The process of ionization of unpaired electrons. unpaired electron
The structure of an atom determines its radius, ionization energy, electron affinity, electronegativity and other parameters of the atom. The electronic shells of atoms determine the optical, electrical, magnetic, and most importantly, the chemical properties of atoms and molecules, as well as most of the properties of solids.
Magnetic characteristics of an atom
The electron has its own magnetic moment, which is quantized in the direction parallel or opposite to the applied magnetic field. If two electrons occupying the same orbital have oppositely directed spins (according to the Pauli principle), then they cancel each other out. In this case, the electrons are said to be paired. Atoms with only paired electrons are pushed out of the magnetic field. Such atoms are called diamagnetic. Atoms that have one or more unpaired electrons are drawn into a magnetic field. They are called diamagnetic.
The magnetic moment of an atom, which characterizes the intensity of the interaction of an atom with a magnetic field, is practically proportional to the number of unpaired electrons.
Features of the electronic structure of atoms of various elements are reflected in such energy characteristics as ionization energy and electron affinity.
Ionization energy
Energy (potential) of ionization of an atom E i is the minimum energy required to remove an electron from an atom to infinity according to the equation
X = X + + e −
Its values are known for atoms of all elements of the Periodic system. For example, the ionization energy of a hydrogen atom corresponds to the transition of an electron from 1 s- energy sublevel (−1312.1 kJ/mol) to the sublevel with zero energy and equals +1312.1 kJ/mol.
In the change of the first ionization potentials, corresponding to the removal of one electron, of atoms, the periodicity is clearly expressed with an increase in the ordinal number of the atom:
When moving from left to right along the period, the ionization energy, generally speaking, gradually increases, while increasing the serial number within the group, it decreases. The alkali metals have the minimum first ionization potentials, the noble gases have the maximum ones.
For the same atom, the second, third and subsequent ionization energies always increase, since the electron has to be detached from a positively charged ion. For example, for a lithium atom, the first, second, and third ionization energies are 520.3, 7298.1, and 11814.9 kJ/mol, respectively.
The sequence of detachment of electrons is usually the reverse of the sequence of population of orbitals by electrons in accordance with the principle of minimum energy. However, the elements that are populated d-orbitals are exceptions - first of all they lose not d-, a s-electrons.
electron affinity
Affinity of an atom for an electron A e - the ability of atoms to attach an additional electron and turn into a negative ion. The measure of electron affinity is the energy released or absorbed in the process. The electron affinity is equal to the ionization energy of the negative ion X − :
X − = X + e −
The halogen atoms have the highest electron affinity. For example, for a fluorine atom, the addition of an electron is accompanied by the release of 327.9 kJ/mol of energy. For a number of elements, the electron affinity is close to zero or negative, which means that there is no stable anion for this element.
Usually, the electron affinity for atoms of various elements decreases in parallel with an increase in their ionization energy. However, there are exceptions for some pairs of elements:
| Element | Ei, kJ/mol | A e, kJ/mol |
| F | 1681 | −238 |
| Cl | 1251 | −349 |
| N | 1402 | 7 |
| P | 1012 | −71 |
| O | 1314 | −141 |
| S | 1000 | −200 |
An explanation for this can be given based on the smaller sizes of the first atoms and the greater electron-electron repulsion in them.
Electronegativity
Electronegativity characterizes the ability of an atom of a chemical element to displace an electron cloud in its direction when a chemical bond is formed (in the direction of an element with a higher electronegativity). The American physicist Mulliken proposed to define electronegativity as the arithmetic mean between the ionization potential and electron affinity:
χ = 1/2 ( Ei + A e)
The difficulty of applying this method lies in the fact that the values of electron affinity are not known for all elements.
The practical significance of this relation is that, knowing μ, which is relatively easy to measure, one can determine D,
which is difficult to determine directly.
Ambipolar diffusion
Both electrons and ions diffuse in the gas discharge plasma. The diffusion process is represented as follows. Electrons, which are more mobile, diffuse faster than ions. This creates an electric field between the electrons and the stray positive ions. This field slows down the further diffusion of electrons, and vice versa - accelerates the diffusion of ions. When the ions are drawn to the electrons, the indicated electric field weakens, and the electrons are again detached from the ions. This process is ongoing. Such diffusion is called ambipolar diffusion, the coefficient of which is
D amb = |
D e μ and + D and μ e |
|||
μ e + μ and |
||||
where D e ,D and |
are the diffusion coefficients of electrons and ions; μ e , μ and – |
|||
mobility of electrons and ions. |
||||
Since D e >> D and and μ e >> μ and , it turns out that |
D andμ e≈ D e μ and , |
|||
therefore D amb ≈ 2D and . Such diffusion takes place, for example, in the positive column of a glow discharge.
1.6. Excitation and ionization of atoms and molecules
It is known that an atom consists of a positive ion and electrons, the number of which is determined by the number of the element in the periodic table of D.I. Mendeleev. The electrons in an atom are at certain energy levels. If an electron receives some energy from the outside, it goes to a higher level, which is called the excitation level.
Usually, the electron is at the level of excitation for a short time, on the order of 10-8 s. When an electron receives significant energy, it moves away from the nucleus to such a large distance that it can lose contact with it and become free. The least bound to the nucleus are the valence electrons, which are at higher energy levels and therefore more easily detached from the atom. The process of detachment of an electron from an atom is called ionization.
On fig. 1.3 shows the energy picture of a valence electron in an atom. Here W o is the ground level of the electron, W mst is the metastable
ny level, W 1 , W 2 - excitation levels (first, second, etc.).
Part I. Chapter 1. Electronic and ionic processes in a gas discharge
Rice. 1.3. Energy picture of an electron in an atom
W ′ = 0 is the state when the electron loses its connection with the atom. The value W and = W ′ − W o is
with the ionization energy. The values of these levels for some gases are given in table. 1.3.
A metastable level is characterized by the fact that electron transitions to and from it are forbidden. This level is filled with the so-called exchange interaction, when an electron from the outside sits on the W mst level, and the excess
the electron leaves the atom. Metastable levels play an important role in the processes occurring in a gas-discharge plasma, since at the normal level of excitation, the electron is within 10-8 s, and at the metastable level - 10-2 ÷ 10-3 s.
Table 1.3
Energy, eV |
CO2 |
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W mst |
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The process of excitation of atomic particles also determines ionization through the so-called phenomenon of diffusion of resonant radiation. This phenomenon consists in the fact that an excited atom, passing into a normal state, emits a quantum of light, which excites the next atom, and so on. The diffusion region of resonant radiation is determined by the photon mean free path λ ν , which depends
sieves on the density of atomic particles n . So, at n= 1016 cm-3 λ ν =10-2 ÷ 1
see The phenomenon of resonance radiation diffusion is also determined by the presence of metastable levels.
Stepwise ionization can occur according to different schemes: a) the first electron or photon excites a neutral
neutral particle, and the second electron or photon imparts additional energy to the valence electron, causing the ionization of this neutral particle;
Part I. Chapter 1. Electronic and ionic processes in a gas discharge
of a given atom, and at this moment the excited atom goes into a normal state and emits a quantum of light, which increases the energy
c) finally, two excited atoms are close to each other. In this case, one of them goes into a normal state and emits a quantum of light, which ionizes the second atom.
It should be noted that stepwise ionization becomes effective when the concentration of fast electrons (with energies close to
to W and ), photons and excited atoms is quite large. This name-
There is a place when the ionization becomes sufficiently intense. In turn, photons incident on atoms and molecules can also produce excitation and ionization (direct or stepwise). The source of photons in a gas discharge is the radiation of an electron avalanche.
1.6.1. Excitation and ionization of molecules
For molecular gases, it is necessary to take into account the possibility of excitation of the molecules themselves, which, unlike atoms, perform rotational and oscillatory motions. These movements are also quantized. The jump energy during rotational motion is 10-3 ÷ 10-1 eV, and during oscillatory motion - 10-2 ÷ 1 eV.
In the elastic collision of an electron with an atom, the electron loses
much of their energy |
W=2 |
≈ 10 |
− 4 W . When coud- |
|||
When an electron interacts with a molecule, the electron excites the rotational and vibrational motion of the molecules. In the latter case, the electron loses a particularly significant energy up to 10-1 ÷ 1 eV. Therefore, the excitation of oscillatory motions of molecules is an effective mechanism for extracting energy from an electron. In the presence of such a mechanism, the acceleration of the electron is difficult, and a stronger field is required in order for the electron to gain enough energy for ionization. Therefore, the breakdown of a molecular gas requires a higher voltage than the breakdown of an atomic (inert) gas at an equal interelectrode distance and equal pressure. This is shown in the data in Table. 1.4, where the comparison of the values \u200b\u200bof λ t, S t and U pr atom-
molecular and molecular gases at atmospheric pressure and d = 1.3 cm.
Part I. Chapter 1. Electronic and ionic processes in a gas discharge
Table 1.4 |
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Characteristic |
Name of gas |
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S t 10 − 16 , cm2 |
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U pr , kV |
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From Table. 1.4 it can be seen that although the transport cross sections S t for the molecule
polar gases and argon are comparable, but the breakdown voltage of argon is much lower.
1.7. Thermal ionization
At high temperatures, gas ionization can occur due to an increase in the kinetic energy of atomic particles, called thermal ionization. So, for Na, K, Cs vapors, thermal ionization is significant at a temperature of several thousand degrees, and for air at a temperature of about 104 degrees. The probability of thermal ionization increases with an increase in temperature and a decrease in the ionization potential of atoms (molecules). At ordinary temperatures, thermal ionization is negligible and can practically have an effect only during the development of an arc discharge.
However, it should be noted that back in 1951, Hornbeck and Molnar discovered that when monoenergetic electrons are passed through cold inert gases, ions are formed at an electron energy sufficient only for excitation, but not for ionization of atoms. This process has been called associative ionization.
Associative ionization sometimes plays an important role in the propagation of ionization waves and spark discharges in places where there are still very few electrons. Excited atoms are formed there as a result of the absorption of light quanta emerging from already ionized regions. In moderately heated air, at temperatures of 4000 ÷ 8000 K, the molecules are sufficiently dissociated, but there are still too few electrons for the development of an avalanche. The main ionization mechanism in this case is a reaction in which unexcited N and O atoms participate.
Associative ionization proceeds according to the following scheme N + O + 2.8 eV ↔ NO + + q . The missing energy of 2.8 eV is drawn from the kinetic energy of the relative motion of atoms.
FOUNDATIONS OF THEORETICAL CHEMISTRY
5. Chemical bond
According to the theory of chemical bonding, outer shells of two or eight electrons (electronic groupings of noble gases) have the greatest stability. Atoms with fewer than eight (or sometimes two) electrons in their outer shell tend to acquire the structure of noble gases. This pattern allowed V. Kossel and G. Lewis to formulate a position that is fundamental when considering the conditions for the formation of a molecule: “When a molecule is formed during a chemical reaction, atoms tend to acquire a stable eight-electron (octet) or two-electron (doublet) shell.”
The formation of a stable electronic configuration can occur in several ways and lead to molecules (and substances) of various structures, therefore, several types of chemical bonds are distinguished. These are ionic, covalent and donor-acceptor (coordination) communications. In addition to these types of bonds, there are others that are not directly related to electron shells. These are hydrogen and metal bonds.
Valency of elements in compounds.
Modern ideas about the nature of a chemical bond are based on electronic (spin) the theory of valency (the greatest contribution to the development of this theory was made by G. Lewis and W. Kossel), according to which atoms, forming bonds, tend to achieve the most stable (i.e., having the lowest energy) electronic configuration. In this case, the electrons that take part in the formation of chemical bonds are called valence.According to spin theory, the valence of an atom is determined by the number of its unpaired electrons that can participate in the formation of chemical bonds with other atoms, that's why valency is always expressed in small integers.
Consider the electronic configuration of the carbon atom. It has two unpaired electrons in the free state and two paired electrons in the 2s state. Under certain conditions (with the expenditure of a certain amount of energy from the outside), this pair of electrons 2s 2 can be separated (“steered”) by transferring one electron from the 2s state to the state 2p and make these electrons also valence:
In this state, the carbon atom can form compounds where it will be tetravalent.
The process of electron depairing requires a certain amount of energy (D E), and it would seem that it is not profitable. But to take into account the energy ratios, it is necessary to consider the entire balance of bond formation. The point is that during the transition of one of the 2s electrons to the state 2p the state of the atom is obtained, in which it can form not two, but four bonds. When a chemical bond is formed, energy is usually released therefore, the appearance of two new valences leads to the release of additional energy, which exceeds the energy D E spent on the depairing of 2s electrons.
Experiments have proven that the energy spent on depairing electrons within one energy level, as a rule, is fully compensated by the energy released during the formation of additional bonds.
In order to obtain, for example, tetravalent oxygen, trivalent lithium, divalent neon in the same way, a very large expenditure of energy is required.
D E associated with the transition 2p® 3s (oxygen). 1s® 2p(lithium), 2p ® 3s (neon). In this case, the energy consumption is so great that it cannot be compensated for by the energy released during the formation of chemical bonds. Therefore, there are no compounds with a variable valence of oxygen, lithium or neon.Achievements in the chemistry of noble (inert) gases can serve as confirmation of this position. It has long been believed that inert gases do not form chemical compounds (hence
their name). However, in 1962, chemists managed to obtain several compounds of "inert" gases, for example, XeF 2 , XeF 4 , XeO 3 . The manifestation of a certain valency by inert gases can be explained only by assuming that the paired electrons of completely filled sublevels can be depaired within the level.Communication energy. An essential characteristic of a chemical bond is its strength. To assess the strength of bonds, the concept is usually used bond energies.
Bond energy is the work required to break a chemical bond in all the molecules that make up one mole of a substance.
Most often, the binding energy is measured in kJ/mol. Ionic and covalent bonds are the strongest; the energies of these bonds range from tens to hundreds of kJ/mol. The metallic bond, as a rule, is somewhat weaker than ionic and covalent bonds, but the values of the binding energies in metals are close to the values of the energies of ionic and covalent bonds. This is evidenced, in particular, by the high boiling points of metals, for example 357 ° C (Hg) , 880 °C (Na), 3000 ° C (Fe), etc. The energies of hydrogen bonds are very small compared to the energy of interatomic bonds. Thus, the energy of a hydrogen bond is usually 20-40 kJ/mol, while the energy of covalent bonds can reach several hundred kJ/mol.Ionic bond.
An ionic bond is an electrostatic interaction between ions with opposite charges.Kossel suggested that an ionic bond is formed as a result of the complete transfer of one or more electrons from one atom to another. This type of bond is possible only between atoms that differ sharply in properties. For example, elements of groups I and II of the periodic system (typical metals) are directly combined with elements of groups VI and VII (typical non-metals). Examples of ionically bonded substances are MgS, NaCl, A 2 O 3 . Such substances under normal conditions are solid, have high melting and boiling points,their melts and solutions conduct electric current.
The valency of elements in compounds with ionic bonds is very often characterized by the degree of oxidation, which, in turn, corresponds to the charge of the ion of the element in this compound.
The use of the concept of oxidation state for atoms of elements that form other types of chemical bonds is not always correct and requires great care.
covalent bond. It is known that non-metals interact with each other. Consider the formation of the simplest molecule H 2 .
Imagine that we have two separate isolated hydrogen atoms H" and H". When these atoms approach each other, the forces of electrostatic interaction - the force of attraction of the electron of the atom H "to the nucleus of the atom H" and the electron of the atom H "to the nucleus of the atom H" - will increase: the atoms will begin to attract each other. However, at the same time, the repulsive forces between like-charged atomic nuclei and between
the electrons of these atoms. This will lead to atoms will be able to approach each other so much that the forces of attraction will be completely balanced by the forces of repulsion. The calculation of this distance ( covalent bond length) shows that the atoms will approach each other so much that the electron shells involved in the formation of the bond will begin to overlap with each other. This, in turn, will lead to the fact that the electron, which previously moved in the field of attraction of only one nucleus, will be able to move in the field of attraction of another nucleus. Thus, at some point in time around one, then around another atom, a filled shell of a noble gas will arise (such a process can occur only with electrons that have oppositely directed spin projections). In this case, a common pair of electrons arises, simultaneously belonging to both atoms.The overlap region between the electron shells has an increased electron density, which reduces the repulsion between the nuclei and promotes the formation of a covalent bond.
Thus, a bond carried out by the formation of electron pairs equally belonging to both atoms is called covalent.
Communication polarity. A covalent bond can occur not only between the same, but also between different atoms. Thus, the formation of an HCl molecule from hydrogen and chlorine atoms also occurs due to a common pair of electrons, however, this pair belongs to a greater extent to the chlorine atom than to the hydrogen atom, since the non-metallic properties of chlorine are much more pronounced than those of hydrogen.
A kind of covalent bond formed by identical atoms is called non-polar, and formed by different atoms - polar.
Bond polarity is quantified dipole moment
m , which is the product of the dipole lengthl-distance between two equal and opposite charges+q and -q-on the absolute value of the charge:= lh q.The dipole moment is a vector quantity and is directed along the dipole axis from a negative charge to a positive one. It is necessary to distinguish between the dipole moments (polarity) of the bond and the molecule as a whole. So, for the simplest diatomic molecules, the dipole moment of the bond is equal to the dipole moment of the molecule.
On the contrary, in a carbon monoxide (IV) molecule, each of the bonds is polar, and the molecule as a whole is non-polar (
m =0), since the O==C==O molecule is linear, and the dipole moments of the C==O bonds compensate each other (see Fig.). The presence of a dipole moment in a water molecule means that it is non-linear, i.e., O-H bonds are located at an angle not equal to 180 ° (see Fig.).
Electronegativity. Along with dipole moments, another common characteristic called electronegativity is used to assess the degree of ionicity (polarity) of a bond.
Electronegativity is the ability of an atom to attract the valence electrons of other atoms to itself. Electronegativity (EO) cannot be measured and expressed in units of any physical quantities, therefore, several scales have been proposed for the quantitative determination of EO, the most widely recognized and widespread of which was the scale of relative EO developed by L. Pauling.
According to the Pauling scale, the EO of fluorine (the most electronegative of all elements) is conditionally taken equal to
4.0; oxygen is in second place, nitrogen and chlorine are in third. Hydrogen and typical non-metals are at the center of the scale; their ER values are close to 2. Most metals have ER values approximately equal to1.7 or less. EO is a dimensionless quantity.Pauling's EO scale roughly resembles the Periodic Table of the Elements. This scale makes it possible to assess the degree of ionicity (polarity) of a bond. To do this, use the relationship between the difference in EO and the degree of ionicity of the bond.
The greater the EO difference, the greater the degree of ionicity. An EO difference of 1.7 corresponds to a 50% ionic nature of the bonds; therefore, bonds with an EO difference of more than 1.7 can be considered ionic, bonds with a smaller difference are classified as covalent polar.
Ionization energy. The ionization energy is the energy required to detach the most loosely bound electron from an atom. It is usually expressed in electron volts. When an electron is detached from an atom, the corresponding cation is formed.
The ionization energy for elements of the same period increases from left to right with increasing nuclear charge. In a subgroup, it decreases from top to bottom due to an increase in the distance of the electron from the nucleus.
The ionization energy is related to the chemical properties of the elements. So, alkali metals, which have low ionization energies, have pronounced metallic properties. The chemical inertness of noble gases is associated with their high ionization energies.
Affinity for an electron. Atoms can not only donate, but also acquire electrons. In this case, the corresponding anion is formed. The energy released when one electron is attached to an atom is called electron affinity. Usually, the electron affinity, like the ionization energy, is expressed in terms of electron volts. Electron affinities are not known for all elements; it is very difficult to measure them. They are greatest in halogens, which have 7 electrons at the outer level. This indicates an increase in the non-metallic properties of the elements as the end of the period is approached.
The degree of oxidation in covalent compounds. For polar compounds, the concept of oxidation state is also often used, conventionally assuming that such compounds consist only of ions. So, in hydrogen halides and water, hydrogen has a formally positive valence equal to 1+, halogens have a formally negative valence 1-, oxygen has a negative valence 2-: H
+ F -, H + Cl -, H 2 + O 2 -.The concept of the degree of oxidation was introduced under the assumption of a complete displacement of pairs of electrons to one or another atom (showing the charge of the ions that form the ionic compound).
Therefore, in polar compounds, the oxidation state means the number of electrons, only displaced from a given atom to the atom associated with it.The concept of “oxidation state” becomes quite formal when it is used when considering a covalent compound, since the oxidation state is the conditional charge of an atom in a molecule, calculated on the assumption that the molecule consists only of ions. It is clear that in reality there are no ions in covalent compounds.
The difference between the concept of oxidation state and valence in covalent compounds can be especially clearly illustrated on the chlorine derivatives of methane: the valency of carbon is everywhere equal to four, and its oxidation state (counting the oxidation states of hydrogen 1+ and chlorine 1- in all compounds) in each compound is different: 4 - CH 4, 2 - CH 3 Cl, 0 CH 2 Cl 2 , 2+ CHCl 3 , 4+ CCl 4 .
Thus, it must be remembered that oxidation state - conditional, formal concept and, most often, does not characterize the real valence state of an atom in a molecule.
Donor-acceptor bond. In addition to the mechanism for the formation of a covalent bond, according to which a common electron pair arises from the interaction of two electrons, there is also a special pre-nor-acceptor mechanism. It lies in the fact that a covalent bond is formed as a result of the transition of an already existing electron pair donor(electron supplier) for the general use of the donor and acceptor. The donor-acceptor mechanism is well illustrated by the scheme for the formation of an ammonium ion (asterisks indicate the electrons of the outer level of the nitrogen atom):
In the ammonium ion, each hydrogen atom is bonded to the nitrogen atom by a common electron pair, one of which is realized by the donor-acceptor mechanism. It is important to note that H-N bonds formed by various mechanisms do not have any differences in properties, i.e., all bonds are equivalent, regardless of the mechanism of their formation. This phenomenon is due to the fact that at the moment of bond formation, the orbitals of the 2s and 2p electrons of the nitrogen atom change their shape. As a result, four completely identical orbitals arise (here, sp 3 hybridization).
The donors are usually atoms with a large number of electrons, but with a small number of unpaired electrons. For elements of period II, in addition to the nitrogen atom, oxygen (two lone pairs) and fluorine (three lone pairs) have such a possibility. For example, the hydrogen ion H
+ in aqueous solutions never exists in a free state, since from water molecules H 2 O and H ions + hydroxonium ion is always formed H 3 O + The hydronium ion is present in all aqueous solutions, although the symbol H is retained for simplicity. + .The donor-acceptor mechanism of bond formation helps to understand the reason for the amphoteric nature of aluminum hydroxide: in Al(OH) molecules
3 around an aluminum atom there are 6 electrons - an unfilled electron shell. Two electrons are missing to complete this shell. And when an alkali solution containing a large amount of hydroxide ions is added to aluminum hydroxide, each of which has a negative charge and three lone pairs of electrons (OH)- , then the hydroxide ions attack the aluminum atom and form the ion [Al(OH) 4 ] - , which has a negative charge (transferred to it by the hydroxide ion) and a fully completed eight-electron shell around the aluminum atom.
The formation of bonds occurs similarly in many other molecules, even in such “simple” ones as the HNO 3 molecule:
At the same time, the nitrogen atom gives up its electron pair to the oxygen atom, which accepts it: as a result, a completely completed eight-electron shell is reached both around the oxygen atom and around nitrogen, but since the nitrogen atom gave up its pair and therefore owns it together with another atom, it acquired the charge is “+”, and the oxygen atom is the charge “-”. C degree of oxidation nitrogen in HNO 3 equals 5+, while valence equals 4.
Spatial structure of molecules. Ideas about the nature of covalent bonds, taking into account the type of orbitals involved in the formation of a chemical bond, allow us to make some judgments about the shape of molecules.
If a chemical bond is formed using the electrons of s-orbitals, as, for example, in the H 2 molecule , then, due to the spherical shape of the s-orbitals, there is no preferred direction in space for the most favorable bond formation. The electron density in the case of p-orbitals is distributed unevenly in space, so a certain preferred direction appears, along which the formation of a covalent bond is most likely.
Consider examples that allow us to understand the general patterns in the direction of chemical bonds. Let us discuss the formation of bonds in the water molecule H 2 O. The H 2 molecule O is formed from an oxygen atom and two hydrogen atoms. The oxygen atom has two unpaired electrons that occupy two orbitals located at an angle of 90° to each other. Hydrogen atoms have unpaired 1s electrons. It is clear that the angles between two O-H bonds formed by the p-electrons of the oxygen atom with the s-electrons of the hydrogen atoms must be straight or close to it (see Fig.).
Similarly, the angles between bonds in H molecules must be straight.
2 O, H 2 S, F 2 O. Cl 2 O, PH 3, PCl 3 etc. The actual values of the angles between the bonds differ markedly from the theoretical ones.The increase in bond angles (> 90°) can be fully explained by the mutual repulsion of atoms not bonded to each other, which we did not take into account when predicting the angles between bonds. Thus, the mutual repulsion of hydrogen atoms in the molecule H
2 S is weaker than in the H molecule 2 O (since the radius of the sulfur atom is greater than the radius of the oxygen atom), therefore, the H-S-H bond angles are closer to 90 ° than the H-O-H angles.Thus, a divalent non-metal atom with two valence p-orbitals forms a bent (angular, with an angle close to 90°) molecule, and a trivalent atom with three valence p-orbitals forms a pyramid-shaped molecule.
hybridization of orbitals. Consider the formation of a methane molecule CH
4 . An excited carbon atom has four unpaired electrons: one s-electron and three p-electrons - ls 2 2s l 2p 3 .Arguing as in the case of H
2 O, one would think that the carbon atom would form three C-H bonds directed at right angles to each other (p-electrons), and one bond formed by an s-electron, the direction of which would be arbitrary, since the s-orbital has spherical symmetry.Therefore, one would expect that the three C-H bonds in CH
4 are directed p-bonds and are exactly the same, and the fourth bond is an undirected s-s bond and differs from the first three.However, experimental data showed that all four CH bonds in the methane molecule CH
4 are identical and directed towards the vertices of the tetrahedron (the angle between them is 109.5°).Due to the relative closeness of the energy values of 2s- and 2p-electrons, these electrons can interact with each other during the formation of a chemical bond with electrons of another atom, giving four new equivalent hybrid electron clouds.
3 -hybrid orbitals of the carbon atom are located at an angle of 109.5° to each other, they are directed to the vertices of the tetrahedron, in the center of which is the carbon atom. The hybrid orbital is strongly elongated in one direction from the nucleus (see Fig.).
This causes more strong overlap such orbitals with the orbitals of electrons of other atoms compared with the overlap of s- and p-orbitals and leads to the formation of stronger bonds.
Thus, during the formation of a methane molecule, different orbitals of the valence electrons of the carbon atom - one s-orbital and three p-orbitals - turn into four identical “hybrid” sp
3 -orbitals (sp 3 -hybridization). This explains the equivalence of the four bonds of the carbon atom in the molecule.Hybridization turns out to be characteristic not only for compounds of the carbon atom. Hybridization of orbitals can occur when electrons belonging to different types of orbitals simultaneously participate in the formation of bonds.
Consider examples of different types of hybridization s
- and p-orbitals. Hybridization of one s- and one p-orbitals (sp-hybridization) occurs during the formation of beryllium halides, for example, BeF 2 , zinc, mercury, acetylene molecules, etc. The atoms of these elements in the ground state have two paired s-electrons on the outer layer. As a result of excitation, one of the electrons of the s-orbital passes to a p-orbital close in energy, i.e., two unpaired electrons appear, one of which is an s-electron, andanother p-electron. When a chemical bond occurs, these two different orbitals turn into two identical hybrid orbitals (hybridization type - sp), directed at an angle of 180 ° to each other, i.e. these two bonds have the opposite direction (see Fig. rice.).
Experimental determination of the structure of BeX molecules
2 , ZnX 2 , HgX 2 , C 2 H 2 etc. (X - halogen) showed that these molecules are indeed linear.Let us dwell in more detail on the structure of the acetylene C molecule
2 H 2 . In the acetylene molecule, each carbon atom forms two hybridized bonds directed at an angle of 180° to each other (see Fig.).As in making connections
S-S , in the same way, during the formation of C-H bonds, a common two-electron cloud arises, forming s -connections. In general s A bond can be called a bond that occurs when the electron clouds of two atoms are socialized, if the clouds overlap along the line connecting the atoms.
But in the acetylene molecule, the carbon atoms are in sp-hybrid states, i.e., each of the carbon atoms contains two more p-electrons that did not take part in the formation
s -connections. The acetylene molecule has a flat linear skeleton; therefore, both p-electron clouds in each of the carbon atoms protrude from the plane of the molecule in a direction perpendicular to it. In this case, there is also some interaction of electron clouds, but less strong than in the formation s -connections. Thus, two more covalent carbon-carbon bonds are formed in the acetylene molecule, called p -connections (see fig.).
Education case multiple bonds
between carbon atoms for an acetylene molecule - the case of the formation of a triple bond, which consists of one s - and two p-bonds . s - Bonds are stronger than p-bonds.Another type of hybridization of s- and p-orbitals is carried out, for example, in compounds of boron, aluminum or carbon (ethylene benzene). An excited boron atom has one s and two p electrons. In this case, the formation of boron compounds is accompanied by hybridization of one s- and two p-orbitals (ps 2 -hybridization), thus forming three identical sp 2 -hybrid orbitals located in the same plane at an angle of 12 0 ° to each other (see fig.).
Experiments have shown that compounds such as BF
3, AlCl 3, as well as ethylene and benzene have a planar structure and all three bonds B* F(in the BF 3 molecule ) are located at an angle of 120° to each other.Through the formation of sp
2 -hybrid orbitals are explained and the structure of unsaturated hydrocarbons.Hydrogen bond. The very name of this type of bond emphasizes that a hydrogen atom takes part in its formation. Hydrogen bonds can form when a hydrogen atom is bonded to an electronegative atom, which displaces an electron cloud towards itself, thereby creating a positive charge.
d + on hydrogen.The hydrogen bond, like the other types of bonds we have considered, is due to electrostatic interaction, but this interaction is no longer carried out between atoms, but between molecules. Thus, a hydrogen bond is an example of an intermolecular bond.
As an example, consider the formation of a hydrogen bond between two water molecules. O-H bonds in H
2 О have a noticeable polar character with an excess of negative charge d - on the oxygen atom. The hydrogen atom, on the other hand, acquires a small positive charge. d + and can interact with lone pairs of electrons of the oxygen atom of the neighboring water molecule.
The hydrogen bond is usually schematically represented by dots.
The interaction between water molecules turns out to be quite strong, such that even in water vapor there are dimers and trimers of the composition (H
2 O) 2, (H 2 O) 3 etc. In solutions, long chains of associates of the following type may appear:because the oxygen atom has two lone pairs of electrons.
Thus, hydrogen bonds can form if there is a polar X-H bond and a free pair of electrons. For example, molecules of organic compounds containing groups -OH, -COOH, -CONH 2, -NH 2 and others, are often associated due to! formation of hydrogen bonds.
Typical cases of association are observed for alcohols and organic acids. For example, for acetic acid, the occurrence of a hydrogen bond can lead to to combination of molecules in pairs with the formation of a cyclic dimeric structure, and the molecular weight of acetic acid, measured by the vapor density, is doubled (120 instead of 60).
Hydrogen bonds can occur both between different molecules and within a molecule if this molecule contains groups with donor and acceptor abilities. For example, it is intramolecular hydrogen bonds that play the main role in the formation of peptide chains that determine the structure of proteins. Perhaps the most important and undoubtedly one of the best known examples of the influence the intramolecular hydrogen bond on the structure is deoxyribonucleic acid (DNA). The DNA molecule is folded into a double helix. The two strands of this double helix are linked to each other by hydrogen bonds.
Metal connection. Most metals have a number of properties that are of a general nature and differ from the properties of other simple or complex substances. Such properties are relatively high melting points, the ability to reflect light, and high thermal and electrical conductivity. These features are due to the existence in metals of a special type of bond - metallic bond.
In accordance with the position in the periodic system, metal atoms have a small number of valence electrons. These electrons are rather weakly bound to their nuclei and can easily break away from them. As a result, positively charged ions and free electrons appear in the crystal lattice of the metal. Therefore, in the crystal lattice of metals there is a large freedom of movement of electrons: some of the atoms will lose their electrons, and the resulting ions can take these electrons from the “electron gas”. As a consequence of this, the metal is a series of positive ions localized in certain positions of the crystal lattice, and a large number of electrons moving relatively freely in the field of positive centers. This is an important difference between metallic bonds and covalent bonds, which have a strict orientation in space. AT In the case of metals, it is impossible to talk about the direction of the bonds, since the valence electrons are distributed almost uniformly throughout the crystal. This is precisely what explains, for example, the plasticity of metals, i.e., the possibility of displacement of ions and atoms in any direction without breaking the bond.
Lectures for students of general technical areas and specialties lecture 3 Topic 4
Lectures for students of general technical areas and specialties lecture 4 Topic 5
Lectures for students of general technical areas and specialties module II. Reaction patterns
Lectures for students of general technical areas and specialties lecture 7 Topic Fundamentals of chemical kinetics
Lectures for students of general technical areas and specialties lecture 8 Topic Chemical equilibrium On this topic, you need to know and be able to:
Foreword for teachers
Lectures for students of general technical areas and specialties module III. Solutions and electrochemical processes
7. Paired and unpaired electrons
Electrons that fill orbitals in pairs are called paired, and single electrons are called unpaired. Unpaired electrons provide the chemical bond of an atom with other atoms. The presence of unpaired electrons is established experimentally by studying the magnetic properties. Substances with unpaired electrons paramagnetic(they are drawn into the magnetic field due to the interaction of electron spins, like elementary magnets, with an external magnetic field). Substances that have only paired electrons diamagnetic(external magnetic field does not act on them). Unpaired electrons are located only on the outer energy level of an atom and their number can be determined from its electronic graphic scheme.
Example 4 Determine the number of unpaired electrons in a sulfur atom.
Solution. The atomic number of sulfur is Z = 16, therefore, the full electronic formula of the element is: 1s 2 2s 2 2p 6 3s 2 3p 4. The electronic graphic scheme of external electrons is as follows (Fig. 11).
Rice. 11. Electron-graphic scheme of valence electrons of a sulfur atom
It follows from the electron-graphic scheme that there are two unpaired electrons in the sulfur atom.
8. Electron slip
All sublevels have increased stability when they are completely filled with electrons (s 2 , p 6 , d 10 , f 14), and sublevels p, d and f, in addition, when they are half filled, i.e. p 3 , d 5 , f 7 . States d 4 , f 6 and f 13 , on the contrary, have reduced stability. In this regard, some elements have the so-called slip electron, which contributes to the formation of a sublevel with increased stability.
Example 5 Explain why in chromium atoms the 3d sublevel is filled with electrons when the 4s sublevel is not completely filled? How many unpaired electrons are in a chromium atom?
Solution. Chromium atomic number Z = 24, electronic formula: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5. An electron jump from the 4s to the 3d sublevel is observed, which ensures the formation of a more stable state 3d 5 . From the electron-graphic scheme of external electrons (Fig. 12) it follows that there are six unpaired electrons in the chromium atom.
Rice. 12. Electron-graphic scheme of valence electrons of a chromium atom
9. Abbreviated electronic formulas
Electronic formulas of chemical elements can be written in abbreviated form. In this case, the part of the electronic formula corresponding to the stable electron shell of the atom of the previous noble gas is replaced by the symbol of this element in square brackets (this part of the atom is called skeleton atom), and the rest of the formula is written in the usual form. As a result, the electronic formula becomes short, but its information content does not decrease from this.
Example 6 Write the abbreviated electronic formulas for potassium and zirconium.
Solution. Potassium atomic number Z = 19, full electronic formula: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1, the previous noble gas is argon, abbreviated electronic formula: 4s 1.
Zirconium atomic number Z = 40, full electronic formula: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 2, the previous noble gas is krypton, abbreviated electronic formula: 5s 2 4d 2.
10. Families of chemical elements
Depending on which energy sublevel in an atom is filled with electrons last, the elements are divided into four families. In the periodic table, the symbols of elements of different families are highlighted in different colors.
1. s-elements: in the atoms of these elements, the ns-sublevel is the last to be filled with electrons;
2. p-Elements: the np-sublevel is filled with electrons last;
3. d-Elements: the (n - 1) d-sublevel is filled with electrons last;
4. f-Elements: the last to be filled with electrons is the (n - 2) f-sublevel.
Example 7 Using the electronic formulas of atoms, determine which families of chemical elements include strontium (z = 38), zirconium (z = 40), lead (z = 82) and samarium (z = 62).
Solution. We write down the abbreviated electronic formulas of these elements
Sr: 5s 2 ; Zr: 5s 2 4d 2 ; Pb: 6s 2 4f 14 5d 10 6p 2 ; Sm: 6s 2 4f 6 ,
from which it can be seen that the elements belong to the families s (Sr), p (Pb), d (Zr), and f (Sm).
11. Valence electrons
The chemical bond of a given element with other elements in compounds is provided by valence electrons. Valence electrons are determined by the belonging of elements to a particular family. So, for s-elements, the electrons of the outer s-sublevel are valence, for p-elements, the outer sublevels s and p, and for d-elements, the valence electrons are at the outer s-sublevel and pre-external d-sublevel. The question of the valence electrons of the f-elements is not unambiguously resolved.
Example 8 Determine the number of valence electrons in aluminum and vanadium atoms.
Solution. 1) Abbreviated electronic formula of aluminum (z = 13): 3s 2 3p 1. Aluminum belongs to the p-element family, therefore, there are three valence electrons in its atom (3s 2 3p 1).
2) Electronic formula of vanadium (z = 23): 4s 2 3d 3. Vanadium belongs to the family of d-elements, therefore, there are five valence electrons in its atom (4s 2 3d 3).
12. The structure of atoms and the periodic system
12.1. Discovery of the periodic law
The basis of the modern theory of the structure of matter, the study of the whole variety of chemical substances and the synthesis of new elements are the periodic law and the periodic system of chemical elements.
The Periodic Table of Elements is a natural systematization and classification of chemical elements developed by the outstanding Russian chemist D.I. Mendeleev on the basis of the periodic law discovered by him. The periodic system is a graphical representation of the periodic law, its visual expression.
The periodic law was discovered by Mendeleev (1869) as a result of the analysis and comparison of the chemical and physical properties of 63 elements known at that time. Its original wording:
the properties of the elements and the simple and complex substances formed by them are in a periodic dependence on the atomic mass of the elements.
Developing the periodic system, Mendeleev specified or corrected the valency and atomic masses of some known but poorly understood elements, predicted the existence of nine elements not yet discovered, and described the expected properties for three of them (Ga, Ge, Sc). With the discovery of these elements (1875–1886), the periodic law was universally recognized and formed the basis of all subsequent development of chemistry.
For almost 50 years after the discovery of the periodic law and the creation of the periodic system, the very reason for the periodicity of the properties of the elements was unknown. It was not clear why the elements of one group have the same valence and form compounds with oxygen and hydrogen of the same composition, why the number of elements in periods is not the same, why in some places of the periodic system the arrangement of elements does not correspond to an increase in atomic mass (Ar - K, Co - Ni, Te-I). Answers to all these questions were obtained by studying the structure of atoms.
12.2. Explanation of the periodic law
In 1914, the charges of atomic nuclei were determined (G. Moseley) and it was found that element properties are in periodic dependence not from the atomic mass of the elements, but from positive charge of the nuclei of their atoms. But after changing the formulation of the periodic law, the form of the periodic system did not fundamentally change, since the atomic masses of the elements increase in the same sequence as the charges of their atoms, except for the above sequences argon - potassium, cobalt - nickel and tellurium - iodine.
The reason for the increase in the charge of the nucleus with an increase in the number of the element is understandable: in the nuclei of atoms, when moving from element to element, the number of protons monotonically increases. But the structure of the electron shell of atoms with a successive increase in the values of the main quantum number repeats periodically renewal of similar electronic layers. In this case, new electron layers are not only repeated, but also become more complicated due to the appearance of new orbitals, so the number of electrons in the outer shells of atoms and the number of elements in periods increases.
First period: the first energy level, which has only one orbital (orbital 1s), is being filled with electrons, therefore there are only two elements in the period: hydrogen (1s 1) and helium (1s 2).
Second period: the filling of the second electronic layer (2s2p) is in progress, in which the first layer (2s) is repeated and its complication (2p) is in progress - in this period there are 8 elements: from lithium to neon.
Third period: the third electron layer (3s3p) is being filled, in which the second layer is repeated, and there is no complication, since the 3d sublevel does not belong to this layer; there are also 8 elements in this period: from sodium to argon.
The fourth period: the fourth layer (4s3d4p) is being filled with electrons, which is more complicated than the third appearance of five d-orbitals of the 3d-sublevel, so there are 18 elements in this period: from potassium to krypton.
Fifth period: the fifth layer (5s4d5p) is filled with electrons, the complication of which does not occur in comparison with the fourth, therefore there are also 18 elements in the fifth period: from rubidium to xenon.
Sixth period: the sixth layer (6s4f5d6p) is being filled, which is more complicated than the fifth due to the appearance of seven orbitals of the 4f sublevel, so there are 32 elements in the sixth period: from cesium to radon.
Seventh period: the seventh layer (7s5f6d7p), similar to the sixth, is filled with electrons, so there are also 32 elements in this period: from francium to an element with atomic number 118, which has been obtained, but does not yet have a name.
Thus, the patterns of formation of the electron shells of atoms explain the number of elements in the periods of the periodic system. Knowledge of these patterns allows us to formulate the physical meaning of the atomic number of a chemical element in the periodic system, period and group.
atomic number element z is the positive charge of the atomic nucleus, equal to the number of protons in the nucleus, and the number of electrons in the electron shell of the atom.
Period is a horizontal sequence of chemical elements whose atoms have an equal number of energy levels, partially or completely filled with electrons.
The period number is equal to the number of energy levels in atoms, the number of the highest energy level and the value of the main quantum number for the highest energy level.
Group - This is a vertical sequence of elements that have the same type of electronic structure of atoms, an equal number of external electrons, the same maximum valency and similar chemical properties.
The group number is equal to the number of external electrons in atoms, the maximum value of the stoichiometric valency and the maximum value of the positive oxidation state of the element in compounds. By the group number, you can also determine the maximum value of the negative oxidation state of the element: it is equal to the difference between the number 8 and the number of the group in which this element is located.
12.3. Basic forms of the periodic system
There are about 400 forms of the periodic system, but two are most common: long (18-cell) and short (8-cell).
AT long The (18-cell) system (represented in this room and in the handbook) has three short periods and four long periods. In short periods (first, second and third) there are only s- and p-elements, so they have 2 (first period) or 8 elements. In the fourth and fifth periods, in addition to s- and p-elements, 10 d-elements appear each, therefore these periods contain 18 elements each. In the sixth and seventh periods, f-elements appear, so the periods have 32 elements each. But the f-elements are taken out of the table and are given below (in the form of an appendix) in two lines, and their place in the system is indicated by asterisks. The first row contains 14 f-elements that follow lanthanum, so they are collectively called lanthanides, and the second row contains 14 f-elements that follow actinium, so they are collectively called actinides. This form of the periodic system is recommended by IUPAC for use in all countries.
AT short(8-cell) system (it is also available in this classroom and in the reference book), f-elements are also placed in the appendix, and large periods (4th, 5th, 6th and 7th), containing 18 elements each (without f-elements), divided in a ratio of 10:8, and the second part is placed under the first. Thus, large periods consist of two rows (lines) each. In this version, there are eight groups in the periodic system, and each of them consists of a main and secondary subgroup. The main subgroups of the first and second groups contain s-elements, and the rest contain p-elements. The secondary subgroups of all groups contain d-elements. The main subgroups contain 7-8 elements each, and the secondary ones contain 4 elements each, except for the eighth group, in which the secondary subgroup (VIII-B) consists of nine elements - three "triads".
In this system, the elements of subgroups are full electronic counterparts. Elements of the same group, but different subgroups, are also analogues (they have the same number of outer electrons), but this analogy is incomplete, because outer electrons are on different sublevels. The short form is compact and therefore more convenient to use, but it does not have that one-to-one correspondence between the shape and the electronic structure of atoms that is inherent in the long system.
Example 9 Explain why chlorine and manganese are in the same group but in different subgroups of the 8-cell periodic table.
Solution. The electronic formula of chlorine (atomic number 17) is 3s 2 3p 5, and manganese (atomic number 25) is 4s 2 3d 5. The atoms of both elements have seven outer (valence) electrons, so they are in the same group (seventh), but in different subgroups, since chlorine is
p-element, and manganese - d-element.
12.4. Periodic properties of elements
Periodicity is expressed in the structure of the electron shell of atoms, therefore, properties that depend on the state of electrons are in good agreement with the periodic law: atomic and ionic radii, ionization energy, electron affinity, electronegativity and valency of elements. But the composition and properties of simple substances and compounds depend on the electronic structure of atoms, therefore, periodicity is observed in many properties of simple substances and compounds: the temperature and heat of melting and boiling, the length and energy of a chemical bond, electrode potentials, standard enthalpies of formation and entropy of substances, etc. d. The periodic law covers more than 20 properties of atoms, elements, simple substances and compounds.
1) Atomic and ionic radii
According to quantum mechanics, an electron can be located at any point around the nucleus of an atom, both near it and at a considerable distance. Therefore, the boundaries of atoms are vague, indefinite. At the same time, quantum mechanics calculates the probability of distribution of electrons around the nucleus and the position of the maximum electron density for each orbital.
Orbital radius of an atom (ion)is the distance from the nucleus to the maximum electron density of the most distant outer orbital of this atom (ion).
Orbital radii (their values are given in the handbook) decrease in periods, because an increase in the number of electrons in atoms (ions) is not accompanied by the appearance of new electron layers. The electron shell of an atom or ion of each subsequent element in the period becomes denser compared to the previous one due to an increase in the charge of the nucleus and an increase in the attraction of electrons to the nucleus.
Orbital radii in groups increase as an atom (ion) of each element differs from the parent by the appearance of a new electronic layer.
Change of orbital atomic radii for five periods is shown in fig. 13, from which it can be seen that the dependence has a “sawtooth” form characteristic of the periodic law.
Rice. 13. Dependence of the orbital radius
But in periods, the decrease in the size of atoms and ions does not occur monotonically: individual elements have small “bursts” and “dips”. In the "dips" there are, as a rule, elements whose electronic configuration corresponds to a state of increased stability: for example, in the third period it is magnesium (3s 2), in the fourth - manganese (4s 2 3d 5) and zinc (4s 2 3d 10) etc.
Note. Calculations of orbital radii have been carried out since the mid-seventies of the last century due to the development of electronic computers. Previously used effective the radii of atoms and ions, which are determined from experimental data on internuclear distances in molecules and crystals. It is assumed that the atoms are incompressible balls that touch their surfaces in compounds. The effective radii determined in covalent molecules are called covalent radii, in metal crystals - metal radii, in compounds with ionic bond - ionic radii. The effective radii differ from the orbital radii, but their change depending on the atomic number is also periodic.
2) Energy and ionization potential of atoms
Ionization energy(E ion) is called the energy expended in detaching an electron from an atom and turning the atom into a positively charged ion.
Experimentally, the ionization of atoms is carried out in an electric field by measuring the potential difference at which ionization occurs. This potential difference is called ionization potential(J). The unit of measurement of the ionization potential is eV/atom, and the ionization energy is kJ/mol; the transition from one value to another is carried out according to the relation:
E ion = 96.5 J
The detachment of the first electron from the atom is characterized by the first ionization potential (J 1), the second - by the second (J 2), etc. Successive ionization potentials increase (Table 1), since each subsequent electron must be detached from an ion with a positive charge increasing by one. From Table. Table 1 shows that for lithium a sharp increase in the ionization potential is observed for J 2 , for beryllium for J 3 , for boron for J 4 , etc. A sharp increase in J occurs when the detachment of the outer electrons ends and the next electron is at the pre-outer energy level.
Table 1
Ionization potentials of atoms (eV/atom) of elements of the second period
| Element | J1 | J2 | J3 | J4 | J5 | J6 | J7 | J8 |
| Lithium | 5,39 | 75,6 | 122,4 | – | – | – | – | – |
| Beryllium | 9,32 | 18,2 | 158,3 | 217,7 | – | – | – | – |
| Bor | 8,30 | 25,1 | 37,9 | 259,3 | 340,1 | – | – | – |
| Carbon | 11,26 | 24,4 | 47,9 | 64,5 | 392,0 | 489,8 | – | – |
| Nitrogen | 14,53 | 29,6 | 47,5 | 77,4 | 97,9 | 551,9 | 666,8 | – |
| Oxygen | 13,60 | 35,1 | 54,9 | 77,4 | 113,9 | 138,1 | 739,1 | 871,1 |
| Fluorine | 17,40 | 35,0 | 62,7 | 87,2 | 114,2 | 157,1 | 185,1 | 953,6 |
| Neon | 21,60 | 41,1 | 63,0 | 97,0 | 126,3 | 157,9 |
The ionization potential is an indicator of the “metallicity” of an element: the smaller it is, the easier it is for an electron to detach from an atom and the stronger the metallic properties of the element should be expressed. For elements with which periods begin (lithium, sodium, potassium, etc.), the first ionization potential is 4–5 eV/atom, and these elements are typical metals. For other metals, the values of J 1 are greater, but not more than 10 eV / atom, and for non-metals usually more than 10 eV / atom: for nitrogen 14.53 eV / atom, oxygen 13.60 eV / atom, etc.
The first ionization potentials increase in periods, and decrease in groups (Fig. 14), which indicates an increase in non-metallic properties in periods and metallic properties in groups. Therefore, non-metals are in the upper right part, and metals are in the lower left part of the periodic table. The boundary between metals and non-metals is "blurred", because most elements have amphoteric (dual) properties. Nevertheless, such a conditional boundary can be drawn, it is shown in the long (18-cell) form of the periodic system, which is available here in the classroom and in the reference book.
Rice. 14. Dependence of the ionization potential
from the atomic number of the elements of the first - fifth periods.
The discovery of radioactivity confirmed the complexity of the structure of not only atoms, but also their nuclei. In 1903, E. Rutherford and F. Soddy proposed the theory of radioactive decay, which radically changed the old views on the structure of atoms. According to this theory, radioactive elements spontaneously decay with the release of α- or β-particles and the formation of atoms of new elements, chemically different from the original ones. At the same time, the stability of the mass of both the initial atoms and those that were formed as a result of the course of the decay process is preserved. E. Rutherford in 1919 was the first to investigate the artificial transformation of nuclei. During the bombardment of nitrogen atoms with α-particles, he isolated the nuclei of hydrogen atoms (protons) and atoms of the oxygen nuclide. Such transformations are called nuclear reactions, since the nuclei of atoms of other elements are obtained from the nuclei of atoms of one element. Nuclear reactions are written using equations. So, the nuclear reaction discussed above can be written as follows:
Definitions of the phenomenon of radioactivity can be given using the concept of isotopes: radioactivity is the transformation of unstable nuclei of atoms of one chemical element into the nuclei of atoms of another element, which is accompanied by the release of elementary particles. The radioactivity exhibited by isotopes of elements that exist in nature is called natural radioactivity. The rate of radioactive transformations is different for different isotopes. It is characterized by the radioactive decay constant, which shows how many atoms of a radioactive nuclide decay in 1 s. It has been established that the number of atoms of a radioactive nuclide that decays per unit time is proportional to the total number of atoms of this nuclide and depends on the magnitude of the radioactive decay constant. For example, if during a certain period half of the total number of atoms of a radioactive nuclide decayed, then in the next such period half of the remainder will decay, that is, half as much as in the previous period, and so on.
The lifespan of a radioactive nuclide is characterized by its half-life, that is, such a period of time during which half of the initial amount of this nuclide decays. For example, the half-life of Radon is 3.85 days, Radium - 1620 years, Uranus - 4.5 billion years. Known such types of radioactive transformations: α-decay, β-decay, spontaneous (unauthorized) nuclear fission. These types of radioactive transformations are accompanied by the release of α-particles, electrons, positrons, γ-ray. In the process of α-decay, the nucleus of an atom of a radioactive element releases the nucleus of a Helium atom, as a result of which the charge of the nucleus of an atom of the original radioactive element decreases by two units, and the mass number - by four. For example, the transformation of a Radium atom to a Radon atom can be written by the equation
The nuclear reaction of β-decay, which is accompanied by the release of electrons, positrons, or the drag of orbital electrons, can also be written by the equation
where e is an electron; hν - quantum of γ-radiation; ν o - antineutrino (an elementary particle whose rest mass and charge are equal to zero).
The possibility of β-decay is due to the fact that, in accordance with modern concepts, a neutron can turn into a proton under certain conditions, releasing an electron and an antineutrino. Proton and neutron are two states of the same nuclear particle - nucleon. This process can be represented by a diagram
Neutron -> Proton + Electron + Antineutrino
In the process of β-decay of atoms of a radioactive element, one of the neutrons, which is part of the atomic nucleus, releases an electron and an antineutrino, turning into a proton. In this case, the positive charge of the nucleus increases by one. This type of radioactive decay is called electronic decay (β - decay). So, if the nucleus of an atom of a radioactive element releases one α-particle, the nucleus of an atom of a new element is obtained with a proton number two units less, and when a β-particle is released, the nucleus of a new atom with a proton number one greater than that of the original one. This is the essence of the Soddy-Faience displacement law. The atomic nuclei of some unstable isotopes can release particles that have a positive +1 charge and a mass close to that of an electron. This particle is called a positron. So, the possible transformation of a proton into a neutron according to the scheme:
Proton → Neutron + Positron + Neutrino
The transformation of a proton into a neutron is observed only when the instability of the nucleus is caused by an excess content of protons in it. Then one of the protons turns into a neutron, and the positron and neutrinos that arise in this case fly out of the boundaries of the nucleus; the nuclear charge is reduced by one. This type of radioactive decay is called positron decay (β+ decay). So, due to the β-decay of the nucleus of an atom of a radioactive element, an atom of an element is obtained that is shifted one place to the right (β-decay) or to the left (β+-decay) from the original radioactive element. A decrease in the charge of the nucleus of a radioactive atom by one can be caused not only by β+ decay, but also by electron drag, as a result of which one of the electrons of the electron ball closest to the nucleus is captured by the nucleus. This electron with one of the protons of the nucleus forms a neutron: e - + p → n
The theory of the structure of the atomic nucleus was developed in the 30s of the XX century. Ukrainian scientists D.D. Ivanenko and E.M. Gapon, as well as the German scientist W. Heisenberg. According to this theory, the nuclei of atoms are composed of positively charged protons and electrically neutral neutrons. The relative masses of these elementary particles are almost the same (the proton mass is 1.00728, the neutron mass is 1.00866). Protons and neutrons (nucleons) are contained in the nucleus by very strong nuclear forces. Nuclear forces act only at very small distances - about 10 -15 m.
The energy that is released during the formation of a nucleus from protons and neutrons is called the binding energy of the nucleus and characterizes its stability.
