The periodic table of chemical elements is a graphical expression of the periodic law. Group. the covalent nature of the bond is enhanced. Sodium chloride solution
Group a set of elements with the same oxygen valence is called. This highest valency is determined by the group number. Since the sum of the highest oxygen valences for non-metallic elements is eight, it is easy to determine the formula of the higher hydrogen compound by the group number. So, for phosphorus- an element of the fifth group - the highest valence for oxygen is five, the formula for the higher oxide is P2O5, and the formula for the compound with hydrogen is PH3. For sulfur, an element of the sixth group, the formula of the higher oxide is SO3, and the higher compound with hydrogen is H2S.
Some elements have a higher valency that is not equal to the number of their groups. Such exceptions are copper Cu, silver Ag, gold Au... They are in the first group, but their valencies vary from 1 to 3. For example, there are compounds: CuO; AgO; Cu2O3; Au2O3. Oxygen placed in the sixth group, although its compounds with a valency higher than two are almost never found. Fluorine P - an element of group VII - is monovalent in its most important compounds; bromine Br - an element of group VII - is maximally pentavalent. There are especially many exceptions in the VIII group. There are only two elements in it: ruthenium Ru and osmium Os exhibit a valency equal to eight, their higher oxides have the formulas RuO4 and OsO4 The valence of the other elements of group VIII is much lower.
The elements of the group are distributed by subgroups... The subgroup unites elements of this group, which are more similar in their chemical properties. The similarity depends on the analogy in the structure electronic shells atoms of elements. In the periodic table, the symbols of the elements of each of the subgroups are placed strictly vertically.
The first seven groups have one main and one secondary subgroup; in the eighth group there is one main subgroup, "inert" elements, and three secondary ones. The name of each subgroup is usually given by the name of the top element, for example: lithium subgroup (Li-Na-K-Rb-Cs-Fr), chromium subgroup (Cr-Mo-W) While elements of the same subgroup are chemical analogs, elements of different subgroups of the same group sometimes differ very sharply in their properties. A common property for elements of the main and secondary subgroups of the same group is basically only their identical highest valence for oxygen. So, manganese Mn and chlorine C1, which are in different subgroups of group VII, chemically have almost nothing in common: manganese is a metal, chlorine is a typical non-metal. However, the formulas of their higher oxides and the corresponding hydroxides are similar: Mn2O7 - Cl2O7; НМnО4 - НС1О4.
In the periodic table, there are two horizontal rows of 14 elements, located outside the groups. They are usually placed at the bottom of the table. One of these series is composed of elements called lanthanides (literally: similar to lanthanum), the other series - elements of actinides (similar to anemones). The actinide symbols are located below the lanthanide symbols. This arrangement reveals 14 shorter subgroups consisting of 2 elements each: these are the second side, or lanthanoid-actinoid subgroups.
On the basis of all that has been said, there are: a) main subgroups, b) side subgroups, and c) second side (lanthanoid-actinoid) subgroups.
It should be noted that some of the main subgroups also differ from each other in the structure of the atoms of their elements. Based on this, all subgroups periodic system can be divided into 4 categories.
I. Main subgroups of groups I and II (subgroups of lithium and beryllium).
II. Six main subgroups III - IV - V - VI - VII - VIII of groups (subgroups of boron, carbon, nitrogen, oxygen, fluorine and neon).
III. Ten side subgroups(one each in groups I-VII and three in group VIII). Jfc,
IV. Fourteen lanthanoid-actinoid subgroups.
The numbers of subgroups of these 4 categories make up an arithmetic progression: 2-6-10-14.
It should be noted that the top element of any main subgroup is in period 2; any secondary upper element - in the 4th period; the top element of any lanthanoid-actinoid subgroup - in the 6th period. Thus, with each new even period of the periodic system, new categories of subgroups appear.
Each element, except for being in one or another group and subgroup, is in another of the seven periods.
A period is a sequence of elements during which their properties change in an order of gradual enhancement from typically metallic to typically non-metallic (metalloid). Each period ends with an inert element. As the metallic properties weaken, non-metallic properties begin to appear and gradually increase in the elements; in the middle of the periods, there are usually elements that combine, to one degree or another, both metallic and non metallic properties... These elements are often referred to as amphoteric.
The composition of the periods.
The periods are not uniform in terms of the number of elements included in them. The first three are called small, the other four are large.
The number of elements in any period is expressed by the formula 2n2 where n is an integer. In periods 2 and 3, there are 8 elements each; 4 and 5 - 18 elements each; 6-32 elements; in 7, not yet completed, so far 18, elements, although theoretically there should also be 32 elements.
1 period is original. It contains only two elements: hydrogen H and helium He. The transition of properties from metallic to non-metallic occurs: here in one typical amphoteric element - hydrogen. The latter, according to its inherent metallic properties, heads the subgroup alkaline metals, according to its inherent non-metallic properties - a subgroup halogens... Therefore, hydrogen is often placed twice in the periodic table - in groups 1 and VII.
The different quantitative composition of the periods leads to an important consequence: neighboring elements of small periods, for example, carbon C and nitrogen N, comparatively sharply differ from each other in their properties: neighboring elements of long periods, for example, lead Pb and bismuth Bi, are much closer in properties to each other. friend, since the change in the nature of the elements in large periods occurs in small jumps. In some areas of large periods, there is even such a slow decline in metallicity that adjacent elements are very similar in their chemical properties. Such, for example, is the triad of elements of the fourth period: iron Fe - cobalt Co - nickel Ni, which is often called the "iron family". The horizontal similarity (horizontal analogy) overrides even vertical similarity (vertical analogy); so, the elements of the iron subgroup - iron, ruthenium, osmium - are less chemically similar to each other than the elements of the "iron family".
The lanthanides are the most striking example of a horizontal analogy. All of them are chemically similar to each other and to lanthanum La. In nature, they are found in companies, are difficult to separate, the typical highest valence of most of them is 3. The lanthanides have a special internal periodicity: every eighth of them, in order of arrangement, repeats to some extent the properties and valence states of the first, i.e. the one from which the countdown begins. Thus, terbium Tb is similar to cerium Ce; lutetium Lu - to gadolinium Gd.
Actinides are similar to lanthanides, but their horizontal analogy is manifested to a much lesser extent. The highest valence of some actinides (for example, uranium U) reaches six. The internal periodicity, which is possible in principle and among them, has not yet been confirmed.
Arrangement of elements in the periodic table. Moseley's Law.
DI Mendeleev arranged the elements in a certain sequence, sometimes called the "Mendeleev series." In general, this sequence (numbering) is associated with an increase in the atomic masses of elements. However, there are exceptions. Sometimes the logical course of change in valence is in conflict with the course of change in atomic masses In such cases, the need demanded to give preference to any one of these two bases of systematization.DI Mendeleev in some cases violated the principle of arrangement of elements with increasing atomic masses and relied on a chemical analogy between the elements. Co, iodine I before tellurium Te, then these elements would fall into subgroups and groups that do not correspond to their properties and their highest valency.
In 1913, the English scientist G. Moseley, studying the spectra of X-rays for various elements, noticed a pattern connecting the numbers of elements in the periodic system of Mendeleev with the wavelengths of these rays, resulting from irradiation of certain elements with cathode clouds. It turned out that the square roots of the reciprocal values of the wavelengths of these rays are linearly related to the serial numbers of the corresponding elements. H. Moseley's law made it possible to check the correctness of the "Mendeleev series" and confirmed its impeccability.
For example, let us know the values for elements No. 20 and No. 30, the numbers of which in the system do not cause any doubts in our minds. These values are linearly related to the indicated numbers. To check, for example, the correctness of the number assigned to cobalt (27), and judging by the atomic mass, this number should have had nickel, it is irradiated with cathode rays: as a result, X-rays are released from the cobalt. By decomposing them on suitable diffraction gratings (on crystals), we obtain the spectrum of these rays and, choosing the clearest of the spectral lines, we measure the wavelength () of the ray corresponding to this line; then we postpone the value on the ordinate. From the resulting point A, draw a straight line parallel to the abscissa axis, until it intersects with the previously identified straight line. From the intersection point B, we lower the perpendicular to the abscissa axis: it will accurately indicate to us the cobalt number, equal to 27. The modern formulation of the periodic law. The physical meaning of the serial number of the element.
After the works of G. Moseley, the atomic mass of an element gradually began to give way to its leading role to a new, not yet clear in its internal (physical) meaning, but clearer constant - the ordinal or, as is now called, the atomic number of the element. The physical meaning of this constant was revealed in 1920 by the works of the English scientist D. Chadwick. D. Chadwick experimentally established that the ordinal number of an element is numerically equal to the value of the positive charge Z of the nucleus of the atom of this element, that is, to the number of protons in the nucleus. It turned out that D.I. Mendeleev, without suspecting it, arranged the elements in a sequence that exactly corresponds to the increase in the charge of the nuclei of their atoms.
By the same time, it was also established that the atoms of the same element can differ from each other in their mass; such atoms are called isotopes. An example is the atoms: and. In the periodic table, isotopes of the same element occupy one cell. In connection with the discovery of isotopes, the concept of a chemical element was clarified. At present, a chemical element is called the kind of atoms that have the same nuclear charge - the same number of protons in the nucleus. The wording of the periodic law was also clarified. The modern formulation of the law says: the properties of elements and their compounds are periodically dependent on the size and charge of the nuclei of their atoms.
Other characteristics of elements associated with the structure of the outer electron layers of atoms, atomic volumes, ionization energy, and other properties also change periodically.
Periodic table and the structure of the electron shells of atoms of elements.
Later it was found that not only the ordinal number of the element has a deep physical meaning, but also other concepts previously considered earlier also gradually acquired physical meaning. For example, the group number, indicating the highest valence of an element, thereby reveals the maximum number of electrons of an atom of an element that can participate in the formation of a chemical bond.
The period number, in turn, turned out to be related to the number of energy levels available in the electron shell of an atom of an element of a given period.
Thus, for example, the "coordinates" of tin Sn (serial number 50, period 5, main subgroup of group IV) mean that there are 50 electrons in the tin atom, they are distributed over 5 energy levels, only 4 electrons are valence.
The physical meaning of finding elements in subgroups of different categories is extremely important. It turns out that for elements located in subgroups of category I, the next (last) electron is located on the s-sublevel of the external level. These elements belong to the electronic family. For atoms of elements located in subgroups of category II, the next electron is located at the p-sublevel of the external level. These are elements of the electronic family “p.” So, the next 50th electron of tin atoms is located on the p-sublevel of the external, that is, the 5th energy level.
In the atoms of the elements of subgroups of the III category, the next electron is located on the d-sublevel, but already before the outer level, these are elements of the electronic family "d". In the atoms of lanthanides and actinides, the next electron is located at the f-sublevel, before the outer level. These are elements of the "f" electronic family.
It is no coincidence, therefore, that the above-mentioned numbers of subgroups of these 4 categories, that is, 2-6-10-14, coincide with the maximum numbers of electrons on the s-p-d-f sublevels.
But it turns out that it is possible to solve the question of the order of filling the electron shell and derive an electronic formula for an atom of any element and on the basis of the periodic system, which indicates with sufficient clarity the level and sublevel of each successive electron. The periodic table also indicates the placement of one after another of the elements by periods, groups, subgroups and the distribution of their electrons by levels and sublevels, because each element has its own corresponding last electron that characterizes it. As an example, let us analyze the compilation of an electronic formula for an atom of the element zirconium (Zr). The periodic table gives indicators and "coordinates" of this element: serial number 40, period 5, group IV, side subgroup. First conclusions: a) all electrons 40, b) these 40 electrons are distributed at five energy levels; c) out of 40 electrons only 4 are valence, d) the next 40th electron entered the d-sublevel before the external, ie, the fourth energy level.Similar conclusions can be drawn about each of the 39 elements preceding zirconium, only the indicators and coordinates will be different each time.
Therefore, the methodological method of drawing up the electronic formulas of elements on the basis of the periodic system consists in the fact that we sequentially consider the electron shell of each element along the path to the given one, identifying by its "coordinates" where the next electron went in the shell.
The first two elements of the first period, hydrogen H and helium, do not belong to the s-family. Their electrons, including two, enter the s-sublevel of the first level. We write down: The first period ends here, the first energy level also. The next two elements of the second period - lithium Li and beryllium Be - are in the main subgroups of groups I and II. They are also s-elements. Their next electrons will be located on the s sublevel of the 2nd level. We write Next 6 elements of the 2nd period follow in a row: boron B, carbon C, nitrogen N, oxygen O, fluorine F and neon Ne. According to the location of these elements in the main subgroups III - Vl groups, their next electrons among six will be located on the p-sublevel of the 2nd level. We write down: The second period ends with an inert element with neon, the second energy level is also over. This is followed by two elements of the third period of the main subgroups of groups I and II: sodium Na and magnesium Mg. These are s-elements and their next electrons are located on the s-sublevel of the 3rd level. Then there are six elements of the 3rd period: aluminum Al, silicon Si, phosphorus P, sulfur S, chlorine C1, argon Ar. According to the finding of these elements in the main subgroups of III - VI groups, their next electrons among six will be located on the p-sublevel of the 3rd level - The 3rd period is over with an inert element argon, but the 3rd energy level is not finished yet, while there are no electrons on its third possible d-sublevel.
This is followed by 2 elements of the 4th period of the main subgroups of groups I and II: potassium K and calcium Ca. These are s-elements again. Their next electrons will be at the s-sublevel, but already at the 4th level. It is energetically more favorable for these next electrons to start filling the 4th level farther from the nucleus than to fill the 3d sublevel. We write down: The ten following elements of the 4th period from No. 21 scandium Sc to No. 30 zinc Zn are in the side subgroups III - V - VI - VII - VIII - I - II groups. Since they are all d-elements, their next electrons are located on the d-sublevel before the outer level, i.e., the third from the nucleus. We write down:
The next six elements of the 4th period: gallium Ga, germanium Ge, arsenic As, selenium Se, bromine Br, krypton Kr - are in the main subgroups III - VIIJ of groups. Their next 6 electrons are located on the p-sublevel of the external, that is, the 4th level: 3b elements are considered; the fourth period ends with the inert element krypton; the 3rd energy level is also finished. However, at the 4th level, only two sublevels are completely filled: s and p (out of 4 possible).
This is followed by 2 elements of the 5th period of the main subgroups of groups I and II: rubidium No. 37 Rb and strontium No. 38 Sr. These are elements of the s-family, and their next electrons are located on the s-sublevel of the 5th level: The last 2 elements - No. 39 yttrium YU No. 40 zirconium Zr - are already in side subgroups, that is, they belong to the d-family. Two of their next electrons will go to the d-sublevel, before the external one, i.e. Of the 4th level Summing up all the records sequentially, we compose the electronic formula for the zirconium atom no. 40 The derived electronic formula for the zirconium atom can be slightly modified by arranging the sublevels in the order of their level numbering:
The derived formula can, of course, be simplified into the distribution of electrons only over energy levels: Zr - 2 | 8 | 18 | 8 + 2 | 2 (the arrow indicates the place of entry of the next electron; the valence electrons are underlined). The physical meaning of the category of subgroups lies not only in the difference in the place of entry of the next electron into the shell of the atom, but also in the levels at which the valence electrons are located. From a comparison of simplified electronic formulas, for example, chlorine (3rd period, main subgroup of group VII), zirconium (5th period, secondary subgroup of group IV) and uranium (7th period, lanthanoid-actinoid subgroup)
No. 17, C1-2 | 8 | 7
№40, Zr - 2 | 8 | 18 | 8+ 2 | 2
№92, U - 2 | 8 | 18 | 32 | 18 + 3 | 8 + 1 | 2
it can be seen that for elements of any main subgroup, only electrons of the outer level (s and p) can be valence. Elements of side subgroups can have valence electrons of the outer and partially pre-outer levels (s and d). In lanthanides and especially actinides, valence electrons can be at three levels: external, before external and before external. Typically, the total number of valence electrons is equal to the group number.
Properties of elements. Ionization energy. Energy of electron affinity.
Comparative consideration of the properties of elements is carried out in three possible directions of the periodic system: a) horizontal (by period), b) vertical (by subgroup), c) diagonal. To simplify the reasoning, we exclude the 1st period, the unfinished 7th, as well as the entire VIII group. The main parallelogram of the system will remain, in the upper left corner of which there will be lithium Li (no. 3), in the lower left corner - cesium Cs (no. 55). In the upper right - fluorine F (no. 9), in the lower right - astatine At (no. 85).
directions. In the horizontal direction from left to right, the volumes of atoms gradually decrease; occurs, this is as a result of the influence of an increase in the charge of the nucleus on the electron shell. Along the vertical direction from top to bottom, as a result of an increase in the number of levels, the volumes of atoms gradually increase; in the diagonal direction - much less pronounced and shorter - remain close. These are general patterns, of which, as always, there are exceptions.
In the main subgroups, as the volumes of atoms increase, that is, from top to bottom, the elimination of external electrons becomes easier and the attachment of new electrons to the atoms becomes more difficult. The recoil of electrons characterizes the so-called reductive ability of elements, which is especially typical for metals. The attachment of electrons characterizes the oxidizing ability, which is typical for non-metals. Hence, from top to bottom in the main subgroups, the reductive ability of atoms of elements increases; the metallic properties of simple bodies corresponding to these elements also increase. The oxidizing ability decreases.
From left to right in terms of periods, the picture of changes is opposite: the reducing ability of atoms of elements decreases, while the oxidizing ability increases; increase non-metallic properties simple bodies corresponding to these elements.
In the diagonal direction, the properties of the elements remain more or less close. Let's consider this direction with an example: beryllium-aluminum
From beryllium Be to aluminum Al, one can go directly along the diagonal Be → A1, or through boron B, that is, along two legs Be → B and B → A1. The strengthening of non-metallic properties from beryllium to boron and their weakening from boron to aluminum explains why the elements beryllium and aluminum, located diagonally, have some analogy in properties, although they do not belong to the same subgroup of the periodic system. Thus, there is a close relationship between the periodic table, the structure of the atoms of the elements and their chemical properties.
The properties of an atom of any element - to donate an electron and turn into a positively charged ion - are quantified by the expenditure of energy, called the ionization energy I *. It is expressed in kcal / g-atom or xJ / g-atom.
The less this energy, the stronger the atom of the element exhibits reducing properties, the more metallic the element; the more this energy, the weaker the metallic properties, the more the element exhibits non-metallic properties. The property of an atom of any element to accept an electron and transform itself into a negatively charged ion is estimated by the amount of released energy, which is called more energetic than the electron affinity E; it is also expressed in kcal / g-atom or kJ / g-atom.
Electron affinity can be a measure of the ability of an element to exhibit non-metallic properties. The greater this energy, the more non-metallic the element, and, conversely, the lower the energy, the more metallic the element.
Often, to characterize the properties of elements, a value is used, which is called electronegativity.
It: represents the arithmetic sum of the values of the ionization energy and the energy of affinity for an electron
The constant is a measure of the non-metallicity of the elements. The larger it is, the stronger the element exhibits non-metallic properties.
It should be borne in mind that all elements are essentially dual in nature. The division of elements into metals and non-metals is to a certain extent arbitrary, for there are no sharp edges in nature. With the strengthening of the element's metallic properties, its nonmetaglic properties are weakened and vice versa. The most "metallic" of the elements - francium Fr - can be considered the least non-metallic, the most "non-metallic" - fluorine F - can be considered the least metallic.
Summing up the values of the calculated energies - the ionization energy and the electron affinity energy - we get: for cesium the value is 90 kcal / g-a., For lithium 128 kcal / g-a., For fluorine = 510 kcal / g-a. (the value is also expressed in kJ / g-a.). These are the absolute values of electronegativity. For simplicity, the relative values of electronegativity are used, taking the electronegativity of lithium (128) as unity. Then for fluorine (F) we get:
For cesium (Cs), the relative electronegativity will be
On the graph of changes in the electronegativity of elements of the main subgroups
I-VII groups. the electronegativities of the elements of the main subgroups of groups I-VII are compared. The given data indicate the true position of hydrogen in the 1st period; unequal increase in metallicity of elements, from top to bottom in various subgroups; some similarity of elements: hydrogen - phosphorus - tellurium (= 2.1), beryllium and aluminum (= 1.5) and a number of other elements. As can be seen from the above comparisons, using the values of electronegativity, one can approximately compare with each other, even elements of different subgroups, and different periods.
Graph of changes in electro-negative elements of the main subgroups of I-VII groups.
The periodic law and the periodic table of elements are of great philosophical, scientific and methodological significance. They are: a means of knowing the world around us. The periodic law reveals and reflects the dialectical-materialistic essence of nature. Periodic, law and periodic system of elements with all convincing evidence of the unity and materiality of the world around us. They are the best confirmation of the validity of the main features of the Marxist dialectical method of cognition: a) the relationship and interdependence of objects and phenomena, b) the continuity of movement and development, c) transition quantitative changes in qualitative, d) struggle and unity of opposites.
Electronegativity values
|
Group |
I A |
II A |
III B |
IV B |
V B |
VI B |
VII B |
VIII B |
VIII B |
VIII B |
I B |
II B |
III A |
IV A |
V A |
VI A |
VII A |
VIII A | |
|
Period | |||||||||||||||||||
|
1 |
H 2.0 |
He 4.5 | |||||||||||||||||
|
2 |
Li 0.98 |
Be 1.57 |
B 2.04 |
C 2.55 |
N 3.2 |
O 3.44 |
F 3.98 |
Ne 4.4 | |||||||||||
|
3 |
Na 0.93 |
Mg 1.31 |
Al 1.61 |
Si 1.6 |
P 2.49 |
S 2.58 |
Cl 3.0 |
Ar 4.3 | |||||||||||
|
4 |
K 0.82 |
Ca 1.00 |
Sc 1.36 |
Ti 1.54 |
V 1.63 |
Cr 1.66 |
Mn 1.55 |
Fe 1.83 |
Co 1.88 |
Ni 1.91 |
Cu 1.90 |
Zn 1.65 |
Ga 1.81 |
Ge 2.01 |
As 2.4 |
Se 2.55 |
Br 2.96 |
Kr 3.00 | |
|
5 |
Rb 0.82 |
Sr 0.95 |
Y 1.22 |
Zr 1.33 |
Nb 1.6 |
Mo 2.16 |
Tc 1.9 |
Ru 2.2 |
Rh 2.28 |
Pd 2.20 |
Ag 1.93 |
Cd 1.69 |
In 1.78 |
Sn 1.96 |
Sb 2.21 |
Te 2.4 |
I 2.66 |
Xe 2.60 | |
|
6 |
Cs 0.79 |
Ba 0.89 |
* |
Hf 1.3 |
Ta 1.5 |
W 2.36 |
Re 1.9 |
Os 2.2 |
Ir 2.20 |
Pt 2.28 |
Au 2.64 |
Hg 2.2 |
Tl 1.62 |
Pb 2.33 |
Bi 2.02 |
Po 2.3 |
At 2.5 |
Rn 2.2 | |
|
7 |
Fr 0.8 |
Ra 0.9 |
** |
Rf |
Db |
Sg |
Bh |
Hs |
Mt |
Ds |
Rg |
Cn |
Uut |
Uuq |
Uup |
Uuh |
Uus |
Uuo | |
|
Lanthanides |
* |
La 1.1 |
Ce 1.12 |
Pr 1.13 |
Nd 1.14 |
Pm 1.13 |
Sm 1.17 |
Eu 1.2 |
Gd 1.2 |
Tb 1.1 |
Dy 1.22 |
Ho 1.23 |
Er 1.24 |
Tm 1.25 |
Yb 1.1 |
Lu 1.27 | |||
|
Actinides |
** |
Ac 1.1 |
Th 1.3 |
Pa 1.5 |
U 1.38 |
Np 1.36 |
Pu 1.28 |
Am 1.13 |
Cm 1.28 |
Bk 1.3 |
Cf 1.3 |
Es 1.3 |
Fm 1.3 |
Md 1.3 |
# 1.3 |
Lr 1.291 | |||
The main subgroup of group V of the periodic system includes nitrogen, phosphorus, arsenic, antimony and bismuth.
These elements, having five electrons in the outer layer of the atom, are generally characterized as non-metals. However, the ability to attach electrons is much less pronounced in them than in the corresponding elements of groups VI and VII. Due to the presence of five outer electrons, the highest positive oxidation of the elements of this subgroup is -5, and the negative one is 3. Due to the relatively lower electronegativity, the bond of the considered elements with hydrogen is less polar than the bond with hydrogen of the elements of VI and VII groups. That's why hydrogen compounds These elements do not remove hydrogen ions H in an aqueous solution, thus, do not have acidic properties.
The physical and chemical properties of the elements of the nitrogen subgroup change with an increase in the serial number in the same sequence that was observed in the previously considered groups, But since the non-metallic properties are expressed weaker than that of oxygen and even more so for fluorine, then the weakening of these properties when going to the next elements entails the appearance and growth of metallic properties. The latter are already noticeable in arsenic, antimony has approximately the same properties, and in bismuth, metallic properties prevail over non-metallic ones.
DESCRIPTION OF ELEMENTS.
NITROGEN(from the Greek ázōos - lifeless, Latin Nitrogenium), N, a chemical element of the V group of the periodic system of Mendeleev, atomic number 7, atomic mass 14.0067; colorless gas, odorless and tasteless.
Historical reference. Nitrogen compounds - nitrate, Nitric acid, ammonia - were known long before nitrogen was obtained in a free state. In 1772, D. Rutherford, burning phosphorus and other substances in a glass bell, showed that the gas remaining after combustion, which he called "suffocating air," does not support breathing and combustion. In 1787 A. Lavoisier established that the "vital" and "suffocating" gases that make up the air are simple substances, and suggested the name "nitrogen". In 1784 G. Cavendish showed that nitrogen is a part of saltpeter; from here comes the Latin name nitrogen (from the late Latin nitrum - saltpeter and the Greek gennao - I give birth, I produce), proposed in 1790 by J. A. Chaptal. By the beginning of the 19th century. the chemical inertness of nitrogen in a free state was found out and exceptional role it in compounds with other elements as bound nitrogen. Since then, the "binding" of nitrogen in the air has become one of the most important technical problems in chemistry.
Prevalence in nature. Nitrogen is one of the most abundant elements on Earth, and its main mass (about 4 × 1015 tons) is concentrated in a free state in the atmosphere. In the air, free nitrogen (in the form of N2 molecules) is 78.09% by volume (or 75.6% by weight), not counting its minor impurities in the form of ammonia and oxides. The average nitrogen content in the lithosphere is 1.9 × 10-3% by weight.
Natural nitrogen compounds. - ammonium chloride NH4Cl and various nitrates (see. Saltpeter.) Large accumulations of nitrate are characteristic of the dry desert climate (Chile, Central Asia). For a long time, nitrate was the main supplier of nitrogen for industry (now the industrial synthesis of ammonia from nitrogen in air and hydrogen is of primary importance for binding nitrogen). Small quantities bound nitrogen is found in coal (1-2.5%) and oil (0.02-1.5%), as well as in the waters of rivers, seas and oceans. Nitrogen accumulates in soils (0.1%) and in living organisms (0.3%).
Although the name "nitrogen" means "not life-sustaining," it is actually an essential element for life. The protein of animals and humans contains 16 - 17% nitrogen. In the organisms of carnivores, protein is formed due to the consumed protein substances present in the organisms of herbivores and in plants. Plants synthesize protein by assimilating nitrogenous substances contained in the soil, mainly inorganic ones. Significant amounts of nitrogen enter the soil due to nitrogen-fixing microorganisms capable of converting free nitrogen in the air into nitrogen compounds.
The nitrogen cycle is carried out in nature, the main role in which microorganisms play - nitrifying, denitrifying, nitrogen-fixing, etc. However, as a result of extraction from the soil by plants huge amount bound nitrogen (especially with intensive farming) soils are depleted in nitrogen. Nitrogen deficiency is typical for agriculture in almost all countries, nitrogen deficiency is also observed in animal husbandry ("protein starvation"). Plants do not develop well on soils that are poor in available nitrogen. Nitrogen fertilizers and protein feeding of animals are the most important means of lifting Agriculture. Economic activity a person disrupts the nitrogen cycle. Thus, the combustion of fuel enriches the atmosphere with nitrogen, and the factories producing fertilizers bind the nitrogen of the air. The transport of fertilizers and agricultural products redistributes nitrogen to the surface of the earth.
Nitrogen is the fourth most common element Solar system(after hydrogen, helium and oxygen).
Isotopes, atom, molecule. Natural nitrogen consists of two stable isotopes: 14N (99.635%) and 15N (0.365%). The 15N isotope is used in chemical and biochemical research as a tagged atom. Of artificial radioactive nitrogen isotopes, 13N has the longest half-life (T1 / 2 - 10.08 min), the rest are very short-lived. In the upper layers of the atmosphere, under the action of cosmic neutrons, 14N is converted into the radioactive isotope of carbon 14C. This process is also used in nuclear reactions to produce 14C. The outer electron shell of the nitrogen atom. consists of 5 electrons (one lone pair and three unpaired - configuration 2s22p3). Most often nitrogen. in compounds 3-covalent due to unpaired electrons (as in ammonia NH3). The presence of a lone pair of electrons can lead to the formation of another covalent bond, and nitrogen becomes 4-covalent (as in the ammonium ion NH4 +). The oxidation states of nitrogen vary from +5 (in N205) to -3 (in NH3). Under normal conditions, in a free state, nitrogen forms an N2 molecule, where the N atoms are linked by three covalent bonds. The nitrogen molecule is very stable: the energy of its dissociation into atoms is 942.9 kJ / mol (225.2 kcal / mol), therefore, even at t about 3300 ° C, the degree of dissociation of nitrogen is. is only about 0.1%.
Physical and chemical properties. Nitrogen is slightly lighter than air; density 1.2506 kg / m3 (at 0 ° C and 101,325 n / m2 or 760 mm Hg), tp -209.86 ° C, tboil -195.8 ° C. A. liquefies with difficulty: its critical temperature is rather low (-147.1 ° C), and its critical pressure is high 3.39 MN / m2 (34.6 kgf / cm2); density of liquid nitrogen 808 kg (m3. In water, nitrogen is less soluble than oxygen: at 0 ° C in 1 m3 H2O dissolves 23.3 g of nitrogen. Better than in water, nitrogen is soluble in some hydrocarbons.
Only with such active metals as lithium, calcium, magnesium, nitrogen interacts when heated to relatively low temperatures. Nitrogen reacts with most other elements at high temperatures and in the presence of catalysts. Compounds of nitrogen with oxygen N2O, NO, N2O3, NO2 and N2O5 have been well studied. From them, during the direct interaction of elements (4000 ° C), NO oxide is formed, which, upon cooling, is easily oxidized further to NO2 dioxide. In air, nitrogen oxides are formed during atmospheric discharges. They can also be obtained by the action of ionizing radiation on a mixture of nitrogen and oxygen. When nitrous N2O3 and nitric N2O5 anhydrides are dissolved in water, nitrous acid HNO2 and nitric acid HNO3, respectively, are obtained, which form salts - nitrites and nitrates. Nitrogen combines with hydrogen only at high temperatures and in the presence of catalysts, thus forming ammonia NH3. In addition to ammonia, numerous other nitrogen compounds with hydrogen are known, for example, hydrazine H2N-NH2, diimide HN-NH, hydrazoic acid HN3 (H-N-NºN), octazone N8H14, etc .; most nitrogen compounds with hydrogen are isolated only in the form of organic derivatives. Nitrogen does not directly interact with halogens, therefore, all nitrogen halides are obtained only indirectly, for example, nitrogen fluoride NF3- by the interaction of fluorine with ammonia. As a rule, nitrogen halides are low-stable compounds (with the exception of NF3); more stable nitrogen oxyhalides - NOF, NOCI, NOBr, N02F and NO2CI. There is also no direct nitrogen combination with sulfur; nitrogenous sulfur N4S4 is produced by the reaction of liquid sulfur with ammonia. When hot coke interacts with nitrogen, cyanogen (CN) is formed.;. By heating nitrogen with acetylene C2H2 to 1500 ° C, hydrogen cyanide HCN can be obtained. The interaction of nitrogen with metals at high temperatures leads to the formation of nitrides (eg Mg3N2).
16. Which of the gases taken with the same mass takes up the largest volume under the same conditions:
17. Determine the molar mass equivalent (g / mol) of sulfur in sulfur oxide (VI):
18. What is mass fraction(%) of metal in oxide if the molar mass of the trivalent metal equivalent is 15 g / mol:
19. What is the relative molecular weight of a gas if this gas is 2.2 times heavier than air:
20. Which of the above equations is called the Mendeleev-Clapeyron equation:
3) PV = RT
21. Indicate 3 gases that have the same density for any other gas:
1) CH 4, SO 2, Cl 2
2) C 2 H 4, CH 4, F 2
3) CO, Cl 2, H 2
4) CO, C 2 H 4, N 2
5) N 2, CH 4, H 2
22. How many moles of oxygen are formed from 3 mol of potassium chlorate during its complete thermal decomposition:
23. What amount (mol) of FeS 2 is required to obtain 64 g SO 2 according to the equation:
4 FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2;
24. What mass (g) of calcium carbonate will be consumed to obtain 44.8 liters of carbon dioxide, measured at normal conditions:
1) 200,0;
25. The equivalent of aluminum is:
1) an aluminum atom;
2) 1/2 part of an aluminum atom;
3) 1/3 part is an aluminum atom;
4) two aluminum atoms;
5) 1 mole of aluminum atoms.
26.The law of constancy of the composition of substances is valid for substances:
1) with a molecular structure;
2) with a non-molecular structure;
3) with an ionic crystal lattice;
4) with an atomic crystal lattice;
5) for oxides and salts.
27. The equivalent of magnesium is:
1) a magnesium atom;
2) 1/2 part of a magnesium atom;
3) 1/3 of the magnesium atom;
4) two magnesium atoms;
5) 1 mole of magnesium atoms.
28. The neutralization of 2.45 g of acid consumes 2.80 g of potassium hydroxide. Define
molar mass of acid equivalent:
1) 98 g / mol;
2) 36.5 g / mol;
3) 63 g / mol;
4) 40 g / mol;
G / mol.
Classification and nomenclature of inorganic compounds
1) Na 2 O; CaO; CO 2
2) SO 3; CuO; CrO 3
3) Mn 2 O 7; CuO; CrO 3
4) SO 3; CO 2; P 2 O 5
5) Na 2 O; H 2 O; CO 2
30. Only acid oxides series:
1) CO 2; SiO 2; MnO; CrO 3
2) V 2 O 5; CrO 3; TeO 3; Mn 2 O 7
3) CuO; SO 2; NiO; MnO
4) CaO; P 2 O 3; Mn 2 O 7; Cr 2 O 3
5) Na 2 O; H 2 O; CuO; Mn 2 O 7
31. Cannot be used to neutralize sulfuric acid:
1) sodium bicarbonate;
2) magnesium oxide;
3) hydroxomagnesium chloride;
4) sodium hydrogen sulfate;
5) sodium oxide
32. To neutralize sulfuric acid, you can use:
2) Mg (OH) 2
33. Using a glass tube, exhale carbon dioxide into solutions. The change will be in solution:
3) Ca (OH) 2;
34. By dissolving the corresponding oxide in water, you can get:
35. Under certain conditions, salt is formed in the case of:
1) N 2 O 5 + SO 3;
4) H 2 SO 4 + NH 3;
36. Can form acidic salts:
1) H 3 PO 4;
37. May form basic salts:
2) Ba (OH) 2;
38. The mass of limestone required to obtain 112 kg of quicklime:
39. Reacts with water:
2) CаO;
40. Let's dissolve in water:
3) Ba (OH) 2;
41. To obtain potassium phosphate, potassium hydrogen phosphate must be influenced by:
42. Acidic oxide:
3) Mn 2 O 7;
43. Will directly interact in an aqueous solution:
2) Cu (OH) 2 and ZnO;
3) AI 2 O 3 and HCI;
4) Rb 2 O and NaOH;
5) CaO and K 2 O.
44. All acidic salts in the group:
1) KCI, CuOHCI, NaHSO 4;
2) KAI (SO 4) 2, Na, Ca (HCO 3) 2;
3) CuS, NaHSO 3, Cu (HS) 2;
4) NaHCO 3, Na 2 HPO 4, NaH 2 PO 4;
5) AIOHCI 2, NaHCO 3, NaCN.
45. Does not form acidic salts:
4) HPO 3;
46. Wrong title:
1) ferrous sulfate;
2) potassium sulfate;
3) iron (II) hydrochloride;
4) copper (I) chloride;
5) ammonium sulfate.
47. When water is split off from a monobasic acid with a mass of 16.0 g, formed by an element in an oxidation state of +5, an oxide with a mass of 14.56 g will be obtained. Acid was taken:
1) nitrogen;
2) metavanadium;
3) orthophosphoric;
4) arsenic;
5) chloric.
48. When calcining metal (III) weighing 10.8 g in air, a metal oxide weighing 20.4 g was obtained. For calcining, the following was taken:
2) aluminum AI;
3) iron Fe;
4) scandium Sc;
5) sodium Na.
49. A sign characterizing hydrochloric acid:
1) two-base;
2) weak;
3) volatile;
4) oxygenated;
5) acid is an oxidizing agent.
50. Dibasic acid:
1) nitrogen;
2) salt;
3) acetic;
4) blue;
Selenium
51. Monobasic acid:
1) selenium;
2) phosphorous;
3) tellurium;
4) boric;
5) bluish.
52. Forms two types of acidic salts:
1) sulfuric acid;
2) orthophosphoric acid;
3) metaphosphoric acid;
4) selenic acid;
5) sulfurous acid.
53. Does not form acidic salts:
1) sulfuric acid;
2) phosphoric acid;
3) metaphosphoric acid;
4) selenic acid;
5) sulfurous acid.
54. Indicate the cationic complex:
1) Na 3;
3) K 3;
4) CI 3;
5) K 2.
55. Complex non-electrolyte:
1) Na 3;
2) ;
3) K 3;
4) CI 3;
5) K 2.
56. Anionic complex:
1) potassium hexacyanoferrate (III);
2) tetrachlorodiammineplatinum (IV);
3) diamminesilver chloride;
57. Complex non-electrolyte:
1) potassium hexacyanoferrate (III);
2) tetrachlorodiammineplatinum (IV);
3) diamminesilver chloride;
4) tetraamminecopper (II) sulfate;
5) chloride of hexaaquachrome (III).
58. Formula of hexaaquachrome (III) chloride:
1) Na 3;
2) CI
3) CI 2;
4) CI 3;
5) K 2 Cr 2 O 7.
59. Formula of hexaaquachrome (II) chloride:
1) Na 3;
2) CI
3) CI 2; 3bl
4) CI 3;
5) K 2 Cr 2 O 7.
60. Yellow blood salt refers to:
1) To aqua complexes;
2) Hydrates;
3) To acidocomplexes;
4) Ammonia;
5) To chelates.
61. Copper sulfate refers to:
1) To aqua complexes;
2) Hydrates;
3) To acidocomplexes;
4) Ammonia;
5) To chelates.
62. To obtain CaCO 3, add to the Ca (HCO 3) 2 solution:
1) Ca (OH) 2;
“The structure of matter and periodic law DI. Mendeleev "
63. In the nucleus of the most common lead isotope 207 Pb neutrons:
2) 125
64. The maximum number of electrons at the n = 3 level:
65. At the energy level with n = 4 sublevels:
66. The number of energy levels in a tungsten atom:
67. In the nucleus of the osmium atom of protons:
68. The nucleus of a krypton atom contains:
P and 44n
69. The number of electrons in the chromium ion:
70. Ion, which has 18 electrons and 16 protons, has a nuclear charge:
71. The maximum number of electrons that can occupy a 3s-orbital:
72. The electronic configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 has an atom:
73. The orbital designations are incorrect:
3) 1p, 2d
74. Equal to the argon atom electronic configuration has a particle:
1) Ca 2+
75. The electron affinity is called:
1) the energy required to detach an electron from an unexcited atom;
2) the ability of an atom of a given element to pull off the electron density;
3) the transition of an electron to a higher energy level;
4) the release of energy when an electron is attached to an atom or ion;
5) chemical bond energy.
76. As a result of a nuclear reaction 
an isotope is formed:
77. In a hydrogen atom absorption of a photon with a minimum energy requires the transition of an electron:
78. The corpuscular-wave nature of the electron is characterized by the equation:
79. For a valence electron of a potassium atom, the values of quantum numbers (n, l, m l, m s):
1) 4, 1, -1, - :
2) 4, 1, +1, +: 3bm
3) 4, 0, 0, + :
4) 5, 0, +1, + :
80. The charge of the atomic nucleus, in which the configuration of valence electrons in the ground state is… 4d 2 5s 2:
81. The principal quantum number n determines:
1) shape electronic cloud;
2) the energy of the electron;
82. Orbital quantum number l determines:
1) the shape of the electron cloud;
2) the energy of the electron;
3) the orientation of the electron cloud in space;
4) the rotation of the electron around its own axis;
5) hybridization of the electronic cloud.
83. Magnetic quantum number m determines:
1) the shape of the electron cloud;
2) the energy of the electron;
3) the orientation of the electron cloud in space;
4) the rotation of the electron around its own axis;
5) hybridization of the electronic cloud.
84. The spin quantum number m s determines:
1) the shape of the electron cloud;
2) the energy of the electron;
3) the orientation of the electron cloud in space;
4) the rotation of the electron around its own axis;
5) hybridization of the electronic cloud.
85. When - decay, the nucleus of an atom of a radioactive element emits:
1) an electron;
2) positron;
4) two protons;
5) two neutrons.
86. During - - decay, the nucleus of an atom of a radioactive element emits:
1) an electron;
2) positron;
3) two protons and two neutrons, combined into the nucleus of a helium atom;
4) two protons;
5) two neutrons.
87. At + - decay, the nucleus of an atom of a radioactive element emits:
1) an electron;
2) positron;
3) two protons and two neutrons, combined into the nucleus of a helium atom;
4) two protons;
5) two neutrons.
88. Smallest value sum (n + l) has atomic orbital:
89. The highest value of the sum (n + l) has the atomic orbital
90. The nitrogen atom will be more stable if on the 2p-sublevel three electrons are distributed, one in each orbital. This matches the content:
2) Pauli's principle;
3) Gund's Rules;
4) the 1st rule of Klechkovsky;
5) the 2nd rule of Klechkovsky.
91. The twenty-first electron of the scandium atom is located on the 3d-sublevel, and not on the 4p-sublevel. This matches the content:
1) The principle of least energy;
2) Pauli's principle;
3) Gund's Rules;
4) the 1st rule of Klechkovsky;
5) the 2nd rule of Klechkovsky.
92. The nineteenth electron of the potassium atom is located at the 4s-sublevel, and not at the 3d-sublevel. This matches the content:
1) The principle of least energy;
2) Pauli's principle;
3) Gund's Rules;
4) the 1st rule of Klechkovsky;
5) the 2nd rule of Klechkovsky.
93. The only electron of the hydrogen atom in the ground state is located at the first energy level. This matches the content:
1) The principle of least energy;
2) Pauli's principle;
3) Gund's Rules;
4) the 1st rule of Klechkovsky;
5) the 2nd rule of Klechkovsky.
94. The maximum number of electrons at the second energy level of atoms of elements
equals 8. This corresponds to the content:
1) The principle of least energy;
2) Pauli's principle;
3) Gund's Rules;
4) the 1st rule of Klechkovsky;
5) the 2nd rule of Klechkovsky.
95. One of the mechanisms of covalent bond formation:
1) radical;
2) exchange;
3) molecular;
4) ionic;
5) chain.
96. An example of a non-polar molecule having a polar covalent bond would be:
4) CCl 4
97. Non-polar molecule:
98. In the series of molecules LiF - BeF 2 - BF 3 - CF 4 - NF 3 - OF 2 - F 2:
1) the nature of the connection does not change;
2) the ionic nature of the bond is enhanced;
3) the covalent nature of the bond weakens;
4) the covalent nature of the bond is enhanced;
5) correct option No answer.
99. Covalent bond by the donor-acceptor mechanism is formed in the molecule:
2) CCl 4;
3) NH 4 C1;
4) NH 3;
100. The nitrogen molecule produces:
1) only -connections;
2) only -connections;
3) both - and - communication;
4) single bond;
5) double bond.
101. Methane molecule has the structure:
1) flat;
2) tetrahedral;
3) pyramidal;
4) square;
102. The formation of an ionic lattice is typical for:
1) cesium iodide;
2) graphite;
3) naphthalene;
4) diamond;
103. Which of the following substances is characterized by the formation of an atomic lattice:
1) ammonium nitrate;
2) diamond;
4) sodium chloride;
5) sodium.
104. Chemical elements are arranged in ascending order of electronegativity in
1) Si, P, Se, Br, Cl, O;
2) Si, P, Br, Se, Cl, O;
3) P, Si, Br, Se, Cl, O;
4) Br, P, Cl, Si, Se;
5) Si, P, Se, Cl, O, Br
105. The valence orbitals of the beryllium atom in the beryllium hydride molecule ... are hybridized
106. Beryllium hydride molecule has the structure:
1) square
Flat
3) tetrahedral
5) spherical.
107. The valence orbitals of the boron atom in the BF 3 molecule are hybridized according to the type:
108. Which of the molecules is the most durable?
109. Which of the indicated molecules has the greatest dipole?
110. What is the spatial configuration of the molecule during sp 2 hybridization of AO:
1) linear
2) tetrahedron
3) flat square
Flat trigonal
111. The molecule has an octahedral structure if the next hybridization occurs
3) d 2 sp 3
112. The modern theory of the structure of the atom is based on the concepts:
1) classical mechanics;
3) Bohr's theories;
4) electrodynamics;
5) chemical kinetics.
113. Of the following characteristics of the atoms of the elements change periodically:
1) the charge of the nucleus of an atom
2) relative atomic mass;
3) the number of energy levels in the atom;
4) the number of electrons at the external energy level;
5) the total number of electrons.
114. Within a period, an increase in the ordinal number of an element is usually accompanied by:
1) a decrease in the atomic radius and an increase in the electronegativity of the atom;
2) an increase in the atomic radius and a decrease in the electronegativity of the atom;
3) a decrease in the atomic radius and a decrease in the electronegativity of the atom
4) an increase in the atomic radius and an increase in the electronegativity of an atom
5) decrease in electronegativity.
115. The atom of which of the elements most easily gives up one electron:
1) sodium, atomic number 11;
2) magnesium, serial number 12;
3) aluminum, serial number 13;
4) silicon, serial number 14;
5) sulfur, serial number 16.
116. Atoms of elements of the IA group of the periodic table of elements have the same number:
1) electrons at the external electronic level;
2) neutrons;
3) all electrons;
4) electronic shells;
5) protons.
117. Which of the following elements is named after the country:
118. Which series includes only transitional elements:
1) items 11, 14, 22, 42;
2) elements 13, 33, 54, 83;
3) items 24, 39, 74, 80;
4) elements 19, 32, 51, 101;
5) elements 19, 20, 21, 22.
119. The atom of which of the elements of the VA group has the maximum radius:
2) phosphorus;
3) arsenic;
4) bismuth;
5) antimony.
120. What series of elements are presented in ascending order of atomic radius:
1) O, S, Se, Te;
3) Na, Mg, AI, Si;
4) J, Br, CI, F;
5) Sc, Te, V, Cr.
121. Metallic character of the properties of elements in the series Mg - Ca - Sr - Ba
1) decreases;
2) increases;
3) does not change;
4) decreases and then increases;
5) increases and then decreases.
122. The main properties of hydroxides of elements of the JА group as the serial number increases
1) decrease,
2) increase,
3) remain unchanged,
4) decrease and then increase,
5) increase and then decrease.
123. Simple substances of which elements have the greatest similarity between physical and chemical properties:
3) F, CI;
124. The existence of which of the above elements was predicted by D.I. Mendeleev:
3) Sc, Ga, Ge;
125. What distinguishes large periods from small ones:
1) the presence of alkali metals;
2) absence of inert gases;
3) the presence of d- and f-elements;
4) the presence of non-metals;
5) the presence of elements with metallic properties.
126. How, using the electronic formula of an element, determine the period in which this element is located:
1) by the value of the principal quantum number of the external energy level;
2) by the number of valence electrons;
3) by the number of electrons in the external energy level;
4) by the number of sublevels in the external energy level;
5) by the value of the sublevel where the last valence electron is located.
127. Which element has the lowest ionization potential:
128. The chemical element of the third period forms the highest oxide of the composition E 2 O 3. How are electrons distributed in an atom of a given element?
1) 1s 2 2s 2 2p 1
2) 1s 2 2s 2 2p 6 3s 1
3) 1s 2 2s 2 2p 6 3s 2 3p 1
4) 1s 2 2s 2 2p 6 3s 2 3p 6
5) 1s 2 2s 2 2p 3
129. What chemical element forms the base with the most pronounced properties
1) calcium
3) aluminum
Potassium
5) beryllium
130. A chemical element has the following distribution of electrons over the electron layers in the atom 2.8.6. What position does it occupy in the periodic table chemical elements DI. Mendeleev:
1) 6 period 6 group
Period 6 group
3) 2 period 6 group
4) 3 period 2 group
5) 2 period 8 group
131. The quantum numbers of the last electron in the atom of the element are equal to n = 5, l = 1, m = -1, m s = -. Where is this element in the periodic table?
1) 5 period, first group
2) 5 period, main subgroup of 4 groups
3) 4 period, sixth group
period, sixth group main subgroup
5) 5 period, the sixth group is a side subgroup.
132. Formula of the highest oxide of the chemical element EO 2. To which group of the main subgroup of the periodic table of chemical elements D.I. Mendeleev does this element belong?
Fourth
5) sixth.
133. From the above list of elements - Li, Na, Ag, Au, Ca, Ba - alkali metals include:
1) all metals;
2) Li, Na;
3) Li, Na, Ag, Au;
134. In the order from Li to Fr:
1) enhanced metallic properties;
2) metallic properties decrease;
3) decreases atomic radius;
4) the bond of valence electrons with the nucleus is enhanced;
5) decreases activity in relation to water
135. Metals do not include the sequence of elements:
3) B, As, Te;
136. With an increase in the serial number of the element, the acidic properties of oxides N 2 O 3 - P 2 O 3 - As 2 O 3
Sb 2 O 3 - Bi 2 O 3
1) intensify;
2) weaken;
3) remain unchanged;
4) intensify, then weaken;
5) weaken, then intensify.
137. An ammonia molecule has the form:
1) curved;
2) linear;
3) planar;
4) pyramidal;
138. In the series C-Si-Ge-Sn-Pb, non-metallic signs of elements:
1) increase;
2) weaken;
3) do not change;
4) increase and then weaken;
5) weaken and then increase.
139. The valence orbitals at the carbon atom in the CH 4 methane molecule can be described on the basis of
ideas about hybridization of the type (sp; sp 2; sp 3; d 2 sp 3; dsp 2).
In this case, the methane molecule has the form:
1) linear;
2) flat;
3) tetrahedral;
5) square.
140. The valence orbitals at the silicon atom in the silane molecule SiH 4 can be described on the basis of the concept of hybridization of the type (sp; sp 2; sp 3; d 2 sp 3; dsp 2).
Therefore, the silane molecule has the form:
1) linear;
2) flat;
3) tetrahedral;
5) square.
141. What is the maximum number of covalent bonds a nitrogen atom can form:
142. The nitrogen atom of the ammonia molecule with the hydrogen ion forms:
1) ionic bond;
2) covalent bond by the exchange mechanism;
3) non-polar covalent bond;
4) covalent bond by donor - acceptor mechanism;
5) hydrogen bond.
143. Which statement is false:
4) Ionic bond possesses saturability;
144. Which statement is false:
1) The covalent bond is saturable;
2) The covalent bond has directionality;
3) The ionic bond is unsaturated;
4) Ionic bond has directionality;
5) The ionic bond is non-directional.
"Regularities of chemical processes and their energetics"
145. What changes in temperature T and pressure P contributes to the formation of CO by the reaction C (tv.) + CO 2 (g) 2CO (g) -119.8 kJ:
1) an increase in T and an increase in P;
2) an increase in T and a decrease in P;
3) a decrease in T and an increase in P;
4) lowering T and lowering P;
5) increasing P.
146. How many times will the rate of a chemical reaction increase with an increase in temperature by 30 0, if the temperature coefficient of the rate is 2?
147. How many degrees should the temperature be lowered in order for the reaction rate to decrease by a factor of 27, if the temperature coefficient of the rate is 3?
148. How many times will the reaction rate X + 2Y = Z increase with increasing concentration?
Y 3 times?
149. How many times will the rate of the forward reaction be greater in comparison with the rate of the reverse reaction in the 2NO + O 2 2NO 2 system with a 2-fold increase in pressure?
150. What is the correct rate expression for the system: 2Cr + 3Cl 2 = 2CrCl 3
154. The catalyst accelerates the chemical reaction due to:
1) a decrease in the activation energy;
2) increasing the activation energy;
3) a decrease in the heat of reaction;
4) increasing concentration;
5) all answers are not correct.
155. The equilibrium of the reaction Fe 3 O 4 + 4CO «3Fe + 4CO 2 -43.7 kJ is shifted to the left:
1) when the temperature drops;
2) when the temperature rises;
3) with increasing pressure;
4) with an increase in the concentration of starting substances;
5) when adding a catalyst.
156. How many times will the rate of a chemical reaction increase with an increase in temperature by 30 0, if the temperature coefficient of the rate is 3?
157. By how many degrees should the temperature be increased in order for the reaction rate to increase by a factor of 27, if the temperature coefficient of the rate is 3?
158. How many times does the reaction rate X + 2Y = Z increase with an increase in the concentration of X in 3 times?
159. How many times will the rate of the forward reaction be greater in comparison with the rate of the reverse reaction in the 2CO + O 2 2CO 2 system with a 2-fold increase in pressure?
160. How will the rate of the gas reaction 2NO 2 = N 2 O 4 increase with an increase in the concentration of NO 2 by 5 times?
161. How many times will the rate of the gas reaction 2NO + O 2 = 2NO 2 decrease when the mixture of reacting gases is diluted 3 times?
162. How many degrees should the temperature be lowered in order for the reaction rate to decrease 81 times at a temperature coefficient of 3?
163. How many times will the reaction rate 2NO + O 2 = 2NO 2 increase when the pressure in the system increases by 4 times?
164. How many times will the rate of the forward reaction be higher in comparison with the rate of the reverse reaction in the 2NO + O 2 2NO 2 system when the pressure in the system increases by 5 times?
165. How will the reaction rate of 2SO 2, g + O 2, g 2SO 3, g change with increasing concentration 1) will increase by 3 times; 2) will increase 9 times; 3) will decrease by 3 times; 4) will decrease by 9 times; 5) will not change. 166. How will the reaction rate of 2O 3, g 3O 2, g change with a 2-fold increase in pressure? 1) will decrease by 2 times; 2) decrease by 8 times; 3) will increase by 4 times; 4) will decrease by 4 times; 5) will double. 167. How will the reaction rate change 2NO g + O 2, g 2NO 2, g while decreasing concentration of NO and O 2 by 2 times? 1) will increase 2 times; 2) will decrease by 2 times; 3) will increase 2 to 4 times; 4) will decrease 2 to 4 times; It will decrease by 8 times. 168. How will the rate of the direct reaction H 2 O, g H 2, g + O 2, g change if the pressure in the system increases 4 times? 1) will increase by 2 times; 2) will decrease by 2 times; 3) will not change; 4) will increase by 4 times; 5) will decrease by 4 times. 169. The law of action of the masses was discovered: 1) M.V. Lomonosov 2) G.I. Hess 3) J.W. Gibbs K. Guldberg and P. Waage 5) Van't - Goffom 170. Which of these systems is homogeneous Sodium chloride solution 2) ice water 3) saturated solution with sediment 4) coal and sulfur in the air 5) a mixture of gasoline and water 171. The value of the rate constant of a chemical reaction does not depend 1) from the nature of the reactants 2) from temperature 3) from the presence of catalysts From the concentration of substances 5) from any factors 172. The activation energy is 1) the energy required to detach an electron from an atom 2) the excess energy of which the molecules must have per 1 mole for ohm, so that their collision could lead to the formation of a new substance 3) ionization potential 4) the energy that is released as a result of the reaction 5) the energy that is released when an electron is attached to an atom. 173. An increase in the reaction rate with increasing temperature is usually characterized by: 1) the rate constant of the chemical reaction 2) constant of chemical equilibrium The elements of the main subgroup of the sixth group of the periodic system are oxygen, sulfur, selenium, tellurium and polonium. The last of these is a radioactive metal; both natural and artificially obtained isotopes are known. In the outer electron shell, the atoms of the elements under consideration contain six electrons - two in the -orbital and four in the p-orbital. The oxygen atom differs from the atoms of other elements of the subgroup by the absence of a sublevel in the outer electron layer: As indicated in Sec. 41, such an electronic structure of the oxygen atom leads to large energy expenditures for the "steaming" of its electrons, which are not compensated for by the energy of formation of new covalent bonds. Therefore, the covalence of oxygen is usually two. However, in some cases, an oxygen atom with lone electron pairs can act as an electron donor and form additional covalent bonds in a donor-acceptor manner. For sulfur and for other elements of the subgroup, the number of unpaired electrons in the atom can be increased by transferring the s- and p-electrons to the -sublevel of the outer layer. In this regard, these elements exhibit covalence equal not only to 2, but also to 4 and 6. All elements of this subgroup, except for polonium, are non-metals, although less active than halogens. In their compounds, they show both negative and positive oxidation. In compounds with metals and with hydrogen, their oxidation state is usually -2. In compounds with non-metals, such as oxygen, it can be or. The exception is oxygen itself. In terms of the magnitude of electronegativity, it is second only to fluorine (see Table 6 on page 118); therefore, only in combination with this element is its oxidation positive. In compounds with all other elements, oxygen oxidation is negative and usually equal to -2. In hydrogen peroxide and its derivatives (see § 117) it is -1. As in the group of halogens, the physical and chemical properties of the elements under consideration naturally change with an increase in the serial number. The appearance of new electronic layers entails an increase in atomic radii, a decrease in electronegativity, a decrease in the oxidative activity of uncharged atoms, and an increase in the reducing properties of atoms with an oxidation state of -2. Table 25. Some properties of oxygen and its analogs Therefore, the non-metallic properties, pronounced in oxygen, are very weakened in tellurium. Some properties of the elements of the main subgroup of the sixth group are given in table. 25.
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