS, EXTERNAL ELECTRONS, VALENT EPU, VALENT AO, VALENT ELECTRONIC CONFIGURATION OF THE ATOM, VALENT ELECTRONIC FORMULA, VALENCE SUB-LEVEL DIAGRAM.
1. On the energy diagrams you have drawn up and in the complete electronic formulas of the atoms Na, Mg, Al, Si, P, S, Cl, Ar, indicate the outer and valence electrons. Make up the valence electronic formulas of these atoms. On the energy diagrams, highlight the parts corresponding to the energy diagrams of the valence sublevels.
2. What is common between the electronic configurations of atoms a) Li and Na, B and Al, O and S, Ne and Ar; b) Zn and Mg, Sc and Al, Cr and S, Ti and Si; c) H and He, Li and O, K and Kr, Sc and Ga. What are their differences
3.How many valence sublevels in the electron shell of an atom of each of the elements: a) hydrogen, helium and lithium, b) nitrogen, sodium and sulfur, c) potassium, cobalt and germanium
4. How many valence orbitals are completely filled in the atom of a) boron, b) fluorine, c) sodium?
5.How many orbitals with an unpaired electron in an atom of a) boron, b) fluorine, c) iron
6. How many free outer orbitals does a manganese atom have? And how many free valences?
7. For the next lesson, prepare a strip of paper 20 mm wide, divide it into cells (20 × 20 mm), and apply a natural range of elements (from hydrogen to meitnerium) to this strip.
8.In each cell, place the symbol of the element, its ordinal number and the valence electronic formula, as shown in Fig. 6.19 (use Appendix 4).
6.8. Systematization of atoms according to the structure of their electronic shells
The systematization of chemical elements is based on the natural series of elements
and principle of similarity of electronic shells their atoms.
You are already familiar with the natural range of chemical elements. Now let's get acquainted with the principle of similarity of electronic shells.
Considering the valence electronic formulas of atoms in the NRE, it is easy to find that for some atoms they differ only in the values of the principal quantum number. For example 1 s 1 for hydrogen, 2 s 1 for lithium, 3 s 1 for sodium, etc., or 2 s 2 2p 5 for fluorine, 3 s 2 3p 5 for chlorine, 4 s 2 4p 5 for bromine, etc. This means that the outer regions of the clouds of valence electrons of such atoms are very similar in shape and differ only in size (and, of course, in electron density). And if so, then the electron clouds of such atoms and the corresponding valence configurations can be called like... For atoms of different elements with similar electronic configurations, we can write general valence electronic formulas: ns 1 in the first case and ns 2 np 5 in the second. Moving along the natural series of elements, one can find other groups of atoms with similar valence configurations.
In this way, atoms with similar valence electronic configurations are regularly found in the natural series of elements.
This is the principle of the similarity of electronic shells.
Let's try to identify the kind of this regularity. To do this, we will use the natural row of elements you have made.
ERE begins with hydrogen, the valence electronic formula of which is 1 s one . In search of similar valence configurations, we cut the natural series of elements before elements with a common valence electronic formula ns 1 (i.e., before lithium, before sodium, etc.). We got the so-called "periods" of the elements. Let's add the resulting "periods" so that they become the rows of the table (see Fig. 6.20). As a result, only the atoms of the first two columns of the table will have similar electronic configurations.
Let's try to achieve similarity of valence electronic configurations in other columns of the table. To do this, we cut out from the 6th and 7th periods elements with numbers 58 - 71 and 90 - 103 (they are filled with 4 f- and 5 f-sub-levels) and place them below the table. Let's move the symbols of the remaining elements horizontally as shown in the figure. After that, atoms of elements standing in one column of the table will have similar valence configurations, which can be expressed by general valence electronic formulas: ns 1 , ns 2 , ns 2 (n–1)d 1 , ns 2 (n–1)d 2 and so on until ns 2 np 6. All deviations from the general valence formulas are due to the same reasons as in the case of chromium and copper (see paragraph 6.6).
As you can see, using the ERE and applying the principle of similarity of electronic shells, we were able to systematize the chemical elements. Such a system of chemical elements is called natural, since it is based solely on the laws of Nature. The table we received (Fig. 6.21) is one of the ways to graphically depict the natural system of elements and is called long-period table of chemical elements.
THE PRINCIPLE OF SIMILARITY OF ELECTRONIC SHELLS, NATURAL SYSTEM OF CHEMICAL ELEMENTS ("PERIODIC" SYSTEM), TABLE OF CHEMICAL ELEMENTS.
6.9. Long-period table of chemical elements
Let's take a closer look at the structure of the long-period table of chemical elements.
The rows in this table, as you already know, are called "periods" of the elements. The periods are numbered with Arabic numerals from 1 to 7. There are only two elements in the first period. The second and third periods, each containing eight elements, are called short periods. The fourth and fifth periods, each containing 18 elements, are called long periods. The sixth and seventh periods, each containing 32 elements, are called extra-long periods.
The columns in this table are named in groups elements. Group numbers are designated by Roman numerals with Latin letters A or B.
Elements of some groups have their own common (group) names: elements of the IA group (Li, Na, K, Rb, Cs, Fr) - alkaline elements(or alkali metal elements); Group IIA elements (Ca, Sr, Ba and Ra) - alkaline earth elements(or alkaline earth metal elements) (the name "alkali metals" and alkaline earth metals "refers to simple substances formed by the corresponding elements and should not be used as names for groups of elements); elements of group VIA (O, S, Se, Te, Po) - chalcogenes, elements of VIIA group (F, Cl, Br, I, At) - halogens, elements of group VIIIA (He, Ne, Ar, Kr, Xe, Rn) - noble gas elements. (The traditional name "noble gases" also refers to simple substances)
Usually carried out in lower part table elements with serial numbers 58 - 71 (Ce - Lu) are called lanthanides("following lanthanum"), and elements with serial numbers 90 - 103 (Th - Lr) - actinides("following anemones"). There is a variant of the long-period table, in which lanthanides and actinides are not cut out from the NRE, but remain in place in super-long periods. This table is sometimes called superlong-period.
Long period table is divisible by four block(or section).
s-Block includes elements of IA and IIA-groups with common valence electronic formulas ns 1 and ns 2
(s-elements).
r-Block includes elements from IIIA to VIIIA group with general valence electronic formulas from ns 2 np 1 to ns 2 np 6 (p-elements).
d-block includes elements from IIIB to IIB group with general valence electronic formulas from ns 2 (n–1)d 1 to ns 2 (n–1)d 10 (d-elements).
f-Block includes lanthanides and actinides ( f-elements).
The elements s- and p-blocks form A-groups, and the elements d-block - B-group of the system of chemical elements. Everything f-elements are formally included in the IIIB group.
The elements of the first period - hydrogen and helium - are s-elements and can be placed in groups IA and IIA. But helium is more often placed in the VIIIA group as an element that ends the period, which fully corresponds to its properties (helium, like all other simple substances formed by the elements of this group, is a noble gas). Hydrogen, on the other hand, is often placed in the VIIA group, since in its properties it is much closer to halogens than to alkaline elements.
Each of the periods of the system begins with an element with a valence configuration of atoms ns 1, since it is from these atoms that the formation of the next electron layer begins, and ends with an element with a valence configuration of atoms ns 2 np 6 (except for the first period). This makes it easy to distinguish groups of sublevels on the energy diagram, which are filled with electrons from atoms of each of the periods (Fig. 6.22). Do this work with all the sublevels shown in your copy of Figure 6.4. The sublevels highlighted in Figure 6.22 (except for completely filled d- and f-sublevels) are valence for atoms of all elements of a given period.
Appearance in periods s-, p-, d- or f-elements fully correspond to the filling sequence s-, p-, d- or f-sub-level electrons. This feature of the system of elements allows, knowing the period and the group that this element belongs to, immediately write down its valence electronic formula.
LONG-PERIOD TABLE OF CHEMICAL ELEMENTS, BLOCKS, PERIODS, GROUPS, ALKALINE ELEMENTS, ALKALINE EARTH ELEMENTS, CHALCOGENS, HALOGENS, ELEMENTS OF NOBLE GASES, LANTANOIDS.
Write down the general valence electronic formulas of atoms of elements a) IVA and IVB groups, b) IIIA and VIIB groups?
2. What is common between the electronic configurations of atoms of elements A and B groups? How do they differ?
3. How many groups of elements are included in a) s-block, b) R-block, c) d-block?
4. Continue Figure 30 towards increasing the energy of the sublevels and select the groups of sublevels that are filled with electrons in the 4th, 5th and 6th periods.
5. List the valence sublevels of a) calcium, b) phosphorus, c) titanium, d) chlorine, e) sodium. 6.Formulate how s-, p- and d-elements differ from each other.
7. Explain why the belonging of an atom to any element is determined by the number of protons in the nucleus, and not by the mass of this atom.
8. For atoms of lithium, aluminum, strontium, selenium, iron and lead, draw up valence, complete and abbreviated electronic formulas and draw energy diagrams of valence sublevels. 9. Atoms of which elements correspond to the following valence electronic formulas: 3 s 1 , 4s 1 3d 1, 2s 2 2 p 6 , 5s 2 5p 2 , 5s 2 4d 2 ?
6.10. Types of electronic formulas of the atom. Algorithm for their compilation
For different purposes, we need to know either the full or the valence configuration of the atom. Each of these electronic configurations can be represented by both a formula and an energy diagram. That is, total electronic configuration of an atom expressed the complete electronic formula of the atom, or complete energy diagram of the atom... In turn, valence electron configuration of an atom expressed valence(or, as it is often called, " short ") electronic formula of the atom, or diagram of the valence sublevels of the atom(fig. 6.23).
Previously, we made up the electronic formulas of atoms using the ordinal numbers of the elements. In this case, we determined the sequence of filling the sublevels with electrons according to the energy diagram: 1 s, 2s,
2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p,
6s, 4f, 5d, 6p, 7s etc. And only by writing down the complete electronic formula, we could write down the valence formula.
The valence electronic formula of the atom, which is most often used, is more convenient to write based on the position of the element in the system of chemical elements, according to the coordinates period - group.
Let's consider in detail how this is done for elements. s-, p- and d-blocks.
For items s-block valence electronic formula of an atom consists of three symbols. In general, it can be written as follows:
In the first place (in place of the large cell), the period number is put (equal to the main quantum number of these s-electrons), and on the third (in the superscript) is the group number (equal to the number of valence electrons). Taking magnesium atom as an example (3rd period, IIA group), we get:
For items p-block valence electronic formula of an atom consists of six symbols:
Here, in place of large cells, the period number is also put (equal to the main quantum number of these s- and p-electrons), and the group number ( equal to the number valence electrons) turns out to be equal to the sum of the superscripts. For the oxygen atom (2nd period, VIA group) we get:
2s 2 2p 4 .
Valence electronic formula of most elements d-block can be written like this:
As in the previous cases, here, instead of the first cell, the period number is put (equal to the main quantum number of these s-electrons). The number in the second cell turns out to be one less, since the main quantum number of these d-electrons. The group number here is also equal to the sum of the indices. Example - the valence electronic formula of titanium (4th period, IVB group): 4 s 2 3d 2 .
The group number is equal to the sum of the indices and for the elements of the VIB group, but they, as you remember, on the valence s- there is only one electron sublevel, and the general valence electronic formula ns 1 (n–1)d 5 . Therefore, the valence electronic formula, for example, molybdenum (5th period) - 5 s 1 4d 5 .
It is just as easy to compose the valence electronic formula of any element of the IB group, for example, gold (6th period)> -> 6 s 1 5d 10, but in this case it must be remembered that d- the electrons of the atoms of the elements of this group still remain valence, and some of them can participate in the formation of chemical bonds.
The general valence electronic formula of the atoms of group IIB elements is ns 2 (n – 1)d 10 . Therefore, the valence electronic formula, for example, of a zinc atom is 4 s 2 3d 10 .
The valence electronic formulas of the elements of the first triad (Fe, Co and Ni) also obey the general rules. Iron, an element of group VIIIB, has a valence electronic formula of 4 s 2 3d 6. The cobalt atom has one d-electron is larger (4 s 2 3d 7), and for the nickel atom - by two (4 s 2 3d 8).
Using only these rules for writing valence electronic formulas, it is impossible to compose the electronic formulas of atoms of some d-elements (Nb, Ru, Rh, Pd, Ir, Pt), since their filling of valence sublevels with electrons due to the tendency to highly symmetric electron shells has some additional features.
Knowing the valence electronic formula, it is possible to write down the complete electronic formula of the atom (see below).
Often, instead of cumbersome full electronic formulas, one writes abbreviated electronic formulas atoms. To compose them in the electronic formula, all the electrons of the atom except for the valence ones are selected, their symbols are placed in square brackets and the part of the electronic formula corresponding to the electronic formula of the atom of the last element of the previous period (the element that forms the noble gas) is replaced with the symbol of this atom.
Examples of electronic formulas of different types are shown in Table 14.
Table 14. Examples of electronic formulas of atoms
Electronic formulas |
|||
Abbreviated |
Valent |
||
1s 2 2s 2 2p 3 |
2s 2 2p 3 |
2s 2 2p 3 |
|
1s 2 2s 2 2p 6 3s 2 3p 5 |
3s 2 3p 5 |
3s 2 3p 5 |
|
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 |
4s 2 3d 5 |
4s 2 3d 5 |
|
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 3 |
4s 2 4p 3 |
4s 2 4p 3 |
|
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 |
4s 2 4p 6 |
4s 2 4p 6 |
|
Algorithm for drawing up the electronic formulas of atoms (for example, the iodine atom)
№ |
Operation |
Result |
|
Determine the coordinates of the atom in the table of elements. |
Period 5, group VIIA |
||
Make a valence electronic formula. |
5s 2 5p 5 |
||
Complete the symbols of the internal electrons in the sequence of filling the sublevels with them. |
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 5 |
||
Considering the decrease in energy of fully filled d- and f-Sub-levels, write down the complete electronic formula. |
|
||
Note the valence electrons. |
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 5 |
||
Highlight the electron configuration of the preceding noble gas atom. |
|||
Write down the abbreviated electronic formula, combining all in square brackets non-bonded electrons. |
5s 2 5p 5 |
Notes (edit)
1. For elements of the 2nd and 3rd periods, the third operation (without the fourth) immediately leads to a complete electronic formula.
2. (n – 1)d 10 -Electrons remain valence at the atoms of the elements of the IB group.
COMPLETE ELECTRONIC FORMULA, VALENCE ELECTRONIC FORMULA, REDUCED ELECTRONIC FORMULA, ALGORITHM FOR COMPOSING ELECTRONIC FORMULAS OF ATOMS.
1. Make the valence electronic formula of the atom of element a) the second period of the third A group, b) the third period of the second A group, c) the fourth period of the fourth A group.
2. Make the abbreviated electronic formulas of the atoms of magnesium, phosphorus, potassium, iron, bromine and argon.
6.11. Short-period table of chemical elements
For more than 100 years that have passed since the discovery of the natural system of elements, several hundred of the most diverse tables have been proposed, graphically reflecting this system. Of these, in addition to the long-period table, the most widespread is the so-called short-period table of elements by D.I.Mendeleev. A short-period table is obtained from a long-period table if the 4th, 5th, 6th and 7th periods are cut before the elements of the IB group, move apart and fold the resulting rows as we used to add the periods. The result is shown in Figure 6.24.
Lanthanides and actinides are also placed under the main table here.
V groups This table contains elements whose atoms have equal number of valence electrons no matter what orbitals these electrons are in. So, the elements chlorine (a typical element that forms a non-metal; 3 s 2 3p 5) and manganese (metal-forming element; 4 s 2 3d 5), not possessing the semblance of electronic shells, fall here in the same seventh group. The need to distinguish between such elements forces you to select in groups subgroups: the main- analogs of the A-groups of the long-period table and collateral- analogs of B-groups. In Figure 34, the symbols of the elements of the main subgroups are shifted to the left, and the elements of the secondary subgroups are shifted to the right.
True, such an arrangement of elements in the table also has its advantages, because it is the number of valence electrons that primarily determines the valence capabilities of an atom.
The long-period table reflects the regularities of the electronic structure of atoms, the similarity and regularities of changes in the properties of simple substances and compounds by groups of elements, the regular change in a number of physical quantities characterizing atoms, simple substances and compounds throughout the entire system of elements, and much more. The short-period table is less convenient in this respect.
SHORT PERIOD TABLE, MAIN SUBGROUPS, SIDE SUBGROUPS.
1.Convert the long-period table you built from a natural series of elements into a short-period one. Reverse the transformation.
2. Is it possible to draw up a general valence electronic formula of atoms of elements of one group of a short-period table? Why?
6.12. Sizes of atoms. Orbital radii
.The atom has no clear boundaries. What is considered the size of an isolated atom? The nucleus of an atom is surrounded by an electron shell, and the shell consists of electron clouds. The size of the EO is characterized by the radius r eo. All clouds in the outer layer have approximately the same radius. Therefore, the size of an atom can be characterized by this radius. It is called orbital radius of the atom(r 0).
The values of the orbital radii of the atoms are given in Appendix 5.
The radius of the EO depends on the charge of the nucleus and on which orbital the electron that forms this cloud is located in. Consequently, the orbital radius of an atom depends on the same characteristics.
Consider the electron shells of hydrogen and helium atoms. Both in the hydrogen atom and in the helium atom, the electrons are at 1 s-AO, and their clouds would have the same size if the charges of the nuclei of these atoms were the same. But the charge of the nucleus of a helium atom is twice as much as the charge of the nucleus of a hydrogen atom. According to Coulomb's law, the force of attraction acting on each of the electrons of a helium atom is twice the force of attraction of an electron to the nucleus of a hydrogen atom. Consequently, the radius of the helium atom must be much smaller than the radius of the hydrogen atom. And there is: r 0 (He) / r 0 (H) = 0.291 E / 0.529 E 0.55.
The lithium atom has an outer electron at 2 s-AO, that is, it forms a cloud of the second layer. Naturally, its radius should be larger. Really: r 0 (Li) = 1.586 E.
The atoms of the remaining elements of the second period have external electrons (and 2 s, and 2 p) are located in the same second electron layer, and the nuclear charge of these atoms increases with increasing serial number. Electrons are more strongly attracted to the nucleus, and, naturally, the radii of the atoms decrease. We could repeat this reasoning for atoms of elements of other periods, but with one clarification: the orbital radius decreases monotonically only when each of the sublevels is filled.
But if we ignore the particulars, then the general nature of the change in the size of atoms in the system of elements is as follows: with an increase in the ordinal number in the period, the orbital radii of the atoms decrease, and in the group, they increase. The largest atom is a cesium atom, and the smallest is a helium atom, but of the atoms of elements that form chemical compounds (helium and neon do not form them), the smallest is a fluorine atom.
Most of the atoms of the elements that are in the natural row after the lanthanides have orbital radii somewhat smaller than one would expect based on general laws. This is due to the fact that 14 lanthanides are located between lanthanum and hafnium in the system of elements, and, therefore, the charge of the nucleus of the hafnium atom is 14 e more than lanthanum. Therefore, the outer electrons of these atoms are attracted to the nucleus more strongly than they would in the absence of lanthanides (this effect is often called "lanthanide compression").
Please note that when going from the atoms of the elements of the VIIIA group to the atoms of the elements of the IA group, the orbital radius increases abruptly. Consequently, our choice of the first elements of each period (see § 7) turned out to be correct.
ORBITAL RADIUS OF THE ATOM, ITS CHANGE IN THE SYSTEM OF ELEMENTS.
1.According to the data given in Appendix 5, plot on graph paper a graph of the dependence of the orbital radius of an atom on the ordinal number of an element for elements with Z from 1 to 40. The length of the horizontal axis is 200 mm, the length of the vertical axis is 100 mm.
2. How can you characterize the appearance of the resulting broken line?
6.13. Ionization energy of an atom
If you give additional energy to an electron in an atom (how this can be done, you will learn from a physics course), then the electron can go to another AO, that is, the atom will be in excited state... This state is unstable, and the electron will almost immediately return to its original state, and the excess energy will be released. But if the energy imparted to the electron is large enough, the electron can completely detach from the atom, while the atom ionizes, that is, it turns into a positively charged ion ( cation). The energy required for this is called the ionization energy of the atom(E and).
It is rather difficult to tear an electron from a single atom and measure the energy required for this, therefore, it is practically determined and used molar ionization energy(E and m).
The molar ionization energy shows what is the smallest energy required to detach 1 mole of electrons from 1 mole of atoms (one electron from each atom). This value is usually measured in kilojoules per mole. The values of the molar ionization energy of the first electron for most elements are given in Appendix 6.
How does the ionization energy of an atom depend on the position of an element in a system of elements, that is, how does it change in a group and a period?
According to the physical meaning, the ionization energy is equal to the work that needs to be spent on overcoming the force of attraction of the electron to the atom when the electron moves from the atom to an infinite distance from it.
where q- electron charge, Q Is the charge of the cation remaining after the removal of the electron, and r o is the orbital radius of the atom.
AND q, and Q- the quantities are constant, and it can be concluded that, the work on the separation of an electron A, and with it the ionization energy E and, are inversely proportional to the orbital radius of the atom.
Having analyzed the values of the orbital radii of the atoms various elements and the corresponding values of the ionization energy given in Appendices 5 and 6, you can see that the relationship between these quantities is close to proportional, but somewhat different from it. The reason that our conclusion does not agree very well with the experimental data is that we used a very crude model that does not take into account many significant factors. But even this rough model allowed us to draw the correct conclusion that with an increase in the orbital radius, the ionization energy of an atom decreases and, conversely, with a decrease in the radius, it increases.
Since the orbital radius of the atoms decreases in the period with an increase in the ordinal number, the ionization energy increases. In a group, with an increase in the ordinal number, the orbital radius of the atoms, as a rule, increases, and the ionization energy decreases. The largest molar ionization energy is found in the smallest atoms, helium atoms (2372 kJ / mol), and among the atoms capable of forming chemical bonds, in fluorine atoms (1681 kJ / mol). The smallest is for the largest atoms, cesium atoms (376 kJ / mol). In a system of elements, the direction of increasing the ionization energy can be schematically shown as follows: 
In chemistry, it is important that the ionization energy characterizes the tendency of an atom to give up "its" electrons: the greater the ionization energy, the less the atom is inclined to give up electrons, and vice versa.
EXCITED STATE, IONIZATION, CATION, IONIZATION ENERGY, MOLAR IONIZATION ENERGY, CHANGE IN IONIZATION ENERGY IN THE ELEMENT SYSTEM.
1.Using the data given in Appendix 6, determine how much energy needs to be spent in order to take one electron away from all sodium atoms with a total mass of 1 g.
2. Using the data given in Appendix 6, determine how many times more energy needs to be spent to detach one electron from all sodium atoms with a mass of 3 g than from all potassium atoms of the same mass. Why is this ratio different from the ratio of the molar ionization energies of the same atoms?
3.According to the data given in Appendix 6, build a graph of the dependence of the molar ionization energy on the serial number for elements with Z from 1 to 40. The dimensions of the graph are the same as in the task to the previous paragraph. See if this schedule matches the selection of "periods" of the system of elements.
6.14. Electron affinity energy
.The second most important energy characteristic of an atom is electron affinity energy(E With).
In practice, as in the case of the ionization energy, the corresponding molar quantity is usually used - molar electron affinity energy().
The molar energy of affinity for an electron shows what is the energy released when one mole of electrons is attached to one mole of neutral atoms (one electron to each atom). Like the molar ionization energy, this value is also measured in kilojoules per mole.
At first glance, it may seem that energy should not be released in this case, because an atom is a neutral particle, and there are no electrostatic forces of attraction between a neutral atom and a negatively charged electron. On the contrary, approaching an atom, an electron, it would seem, should repel from the same negatively charged electrons that form an electron shell. In fact this is not true. Remember if you have ever had to deal with atomic chlorine. Of course not. After all, it only exists at very high temperatures. Even the more stable molecular chlorine is practically not found in nature - if necessary, it has to be obtained using chemical reactions. And you have to deal with sodium chloride (table salt) all the time. After all, table salt is consumed every day by a person with food. And in nature it occurs quite often. But the composition of table salt includes chloride ions, that is, chlorine atoms, which have attached one "extra" electron. One of the reasons for this prevalence of chloride ions is that chlorine atoms have a tendency to attach electrons, that is, when chloride ions are formed from chlorine atoms and electrons, energy is released.
One of the reasons for the release of energy is already known to you - it is associated with an increase in the symmetry of the electron shell of the chlorine atom during the transition to a singly charged anion... At the same time, as you remember, energy 3 p-sublevel decreases. There are other more complex reasons as well.
Due to the fact that the value of the electron affinity energy is influenced by several factors, the nature of the change in this value in the system of elements is much more complex than the nature of the change in the ionization energy. You can be convinced of this by analyzing the table given in Appendix 7. But since the value of this quantity is determined, first of all, by the same electrostatic interaction as the values of the ionization energy, then its change in the system of elements (at least in A- groups), in general terms, is similar to a change in the ionization energy, that is, the energy of affinity for an electron in a group decreases, and in a period it increases. It is maximal for fluorine (328 kJ / mol) and chlorine (349 kJ / mol) atoms. The nature of the change in the energy of affinity for an electron in a system of elements resembles the nature of the change in the ionization energy, that is, the direction of an increase in the energy of affinity for an electron can be schematically shown as follows:
2. On the same scale along the horizontal axis as in the previous tasks, plot the dependence of the molar energy of affinity for an electron on the ordinal number for atoms of elements with Z from 1 to 40 using app 7.
3.What physical meaning have negative values of electron affinity energy?
4. Why, of all atoms of the elements of the 2nd period, only beryllium, nitrogen and neon have negative values of the molar energy of affinity for an electron?
6.15. The tendency of atoms to give and attach electrons
You already know that the propensity of an atom to give up its own and attach foreign electrons depends on its energy characteristics (ionization energy and energy of affinity for an electron). Which atoms are more inclined to donate their electrons, and which ones are more inclined to accept others?
To answer this question, let us summarize in Table 15 everything that we know about the change in these tendencies in the system of elements.
Table 15. Change in the propensity of atoms to give up their own and attach other people's electrons
Electronic configuration of the atom is a formula showing the arrangement of electrons in an atom by levels and sublevels. After studying the article, you will find out where and how electrons are located, get acquainted with quantum numbers and be able to construct the electronic configuration of an atom by its number, at the end of the article there is a table of elements.
Why study the electronic configuration of elements?
Atoms as a constructor: there are a certain number of parts, they differ from each other, but two parts of the same type are exactly the same. But this constructor is much more interesting than the plastic one and here's why. The configuration changes depending on who is nearby. For example, oxygen next to hydrogen maybe turn into water, next to sodium into gas, and being next to iron completely turns it into rust. To answer the question why this is happening and to predict the behavior of an atom next to another, it is necessary to study the electronic configuration, which will be discussed below.
How many electrons are there in an atom?
An atom consists of a nucleus and electrons revolving around it; the nucleus consists of protons and neutrons. In a neutral state, each atom has the same number of electrons as the number of protons in its nucleus. The number of protons was designated by the ordinal number of the element, for example, sulfur, has 16 protons - the 16th element of the periodic table. Gold has 79 protons - 79th element of the periodic table. Accordingly, in sulfur in the neutral state there are 16 electrons, and in gold there are 79 electrons.
Where to look for an electron?
Observing the behavior of the electron, certain patterns were derived, they are described by quantum numbers, there are four of them:
- Principal Quantum Number
- Orbital quantum number
- Magnetic quantum number
- Spin quantum number
Orbital
Further, instead of the word orbit, we will use the term "orbital", the orbital is the wave function of the electron, roughly it is the area in which the electron spends 90% of the time.
N - level
L - shell
M l - orbital number
M s - the first or second electron in the orbital
Orbital quantum number l
As a result of the study of the electron cloud, they found that depending on the energy level, the cloud takes four basic forms: a ball, dumbbells and the other two, more complex. In ascending order of energy, these shapes are called s-, p-, d- and f-shells. Each of these shells can contain 1 (for s), 3 (for p), 5 (for d) and 7 (for f) orbitals. The orbital quantum number is the shell on which the orbitals are located. The orbital quantum number for s, p, d and f-orbitals takes the values 0,1,2 or 3, respectively.
On the s-shell, one orbital (L = 0) - two electrons
There are three orbitals on the p-shell (L = 1) - six electrons
The d-shell has five orbitals (L = 2) - ten electrons
The f-shell has seven orbitals (L = 3) - fourteen electrons
Magnetic quantum number m l
There are three orbitals on the p-shell, they are denoted by numbers from -L to + L, that is, for the p-shell (L = 1) there are orbitals "-1", "0" and "1". The magnetic quantum number is denoted by the letter m l.
Inside the shell, it is easier for electrons to be located in different orbitals, so the first electrons fill one for each orbital, and then a pair of it is attached to each.
Consider a d-shell:
d-shell corresponds to the value L = 2, that is, five orbitals (-2, -1,0,1 and 2), the first five electrons fill the shell taking the values M l = -2, M l = -1, M l = 0 , M l = 1, M l = 2.
Spin quantum number m s
Spin is the direction of rotation of an electron around its axis, there are two directions, so the spin quantum number has two values: +1/2 and -1/2. One energy sublevel can contain two electrons only with opposite spins. The spin quantum number is denoted by m s
Principal quantum number n
The main quantum number is the energy level at this moment seven energy levels are known, each denoted by an Arabic number: 1,2,3, ... 7. The number of shells at each level is equal to the number of the level: at the first level there is one shell, at the second two, etc.
Electron number
So, any electron can be described by four quantum numbers, a combination of these numbers is unique for each position of the electron, take the first electron, the lowest energy level is N = 1, one shell is located at the first level, the first shell at any level has the shape of a ball (s -shell), i.e. L = 0, the magnetic quantum number can take only one value, M l = 0 and the spin will be +1/2. If we take the fifth electron (in whatever atom it is), then the main quantum numbers for it will be: N = 2, L = 1, M = -1, spin 1/2.
The content of the article
ARSENIC- chemical element of group V of the periodic table, belongs to the nitrogen family. The relative atomic mass is 74.9216. In nature, arsenic is represented by only one stable nuclide 75 As. More than ten of its radioactive isotopes with a half-life from several minutes to several months have also been artificially obtained. Typical oxidation states in compounds are –3, +3, +5. The name of arsenic in Russian is associated with the use of its compounds for the extermination of mice and rats; the Latin name Arsenicum comes from the Greek "arsen" - strong, powerful.
Historical information.
Arsenic belongs to the five "alchemical" elements discovered in the Middle Ages (surprisingly, four of them - As, Sb, Bi and P are in the same group of the periodic table - the fifth). At the same time, arsenic compounds have been known since ancient times, they were used for the production of paints and medicines. The use of arsenic in metallurgy is especially interesting.
Several millennia ago, the Stone Age gave way to the Bronze Age. Bronze is an alloy of copper and tin. Historians believe that the first bronze was cast in the Tigris and Euphrates valleys, somewhere between the 30th and 25th centuries. BC. In some regions, bronze with especially valuable properties was smelted - it was cast better and easier forged. As modern scientists have found out, it was an alloy of copper containing from 1 to 7% arsenic and no more than 3% tin. Probably, at first, during its smelting, the rich copper ore malachite was confused with the weathering products of some also green sulfide copper-arsenic minerals. Having appreciated the remarkable properties of the alloy, the ancient craftsmen then specifically looked for arsenic minerals. For searches, we used the property of such minerals to give a specific garlic smell when heated. However, over time, the smelting of arsenic bronze ceased. Most likely this was due to frequent poisoning during the burning of arsenic-containing minerals.
Of course, arsenic was known in the distant past only in the form of its minerals. So, in ancient China, the hard mineral realgar (sulfide of the composition As 4 S 4, realgar in Arabic means "mine dust") was used for stone carving, but when heated or in the light, it "deteriorated", as it turned into As 2 S 3. In the 4th century. BC. Aristotle described this mineral as sandarak. In the 1st century. AD the Roman writer and scientist Pliny the Elder, and the Roman physician and botanist Dioscorides described the mineral orpiment (arsenic sulfide As 2 S 3). Translated from Latin, the name of the mineral means "golden paint": it was used as a yellow dye. In the 11th century. alchemists distinguished three "varieties" of arsenic: the so-called white arsenic (oxide As 2 O 3), yellow arsenic (sulfide As 2 S 3) and red arsenic (sulfide As 4 S 4). White arsenic was obtained by sublimating arsenic impurities during the roasting of copper ores containing this element. Condensing from the gas phase, arsenic oxide was deposited in the form white bloom... White arsenic has been used since ancient times to kill pests, as well as ...
In the 13th century. Albert von Bolstedt (Albert the Great) obtained a metal-like substance by heating yellow arsenic with soap; it may have been the first arsenic specimen in the form simple substance obtained artificially. But this substance broke the mystical "connection" of seven known metals with seven planets; this is probably why alchemists considered arsenic an "illegitimate metal." At the same time, they discovered its property to give copper a white color, which gave rise to call it "a remedy that whitens Venus (that is, copper)."
Arsenic was unequivocally identified as an individual substance in the middle of the 17th century, when the German pharmacist Johann Schroeder obtained it in a relatively pure form by reducing the oxide charcoal... Later, the French chemist and physician Nicola Lemery obtained arsenic by heating a mixture of its oxide with soap and potash. In the 18th century. arsenic was already well known as an unusual "semimetal". In 1775 the Swedish chemist K.V. Scheele obtained arsenic acid and gaseous arsenous hydrogen, and in 1789 A.L. Lavoisier finally recognized arsenic as an independent chemical element. In the 19th century. organic compounds containing arsenic were discovered.
Arsenic in nature.
There is little arsenic in the earth's crust - about 5 · 10 –4% (that is, 5 g per ton), about the same as germanium, tin, molybdenum, tungsten or bromine. Arsenic is often found in minerals together with iron, copper, cobalt, nickel.
The composition of the minerals formed by arsenic (about 200 of them are known) reflects the "semi-metallic" properties of this element, which can be in both positive and negative oxidation states and combine with many elements; in the first case, arsenic can play the role of a metal (for example, in sulfides), in the second, as a non-metal (for example, in arsenides). The complex composition of a number of arsenic minerals reflects its ability, on the one hand, to partially replace sulfur and antimony atoms in the crystal lattice (the ionic radii S –2, Sb –3, and As –3 are close and amount to 0.182, 0.208, and 0.191 nm, respectively), on the other - metal atoms. In the first case, the arsenic atoms have a rather negative oxidation state, in the second - a positive one.
The electronegativity of arsenic (2.0) is small, but higher than that of antimony (1.9) and most metals; therefore, the oxidation state of –3 is observed for arsenic only in metal arsenides, as well as in stibarsen SbAs and intergrowths of this mineral with crystals of pure antimony or arsenic (mineral allemontite). Many compounds of arsenic with metals, judging by their composition, are more likely to be intermetallic compounds than arsenides; some of them have variable arsenic content. In arsenides, several metals can be present simultaneously, the atoms of which, at a close ion radius, replace each other in the crystal lattice in arbitrary proportions; in such cases, in the formula of the mineral, the symbols of the elements are listed separated by commas. All arsenides have a metallic luster, they are opaque, heavy minerals, and their hardness is low.
An example of natural arsenides (about 25 of them are known) are the minerals lellingite FeAs 2 (an analogue of pyrite FeS 2), skutterudite CoAs 2–3 and nickel skutterudite NiAs 2–3, nickeline (red nickel pyrite) NiAs, rammelsbergite (white nickel pyrite) NiAs 2 , safflorite (speiss cobalt) CoAs 2 and clinosafflorite (Co, Fe, Ni) As 2, langisite (Co, Ni) As, sperrylite PtAs 2, maucherite Ni 11 As 8, oregonite Ni 2 FeAs 2, algodonite Cu 6 As. Due to their high density (more than 7 g / cm 3), many of them are classified by geologists as "superheavy" minerals.
The most common arsenic mineral, arsenopyrite (arsenic pyrite) FeAsS, can be considered as a product of the replacement of sulfur in pyrite FeS 2 by arsenic atoms (ordinary pyrite also always contains a little arsenic). Such compounds are called sulfosalts. The minerals cobaltin (cobalt luster) CoAsS, glaucodot (Co, Fe) AsS, gersdorfite (nickel luster) NiAsS, enargite and luconite of the same composition, but different structure Cu 3 AsS 4, proustite Ag 3 AsS 3 is an important silver ore sometimes called "ruby silver" because of its bright red color, it is often found in upper layers silver veins, where magnificent large crystals of this mineral have been found. Sulfosalts can also contain noble metals of the platinum group; These are minerals osarsite (Os, Ru) AsS, ruarsite RuAsS, irarsite (Ir, Ru, Rh, Pt) AsS, platarsite (Pt, Rh, Ru) AsS, hollingworthite (Rd, Pt, Pd) AsS. Sometimes the role of sulfur atoms in such double arsenides is played by antimony atoms, for example, in seinayokite (Fe, Ni) (Sb, As) 2, arsenopalladinite Pd 8 (As, Sb) 3, arsenpolybasite (Ag, Cu) 16 (Ar, Sb) 2 S 11.
The structure of minerals is interesting, in which arsenic is present simultaneously with sulfur, but rather plays the role of a metal, grouping together with other metals. These are the minerals arsenosulvanite Cu 3 (As, V) S 4, arsenogauhecornite Ni 9 BiAsS 8, freibergite (Ag, Cu, Fe) 12 (Sb, As) 4 S 13, tennantite (Cu, Fe) 12 As 4 S 13, argentothennantite (Ag, Cu) 10 (Zn, Fe) 2 (As, Sb) 4 S 13, goldfieldite Cu 12 (Te, Sb, As) 4 S 13, girodite (Cu, Zn, Ag) 12 (As, Sb) 4 (Se, S) 13. You can imagine how complex the structure has crystal cell all these minerals.
Arsenic has a uniquely positive oxidation state in natural sulfides - yellow orpiment As 2 S 3, orange-yellow dimorphite As 4 S 3, orange-red realgar As 4 S 4, carmine-red getcellite AsSbS 3, as well as in colorless oxide As 2 O 3, which occurs in the form of arsenolite and claudetite minerals with different crystal structures (they are formed as a result of weathering of other arsenic minerals). Usually these minerals are found in small blotches. But in the 30s of the 20th century. In the southern part of the Verkhoyansk ridge, huge crystals of oripigment were found, up to 60 cm in size and weighing up to 30 kg.
In natural salts of arsenic acid H 3 AsO 4 - arsenates (about 90 of them are known), the oxidation state of arsenic is +5; an example is the bright pink erythrine (cobalt color) Co 3 (AsO 4) 2 8H 2 O, green annabergite Ni 3 (AsO 4) 2 8H 2 O, scorodite Fe III AsO 4 2H 2 O and simplesite Fe II 3 (AsO 4) 2 8H 2 O, brown-red gasparite (Ce, La, Nd) ArO 4, colorless guernesite Mg 3 (AsO 4) 2 8H 2 O, Rooseveltite BiAsO 4 and cattigite Zn 3 (AsO 4) 2 · 8H 2 O, as well as many basic salts, for example, olivienite Cu 2 AsO 4 (OH), arsenobismite Bi 2 (AsO 4) (OH) 3. But natural arsenites - derivatives of arsenous acid H 3 AsO 3 are very rare.
In central Sweden, there are the famous Langban iron-manganese quarries, in which more than 50 samples of arsenate minerals have been found and described. Some of them are not found anywhere else. They were formed once as a result of the reaction of arsenic acid H 3 AsO 4 with pyrocroite Mn (OH) 2 at not very high temperatures. Usually, arsenates are oxidation products of sulfide ores. They usually have no industrial application, but some of them are very beautiful and adorn mineralogical collections.
In the names of numerous arsenic minerals one can find place names (Lölling in Austria, Freiberg in Saxony, Seinäjoki in Finland, Skutterud in Norway, Allemont in France, the Canadian Langis mine and the Getchell mine in Nevada, Oregon in the USA, etc.), the names of geologists, chemists, politicians, etc. (German chemist Karl Rammelsberg, Munich mineral trader William Maucher, mine owner Johann von Gersdorff, French chemist F. Claudet, English chemists John Proust and Smithson Tennant, Canadian chemist F.L. Perry, US President Roosevelt, etc.), plant names (for example, the name of the safflorite mineral comes from saffron), the initial letters of the names of the elements - arsenic, osmium, ruthenium, iridium, palladium, platinum, Greek roots ("erythros" - red, "enargon" - visible, "lithos" - stone) and etc. etc.
An interesting old name for the nickel mineral (NiAs) is kupfernickel. Medieval German miners called the evil mountain spirit Nickel, and "Kupfernickel" (from German Kupfer - copper) - "damn copper", "fake copper". The copper-red crystals of this ore looked very much like copper ore; it was used in glass making to color glasses green. But no one was able to get copper from it. This ore was investigated in 1751 by the Swedish mineralogist Axel Kronstedt and isolated a new metal from it, calling it nickel.
Since arsenic is chemically inert enough, it also occurs in its native state - in the form of fused needles or cubes. Such arsenic usually contains from 2 to 16% impurities - most often these are Sb, Bi, Ag, Fe, Ni, Co. It is easy to grind it into powder. In Russia, geologists found native arsenic in Transbaikalia, in the Amur Region, and it is also found in other countries.
Arsenic is unique in that it is found everywhere - in minerals, rocks, soil, water, plants and animals, it is not for nothing that it is called "omnipresent." Arsenic distribution in different regions the globe was largely determined in the processes of formation of the lithosphere by the volatility of its compounds at high temperature, as well as the processes of sorption and desorption in soils and sedimentary rocks. Arsenic easily migrates, which is facilitated by the rather high solubility of some of its compounds in water. In humid climates, arsenic is washed out of the soil and carried away groundwater, and then - rivers. The average content of arsenic in rivers is 3 μg / l, in surface waters- about 10 μg / l, in the water of the seas and oceans - only about 1 μg / l. This is due to the relatively rapid deposition of its compounds from water with accumulation in bottom sediments, for example, in ferromanganese nodules.
In soils, the arsenic content is usually from 0.1 to 40 mg / kg. But in the area of occurrence of arsenic ores, as well as in volcanic regions, the soil can contain a lot of arsenic - up to 8 g / kg, as in some areas of Switzerland and New Zealand. In such places, vegetation dies, and animals get sick. This is typical for steppes and deserts, where arsenic is not washed out of the soil. Clay rocks are also enriched in comparison with the average content - they contain four times more arsenic than the average. In our country, the maximum permissible concentration of arsenic in soil is 2 mg / kg.
Arsenic can be removed from the soil not only by water, but also by the wind. But for this, it must first turn into volatile organo arsenic compounds. This transformation occurs as a result of the so-called biomethylation - the addition of a methyl group with the formation of a C – As bond; this enzymatic process (it is well known for mercury compounds) occurs with the participation of the coenzyme methylcobalamin, a methylated derivative of vitamin B 12 (it is also present in the human body). Biomethylation of arsenic occurs in both fresh and seawater and leads to the formation of organo arsenic compounds - methylarsonic acid CH 3 AsO (OH) 2, dimethylarsic (dimethyl arsenic, or cacodylic) acid (CH 3) 2 As (O) OH, trimethylarsine ( CH 3) 3 As and its oxide (CH 3) 3 As = O, which are also found in nature. Using 14 C-labeled methylcobalamin and 74 As-labeled sodium hydrogen arsenate Na 2 HAsO 4, it was shown that one of the methanobacterial strains reduces and methylates this salt to volatile dimethylarsine. As a result, the air in rural areas contains an average of 0.001 - 0.01 μg / m 3 of arsenic, in cities where there are no specific pollution - up to 0.03 μg / m working on coal with a high arsenic content, etc.), the concentration of arsenic in the air can exceed 1 μg / m 3. The intensity of the fallout of arsenic in the areas where industrial centers are located is 40 kg / km 2 per year.
The formation of volatile arsenic compounds (trimethylarsine, for example, boils at only 51 ° C) caused in the 19th century. numerous poisonings, since arsenic was contained in plaster and even in green wallpaper paint. Scheele greens were previously used as paint Cu 3 (AsO 3) 2 · n H 2 O and Parisian or Swiss greens Cu 4 (AsO 2) 6 (CH 3 COO) 2. In conditions high humidity and the appearance of mold, volatile organo arsenic derivatives are formed from such paint. It is believed that this process could be the reason for the slow poisoning of Napoleon in last years his life (as you know, arsenic was found in Napoleon's hair a century and a half after his death).
Arsenic is found in significant quantities in some mineral waters. Russian standards establish that arsenic in medicinal table mineral waters should not exceed 700 μg / l. V Jermuk it can be several times more. One or two glasses of "arsenic" drunk mineral water They will not bring harm to a person: in order to be fatally poisoned, one must drink three hundred liters at once ... But it is clear that such water cannot be drunk constantly instead of ordinary water.
Chemists have found out that arsenic in natural waters it can be in different forms, which is important from the point of view of its analysis, migration methods, as well as different toxicity of these compounds; thus, compounds of trivalent arsenic are 25–60 times more toxic than pentavalent arsenic. As (III) compounds in water are usually present in the form of weak arsenous acid H 3 AsO 3 ( pK a = 9.22), and the As (V) compounds - in the form of a much stronger arsenic acid H 3 AsO 4 ( pK a = 2.20) and its deprotonated anions H 2 AsO 4 - and HAsO 4 2–.
The living matter of arsenic contains on average 6 · 10 –6%, that is, 6 µg / kg. Some algae are capable of concentrating arsenic to such an extent that they are dangerous to humans. Moreover, these algae can grow and multiply in pure solutions of arsenous acid. Such algae are used in some Asian countries as a rat control agent. Even in clean waters Norwegian fjords algae can contain up to 0.1 g / kg arsenic. In humans, arsenic is found in brain tissue and muscles, it accumulates in hair and nails.
Arsenic properties.
Although arsenic looks like a metal, it is still rather a non-metal: it does not form salts, for example, with sulfuric acid, but itself is an acid-forming element. Therefore, this element is often referred to as a semimetal. Arsenic exists in several allotropic forms and, in this respect, closely resembles phosphorus. The most stable of them is gray arsenic, a very fragile substance that has a metallic luster when it is freshly fractured (hence the name "metallic arsenic"); its density is 5.78 g / cm 3. On strong heating (up to 615 ° C), it sublimes without melting (the same behavior is characteristic of iodine). Under a pressure of 3.7 MPa (37 atm), arsenic melts at 817 ° C, which is significantly higher than the sublimation temperature. The electrical conductivity of gray arsenic is 17 times less than that of copper, but 3.6 times higher than that of mercury. As the temperature rises, its electrical conductivity, like that of typical metals, decreases - approximately to the same extent as that of copper.
If arsenic vapors are very quickly cooled to the temperature of liquid nitrogen (–196 ° C), a transparent soft yellow substance resembling yellow phosphorus is obtained, its density (2.03 g / cm 3) is much lower than that of gray arsenic. Vapors of arsenic and yellow arsenic are composed of tetrahedral As 4 molecules - and here is an analogy with phosphorus. At 800 ° C, a noticeable dissociation of vapors begins with the formation of As 2 dimers, and at 1700 ° C only As 2 molecules remain. When heated and under the influence of ultraviolet radiation, yellow arsenic quickly turns into gray with the release of heat. When arsenic vapors condense in an inert atmosphere, another amorphous form of this black element is formed. If arsenic vapors are deposited on glass, a mirror film is formed.
The structure of the outer electron shell of arsenic is the same as that of nitrogen and phosphorus, but unlike them, it has 18 electrons on the penultimate shell. Like phosphorus, it can form three covalent bonds (configuration 4s 2 4p 3), and a lone pair remains on the As atom. The sign of the charge on the As atom in compounds with covalent bonds depends on the electronegativity of neighboring atoms. The participation of the lone pair in complexation for arsenic is much more difficult in comparison with nitrogen and phosphorus.
If d-orbitals are involved in the As atom, it is possible for the 4s-electrons to unpair with the formation of five covalent bonds. This possibility is practically realized only in combination with fluorine - in pentafluoride AsF 5 (pentachloryl AsCl 5 is also known, but it is extremely unstable and decomposes quickly even at –50 ° C).
In dry air, arsenic is stable, but in humid air it fades and becomes covered with black oxide. Upon sublimation, arsenic vapors easily burn in air with a blue flame with the formation of heavy white vapors of arsenous anhydride As 2 O 3. This oxide is one of the most common arsenic-containing reagents. It has amphoteric properties:
As 2 O 3 + 6HCl ® 2AsCl 3 + 3H 2 O,
2 O 3 + 6NH 4 OH ® 2 (NH 4) 3 AsO 3 + 3H 2 O.
During the oxidation of As 2 O 3, an acidic oxide is formed - arsenic anhydride:
As 2 O 3 + 2HNO 3 ® As 2 O 5 + H 2 O + NO 2 + NO.
When it interacts with soda, sodium hydrogen arsenate is obtained, which is used in medicine:
As 2 O 3 + 2Na 2 CO 3 + H 2 O ® 2Na 2 HAsO 4 + 2CO 2.
Pure arsenic is fairly inert; water, alkalis and acids, which do not have oxidizing properties, do not affect it. Diluted nitric acid oxidizes it to ortho-arsenous acid H 3 AsO 3, and concentrated nitric acid to ortho-arsenic acid H 3 AsO 4:
3As + 5HNO 3 + 2H 2 O ® 3H 3 AsO 4 + 5NO.
Arsenic (III) oxide reacts similarly:
3As 2 O 3 + 4HNO 3 + 7H 2 O ® 6H 3 AsO 4 + 4NO.
Arsenic acid is a medium strength acid, slightly weaker than phosphoric acid. In contrast, arsenous acid is very weak, corresponding in strength to boric acid H 3 BO 3. In its solutions, there is an equilibrium H 3 AsO 3 HAsO 2 + H 2 O. Arsenous acid and its salts (arsenites) are strong reducing agents:
HAsO 2 + I 2 + 2H 2 O ® H 3 AsO 4 + 2HI.
Arsenic reacts with halogens and sulfur. Chloride AsCl 3 is a colorless oily liquid fuming in air; hydrolyzed with water: AsCl 3 + 2H 2 O ® HAsO 2 + 3HCl. Known bromide AsBr 3 and iodide AsI 3, which are also decomposed by water. In the reactions of arsenic with sulfur, sulfides of various compositions are formed - up to Ar 2 S 5. Arsenic sulfides dissolve in alkalis, in ammonium sulfide solution and in concentrated nitric acid, for example:
As 2 S 3 + 6KOH ® K 3 AsO 3 + K 3 AsS 3 + 3H 2 O,
2 S 3 + 3 (NH 4) 2 S ® 2 (NH 4) 3 AsS 3,
2 S 5 + 3 (NH 4) 2 S ® 2 (NH 4) 3 AsS 4,
As 2 S 5 + 40HNO 3 + 4H 2 O ® 6H 2 AsO 4 + 15H 2 SO 4 + 40NO.
In these reactions, thioarsenites and thioarsenates are formed - salts of the corresponding thioacids (analogous to thiosulfuric acid).
In the reaction of arsenic with active metals, salt-like arsenides are formed, which are hydrolyzed by water. The reaction proceeds especially rapidly in an acidic medium with the formation of arsine: Ca 3 As 2 + 6HCl ® 3CaCl 2 + 2AsH 3. Arsenides of low-activity metals - GaAs, InAs, etc. have a diamond-like atomic lattice. Arsine is a colorless, very poisonous, odorless gas, but impurities give it the smell of garlic. Arsine slowly decomposes into elements even at room temperature and quickly when heated.
Arsenic forms many organic arsenic compounds, for example, tetramethyldiarsine (CH 3) 2 As – As (CH 3) 2. Back in 1760, Louis Claude Cadet de Gassicourt, director of the Servian porcelain factory, while distilling potassium acetate with arsenic (III) oxide, unexpectedly obtained a fuming liquid containing arsenic with a disgusting odor, which was called alarsine, or Cadet's liquid. As it turned out later, this liquid contained the first organic derivatives of arsenic obtained in this liquid: the so-called cacodyl oxide, which was formed as a result of the reaction
4CH 3 COOK + As 2 O 3 ® (CH 3) 2 As – O – As (CH 3) 2 + 2K 2 CO 3 + 2CO 2, and dicacodyl (CH 3) 2 As – As (CH 3) 2. Kakodyl (from the Greek "kakos" - bad) was one of the first radicals discovered in organic compounds.
In 1854, the Parisian professor of chemistry Auguste Kaur synthesized trimethylarsine by the action of methyl iodide on sodium arsenide: 3CH 3 I + AsNa 3 ® (CH 3) 3 As + 3NaI.
Subsequently, arsenic trichloride was used for syntheses, for example,
(CH 3) 2 Zn + 2AsCl 3 ® 2 (CH 3) 3 As + 3ZnCl 2.
In 1882, aromatic arsines were obtained by the action of metallic sodium on a mixture of aryl halides and arsenic trichloride: 3C 6 H 5 Cl + AsCl 3 + 6Na ® (C 6 H 5) 3 As + 6NaCl. The chemistry of organic arsenic derivatives developed most intensively in the 1920s, when antimicrobial, as well as irritating and skin-blistering effects were found in some of them. Currently, tens of thousands of organoarsenic compounds have been synthesized.
Obtaining arsenic.
Arsenic is obtained mainly as a by-product of the processing of copper, lead, zinc and cobalt ores, as well as in the extraction of gold. Some polymetallic ores contain up to 12% arsenic. When such ores are heated to 650–700 ° C in the absence of air, arsenic sublimes, and when heated in air, a volatile oxide As 2 O 3 - "white arsenic" is formed. It is condensed and heated with coal, and arsenic is reduced. Obtaining arsenic is hazardous production. Earlier, when the word "ecology" was known only to narrow specialists, "white arsenic" was released into the atmosphere, and it settled on neighboring fields and forests. The exhaust gases of arsenic plants contain from 20 to 250 mg / m 3 As 2 O 3, while usually the air contains about 0.00001 mg / m 3. The average daily allowable concentration of arsenic in the air is considered to be only 0.003 mg / m 3. Paradoxically, even now, it is not factories producing arsenic that pollute the environment much more strongly, but non-ferrous metallurgy enterprises and power plants that burn coal. Bottom sediments near copper smelters contain a huge amount of arsenic - up to 10 g / kg. Arsenic can also get into the soil with phosphorus fertilizers.
And one more paradox: they receive more arsenic than is required; it's pretty rare case... In Sweden, "unnecessary" arsenic was even forced to be buried in reinforced concrete containers in deep abandoned mines.
The main industrial mineral of arsenic is arsenopyrite FeAsS. There are large copper-arsenic deposits in Georgia, Central Asia and Kazakhstan, in the USA, Sweden, Norway and Japan, arsenic-cobalt deposits in Canada, arsenic-tin deposits in Bolivia and England. In addition, there are known gold-arsenic deposits in the USA and France. Russia has numerous arsenic deposits in Yakutia, the Urals, Siberia, Transbaikalia and Chukotka.
Determination of arsenic.
A qualitative reaction to arsenic is the precipitation of yellow sulfide As 2 S 3 from hydrochloric acid solutions. Traces are determined by the Marsh reaction or by the Gutzeit method: strips of paper moistened with HgCl 2 darken in the presence of arsine, which reduces mercury to mercury.
In recent decades, various sensitive analytical methods have been developed, with the help of which it is possible to quantitatively determine negligible concentrations of arsenic, for example, in natural waters. These include flame atomic absorption spectrometry, atomic emission spectrometry, mass spectrometry, atomic fluorescence spectrometry, neutron activation analysis ... If there is very little arsenic in the water, it may be necessary to preconcentrate the samples. Using this concentration, a group of Kharkov scientists from the National Academy of Sciences of Ukraine developed in 1999 an extraction-X-ray fluorescence method for the determination of arsenic (as well as selenium) in drinking water with a sensitivity of up to 2.5–5 µg / L.
For the separate determination of As (III) and As (V) compounds, they are preliminarily separated from each other using well-known extraction and chromatographic methods, as well as using selective hydrogenation. Extraction is usually carried out with sodium dithiocarbamate or ammonium pyrrolidinedithiocarbamate. These compounds form water-insoluble complexes with As (III), which can be removed with chloroform. Then using oxidation nitric acid arsenic can be transferred back to the aqueous phase. In the second sample, arsenate is converted into arsenite with the help of a reducing agent, and then a similar extraction is performed. This is how "total arsenic" is determined, and then, by subtracting the first result from the second, As (III) and As (V) are determined separately. If there are organic arsenic compounds in the water, they are usually converted into methyldiiodarsine CH 3 AsI 2 or dimethyliodarsine (CH 3) 2 AsI, which are determined by one or another chromatographic method. So, using high performance liquid chromatography, nanogram quantities of a substance can be determined.
Many arsenic compounds can be analyzed by the so-called hydride method. It consists in selective reduction of the analyte into volatile arsine. Thus, inorganic arsenites are reduced to AsH 3 at pH 5 - 7, and at pH
The neutron activation method is also sensitive. It consists in irradiating a sample with neutrons, while 75 As nuclei capture neutrons and turn into 76 As radionuclide, which is detected by characteristic radioactivity with a half-life of 26 hours. In this way, up to 10–10% of arsenic can be detected in the sample; 1 mg per 1000 tons of substance
Arsenic use.
About 97% of mined arsenic is used in the form of its compounds. Pure arsenic is rarely used. Only a few hundred tons of metallic arsenic are produced and used a year all over the world. Arsenic in the amount of 3% improves the quality of bearing alloys. The addition of arsenic to lead significantly increases its hardness, which is used in the production of lead-acid batteries and cables. Small additions of arsenic increase corrosion resistance and improve the thermal properties of copper and brass. Highly purified arsenic is used in the manufacture of semiconductor devices, in which it is fused with silicon or germanium. Arsenic is also used as a dopant, which imparts a certain type of conductivity to "classical" semiconductors (Si, Ge).
Arsenic is also used as a valuable additive in nonferrous metallurgy. So, adding 0.2 ... 1% As to lead significantly increases its hardness. It has long been noticed that if a little arsenic is added to the molten lead, then when casting the shot, balls of the correct spherical shape are obtained. The addition of 0.15 ... 0.45% of arsenic to copper increases its tensile strength, hardness and corrosion resistance when working in a gas-polluted environment. In addition, arsenic increases the flowability of copper during casting, facilitates the wire drawing process. Arsenic is added to some types of bronzes, brasses, babbits, printing alloys. And at the same time, arsenic is very often harmful to metallurgists. In the production of steel and many non-ferrous metals, they deliberately make the process more complicated in order to remove all arsenic from the metal. The presence of arsenic in the ore makes the production harmful. Harmful twice: first, for human health; secondly, for metal - significant impurities of arsenic deteriorate the properties of almost all metals and alloys.
Various arsenic compounds are more widely used, which are produced annually in tens of thousands of tons. Oxide As 2 O 3 is used in glass making as a glass clarifier. Even ancient glassmakers knew that white arsenic makes glass "deaf", that is, opaque. However, small additions of this substance, on the contrary, brighten the glass. Arsenic is still included in the formulations of some glasses, for example, "Viennese" glass for thermometers.
Arsenic compounds are used as an antiseptic to prevent deterioration and preservation of hides, furs and stuffed animals, to impregnate wood, as a component of antifouling paints for the bottoms of ships. In this capacity, salts of arsenic and arsenous acids are used: Na 2 HAsO 4, PbHAsO 4, Ca 3 (AsO 3) 2, etc. The biological activity of arsenic derivatives interested veterinarians, agronomists, and sanitary epidemiological service specialists. As a result, arsenic-containing stimulants of the growth and productivity of livestock, antihelminthic agents, drugs for the prevention of diseases of young animals on livestock farms appeared. Arsenic compounds (As 2 O 3, Ca 3 As 2, Na 3 As, Parisian greens) are used to control insects, rodents, and weeds. Previously, this application was widespread, especially in the treatment of fruit trees, tobacco and cotton plantations, to rid livestock of lice and fleas, to stimulate growth in poultry and pig breeding, and to dry cotton before harvesting. Even in ancient China, rice crops were treated with arsenic oxide in order to protect them from rats and fungal diseases and thus increase the yield. And in South Vietnam, American troops used cacodylic acid (Agent Blue) as a defoliant. Now, due to the toxicity of arsenic compounds, their use in agriculture is limited.
Important areas of application of arsenic compounds are the production of semiconductor materials and microcircuits, fiber optics, growing single crystals for lasers, and film electronics. For the introduction of small, strictly dosed amounts of this element into semiconductors, gaseous arsine is used. Gallium arsenides GaAs and indium InAs are used in the manufacture of diodes, transistors, and lasers.
Arsenic also finds limited use in medicine. . Arsenic isotopes 72 As, 74 As, and 76 As with convenient research half-lives (26 h, 17.8 days, and 26.3 h, respectively) are used to diagnose various diseases.
Ilya Leenson
