The concept of the method of molecular orbitals (MO). Open Library - open library of educational information Sun and mo method in chemistry
The theory of valence bonds (MVS) as applied to complex compounds was developed by L. Pauling in 1930. At present, it is used relatively rarely, but it served well for about a quarter of a century in the chemistry of coordination compounds to explain some properties of complexes (spatial structure, magnetic properties). Despite the cumbersomeness of quantitative calculations, there are big problems in interpreting various distortions of octahedral complexes, the absence predictive ability even in cases of highly symmetrical geometric structure of coordination spheres and other shortcomings, AIM remains a handy tool that allows clearly in terms of quality to explain the fact of formation of complexes, which makes it possible to evaluate mutual preferences for binding, the predisposition of complexes to hydrolysis, polycondensation, to predict the composition and some properties of carbonyls and related compounds, and, of course, to explain, and in many cases even predict, the magnetic properties of complexes.
The main provisions of the MVS regarding the structure of the complexes are formulated as follows:
1. The bond between the complexing agent and the ligands is established by the donor-acceptor mechanism, and in σ –bonds The ligand is an electron pair donor (“Lewis acid”), the central atom is an acceptor (“Lewis base”).
2. A measure of bond strength is the degree of orbital overlap. To explain the fact of the formation of strong bonds at a very specific spatial arrangement of ligands around the central atom, which often does not coincide with the spatial arrangement of its own vacant AOs, the concept of hybridization of the complexing agent involved in σ – binding. The type of hybridization is determined by the number and nature of the central atom and ligands. The nature of hybridization determines the geometric shape of the complex.
3. Additional strengthening of the complex is due to the appearance of an additional π – binding. In this case, the electropositive atom of the complexing agent often acts as a donor, while the more electronegative atom, due to which the ligand is coordinated, acts as an acceptor. This donor-acceptor interaction is called dative .
4. The magnetic properties exhibited by the complex are explained by the peculiarities of the population of the orbitals of the complexing agent by electrons. In the presence of unpaired electrons the complex is paramagnetic. The complete absence of unpaired electrons causes complex compound diamagnetism. Approximate value of the magnetic moment μ (in Bohr magnetons, μ V) can be calculated using the formula
, (4.10 )
where n is the number of unpaired electrons.
Before considering several examples of the use of MHS for the analysis of the structure and properties of a number of complexes,
 |
it is useful to recall some information about the electronic structure, valence capabilities of potential complexing agents and ligands, as well as comment on certain provisions of L. Pauling's theory.
Atoms of the second period, acting as complexing agents (Be, B), and therefore, establishing covalent bonds with ligands to a significant extent, are limited in the maximum achievable CN, because at the valence energy level have only four orbitals (2 s - and 2 R -). Elements of III and large periods have vacant n d -orbitals and due to them they can show increased acceptor properties (increase the CN up to 6 or more, establish additional π - bonds with ligands σ - and π -donors). However, as noted earlier (Chapter 1.5), the energy n d – the orbitals are quite large. At the same time, their energy advantage for electrons is enhanced when the atom under consideration is bound to strongly electronegative elements (especially with F - , and ligands coordinated by oxygen atoms: O 2–, OH - , OH 2, etc.). For the first time, an assumption about the possibility of using external d -orbitals was stated in 1937. Higgins, and later it found a calculated confirmation.
Atoms of transition elements have, besides, also (n-1) d -orbitals that are much more valent than n d –orbitals, especially for the first elements of decades, especially in low positive oxidation states. As you fill (n-1) d -orbitals by electrons, their acceptor capabilities are weakened (the probability of using n d -orbitals), but donor properties increase and, accordingly, preferences for binding to ligands increase. σ – donors and π -acceptors.
To distinguish between two types of complexes, the concepts were introduced: outer orbit and intraorbital(Taube), spin-free and spin-coupled(Newholm), high-spin and low spin(Orgel).
The atomic orbitals participating in the covalent bonding must be comparable in energy and correspond to each other in symmetry: they should be arranged in such a way as to provide overlapping areas in which the signs of the wave functions coincide. Because the s – the valence orbital usually has the lowest energy, it is almost always used in bonding, but due to spherical symmetry, it cannot participate in π - overlapping, and σ – interaction can support in any direction (including being involved in hybridization processes). Symmetry R –orbitals allows them to participate both in σ - as well as in π -overlappings. As part of the central atom to maintain its CN (more than one: 6, 4, less often others) R Orbitals are pre-hybridized with s - and, if necessary, with d -orbitals. In addition, the symmetry R π –binding (usually, in the composition of donor atoms of ligands). At high CN (4 and above) in σ –binding may be involved and d -orbitals of suitable symmetry (in squares and octahedrons - located in petals along the rectangular coordinate axes d x 2 - y 2, d z 2 , and in the case of a tetrahedral environment - located along the bisectors of the coordinate angles d xy, d xz , d yz). For reasons to be explained later, the first two orbitals have the group designation d γ (or e g), while the other three are d ε (or t 2g). Symmetry d -orbitals allows them to participate in π – interaction, and, due to some orientation towards a potential partner, they can provide a stronger overlap of electron clouds than what is achieved when used in π -connections R –orbitals of comparable energy (similar in size).
Table 4.11
The shape and relative strength of hybrid bonds ( E* )
The most common types of hybridization, the corresponding (calculated) geometric shapes of the complexes, as well as the relative strength σ –bonds formed with the help of the corresponding hybrid orbitals are shown in Table 4.11.
As for the third point, the reason for postulating this provision was examples of strong binding of some 4 d – and 5 d –elements with ligands whose donor properties are rather weakly expressed. For example, Pt(II), Hg(II), Au(III) bind better with large halide ions than with F - ; they also form fairly strong complexes with: PF 3 and ∶P(C 6 H 5) 3, but do not bind to ∶РН 3 at all (recall that the ∶РН 3 molecule is very reluctant to bind to such an active electron pair acceptor as H + ). These facts were explained by Pauling for several reasons, one of which is the growing multiplicity of the connection due to the additional dative π – interactions of complexing agents with configurations d 8 , d 10 s d -orbitals of Cl, Br, J, P atoms. In turn d –phosphorus orbitals are more actively involved in binding in the composition of such ligands, where their energy is lowered under the influence of their own intraligand strongly electronegative atoms (F) or groups (C 6 H 5).
The existence of various forms of additional π –M–L binding was further supported by a variety of examples. The most important types of π-interactions in complexes can be systematized as follows (Fig. 4.26):
a) π d (M) → p (L) : partial transfer of electrons from d R are the ligand orbitals;
b) π d (M) → d (L) : partial transfer of electrons from d –metal orbitals to vacant d are the ligand orbitals;
in) π p (M) ← p (L) : partial transfer of electrons with R R –metal orbitals;
G) π d (M) ← p (L) : partial transition of electrons with R – ligand orbitals to vacant d metal orbitals.
Now it is possible to consolidate the application of the analyzed provisions of Pauling's theory on specific examples, while analyzing them, we will also consider the magnetic properties of the complexes. Let us first discuss the composition, structure, and some properties of complex compounds. d -metals.
For the first d -elements are characterized by the highest positive oxidation states. This formally means that a fully ionized atom with many empty orbitals acts as a complexing agent and, accordingly, it should preferentially bind to ligands. σ - and π – donors. In particular, for the most stable Ti 4+ complexes are fluoride (to a lesser extent, other halide) and oxygen-containing complexes. If polymeric compounds are not taken into account, then these are the anionic 2– complex and the cationic 2+ complex (the "4+" avocomplex hydrolyzes very strongly under the strong polarizing effect of the central atom; in the +III oxidation state, the aquocomplex is hydrolyzed to a much lesser extent: 3+). Electronic configuration Ti 4+ : 3d 0 4s 0 4p 0 , in σ –binding with ligands involved d 2 sp 3-hybrid orbitals, empty d ε can be involved in an additional multicenter π d (M) ← p (L) - binding:
The Ti 4+ and Ti 3+ cationic complexes are also intraorbital and have octahedral symmetry, but unlike 3+ (and 2–), dihydroxo-diaquotitanium (IV) has a distorted structure: bonds with hydrodroxo groups are shorter than those with water molecules ( CC = 2+4). This can be explained by the unequal π -binding (stronger π -donor properties of OH ions -). At the same time, 3+ is a paramagnetic particle, while 2– and 3– 2+ and + , however, these species (especially the latter) easily enter into substitution reactions for F - or (less willingly) for oxygen-containing ligands:
3– , 3+ , 3–)
· all Cr 3+ complexes must be paramagnetic, because the complexing agent has three electrons;
(magnetic moments of all Cr 3+ complexes correspond to the presence
three unpaired electrons).
Note that due to the partial population d ε orbitals, Cr 3+ cannot show any π -acceptor properties (as part of 3+), nor π -donor (composed of 3-). Curiously, cyanide complexes (carbonyls and other complexes with active ligands) π -acceptors) often exhibit high electron affinity, which makes it possible to stabilize γ in the composition of such compounds. d -elements have abnormally low (sometimes even negative) oxidation states. In particular, K 3 can be reduced to K 6 by reaction with atomic hydrogen (zinc in hydrochloric acid medium). Moreover, as part of the new complex, the chromium atom accepts three additional electrons into its orbitals and, acquiring a zero oxidation state, should thereby reproduce the electronic configuration of the neutral atom 3d 5 4s 1 4p 0 with six unpaired electrons. However, the K 6 complex is diamagnetic. Similar facts gave grounds to assume that in complexes with active π -acceptors change the electronic structure of the complexing agent: d -sub-level is first populated d ε orbitals (at first in accordance with the Hund rule, and with configurations d 4 , d 5 and d 6 - in pairs). This allows, firstly, to save (n-1) d γ orbitals vacant and use them for intraorbital hybridization and σ -bindings, and secondly, pairwise filled d ε orbitals can be used for additional π d (M) → p (L) - interactions, which leads to an increase in the multiplicity of the connection complexing agent–ligand . Taking these considerations into account, the formation of the 6– complex from the point of view of the MHS can be schematically shown as follows:
The peculiarity of cyanide ions as ligands is also confirmed by a comparison of chromium (II) complexes: with the same electronic structure of the central atom ( d 4) magnetic moments 4– , on the one hand, and 2+ , 4– , 4– ,…, on the other hand, differ:
At the same time, VHS turns out to be powerless before explaining the differences in optical properties (coloration) and details of the spatial structure: unlike the cyanide complex, all the others, despite the homogeneous ligand composition and the equivalence of the σ – binding of hybridized orbitals of the central atom, are characterized by a weak tetragonal distortion of octahedral coordination (cn = 4+2).
With a further increase in the charge of the nucleus and a simultaneous increase in the number of electrons in valence orbitals, the following is observed:
ü increasing stabilization of low oxidation states d –elements;
ü amplification π -donor properties of atoms (ions) d -metals. Accordingly, the interaction with ligands gradually weakens. σ - and π donors, there is a growing preference for binding to ligands π -acceptors, as a result - the complexes become more diverse;
ü Gradual transition to outer orbit complexes.
Let us consider some Ni(II), Cu(II), and Cu(I) complexes.
Cu (II) complexes are very diverse in ligand composition: the list of only monodentate ligands, upon binding with which island complexes can be obtained, includes H 2 O, OH - , G - , NH 3 , SCN - , S 2 O 3 2 -, NO 2 - etc. Their color is very diverse: blue, yellow-green, blue-violet, .... At the same time, the magnetic properties of the complexes are the same, and their structures are similar or related:
- at electronic configuration of the central atom d 9 in all complexes of the Cu 2+ ion, one unpaired electron is found;
– in most complexes, tetragonally distorted octahedral coordination is realized (cn = 4+2); sometimes both or one of the weakly bound ligands completely leave the coordination sphere (in this case, either square ones are obtained - CN = 4 ( no tetrahedra!), or square-pyramidal complexes - CC=4+1):
| CN = 4+2 (prolonged octahedron) | CV = 4+1 (square pyramid) | CV = 4 (square) |
| 2+ , 4– , 2+ , 2+ , 4– | 3 – , 2+ | 2– , 2+ , 2– , 2– |
From the point of view of the MVS, all Cu(II) complexes are outer-orbital:
Recall that for the formation of electron clouds oriented towards the vertices of a square pyramid, the orbital d x 2 - y 2 . It is also necessary for the formation of flat-square complexes, while R z-orbital from hybridization is extracted. In addition, it should be noted that, in accordance with the MVS, in the chloride and hydroxo complex, a weak additional π -binding (Cl ions - are weak π – donors and π -acceptors; OH ions - have much more pronounced π -donor properties, but the central atom π -acceptor properties can be realized only at the expense of high-lying 4 d -orbitals). Despite explaining the modes of covalent interaction of the central atom and ligands, the MHS is still powerless to suggest the reasons for both the spectral activity and the structural features of the complexes. Curiously, Cu(I) complexes, on the contrary, are overwhelmingly colorless, but much more structurally diverse, despite the lower coordination numbers (CN: 2, 3, 4; coordination forms: line, triangle, tetrahedron - no squares!):
As for complexes s - and R -elements, we briefly note only some important regularities:
Ions of elements act as complexing agents (see Table 4.7) with an intermediate polarizing effect (electronegativity), however, it is important to understand that for most of the elements under consideration, these characteristics are noticeably higher than for d –metals;
Almost all potential complexing agents form only octahedral complexes (in Be 2+, B 3+ only tetrahedra are known; Al 3+ and Ga 3+, along with octahedrons, also sometimes form tetrahedral complexes; Sn 2+, Pb 2+ have only tetrahedral and trigonal-pyramidal complexes), which requires involvement in hybridization and σ –interaction n d γ orbitals (due to s - and R -orbitals can only be realized CN = 4). This suggests bonding with strongly electronegative atoms, and also that due to vacant n d ε orbitals potential complexing agents are quite active π -acceptors.
The ligands in the vast majority of cases are active σ - and π - donors: OH 2 (only when bound to ions that do not cause strong hydrolysis, i.e. p/d , which are minimal in a given series of elements), OH - (when bonding with ions characterized by an intermediate level p/d in the series of these elements), monatomic ligands: O 2–, F –. R –elements of the VIth, Vth and, to a lesser extent, IVth periods have filled (n-1) d 10 - sublevels and, therefore, can participate in π d (M) → d (L) - interaction. Accordingly, for such elements, even in the aquatic environment, it may turn out to be quite competitive, beneficial relationships M–Cl and Cl as a potential ligand. In some cases, complexes with larger halogens are also stabilized. The same elements, but much less frequently, can form island water-soluble complexes with S2– and SH– ligands.
All complexes R -elements are diamagnetic and the vast majority are colorless. Extremely rare exceptions are possible in the case of complexes with ligands π -acceptors.
Table 4.12
Compositions of the most important island
water-soluble complexes R -elements
| IIa | IIIa | IVa | Va | VIа |
| 2+ 2– 2– | – 2– – | –– | –– | –– |
| 2+ | 3+ – – 3– 3– | 2– 2– | – | –– |
| Same as Al 3+ | 2– 2– 2– | – AsO 4 3– ; – – – | –– | |
| Same as Al 3+ , except hydroxocomplex, additionally – 3– | 2+ 2– 2– 2– ; 2+ – – | – – – ; + – 3– – | ; 2– 2– | |
| 3– ; 2– | 2– 2– ; 2+ – – | 3+ – | 2+ [Ро(OH) 6 ] 2– [РоCl 6 ] 2– ; [Po(OH 2) 6] 2+ |
In conclusion, let us briefly discuss the application of the ideas of MHS to explain the composition, structure, and some properties of rather peculiar compounds: carbonyls and carbonyl complexes. d –elements (polyligand carbonyls are also known: carbonylnitrosyls (M(CO) x (NO) y), carbonyl halides (M(CO) x Г y), carbonyl hydrides (M(CO) x H y), carbonyl-metalocenes (M( CO) x (C 5 H 5) y), etc., including polynuclear ones containing several atoms d -metal). The composition of most of them obeys the rules formulated in the 1920s. at the turn of the formation of a quantum mechanical model of the structure of the atom: the first and modified Sidgwick's rule(18 electron rule ): the most stable are complexes in which the central atom has a completely completed(n-1) d 10n s 2n p 6 -configuration. Valence electrons are taken into account d – elements and electrons of ligands involved in bonds M–L . The rule is based on the assumption of pairwise population of valence orbitals by electrons of the central atom and donor-acceptor interaction complexing agent-ligand (radical ligands, such as NO, are considered as donors of three electrons; ligands with extended π -systems are donors of all their π -electrons).
Table 4.13
Composition of known carbonyls 3 d -elements
The scope and subject matter of this textbook do not allow an analysis of the possible reasons that limit the range of elements prone to the formation of carbonyls (Table 4.14). We only note that, taking into account related compounds, they were obtained for all d - metals except
Table 4.14
A circle d -elements that make up carbonyls
| sc | Ti | V | Cr | Mn | Fe | co | Ni | Cu | Zn |
| Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag | CD |
| La | hf | Ta | W | Re | Os | Ir | Pt | Au | hg |
Nb, Ta, as well as elements of scandium and zinc subgroups. At the same time, the composition and structures of the simplest carbonyls are in perfect agreement with the Sidgwick rule and Pauling's theory. In particular, the alternation of monomeric (for Cr, Fe and Ni) and dimeric molecules (for V, Mn and Co) is the result of the fact that the elements of odd groups have an odd number of valence electrons, therefore monomeric molecules are radicals and are able to combine due to the bond MM (Such connections are called clusters ):
ü Examples of substantiating the composition based on the modified Sidgwick rule:
ü structures based on Pauling theory
Cr Cr = 6 Fe = 5 Ni = 4
octahedron trigonal tetrahedron
bipyramid
d 2 sp 3 dsp 3 sp 3
CN Mn = 6 CN Co = 4+1
octahedron trigonal
bipyramid
d 2 sp 3 dsp 3
Fe, Co and some heavy d -metals are known "complex carbonyls". Convincing explanations for their composition and selective existence have not yet been worked out. At the same time, the features of their structure (the presence of connections MM , the number of bridging or terminal CO molecules, spatial environment) can be predicted using the Sidgwick / Pauling theory (see, for example, the textbook by J. Huey "Inorganic chemistry. The structure of matter and reactivity").
The VS method is based on the following main provisions:
a) a chemical bond between two atoms arises as a result of overlapping AO with the formation of electron pairs (generalized two electrons);
b) atoms forming a chemical bond exchange electrons with each other, which form bonding pairs. The energy of the exchange of electrons between atoms (the energy of attraction of atoms) contributes to the energy of the chemical bond. An additional contribution to the binding energy comes from the Coulomb forces of particle interaction;
c) electrons with antiparallel spins participate in the formation of a chemical bond;
d) the characteristics of the chemical bond (energy, length, polarity, etc.) are determined by the type of AO overlap.
The electronic structure of a molecule differs significantly from the electronic structure of its constituent atoms. For example, the electron orbitals in the hydrogen molecule do not have spherical symmetry, unlike the AO of the hydrogen atom, since the electron pair belongs to a two-center molecular system. At the same time, this bonding electron pair is at a lower energy level than the unpaired electrons of hydrogen atoms.
As a result of the formation of molecules from atoms, only the electronic structure of the outer and pre-outer shells of atoms undergoes changes. Therefore, atoms with the initial electronic structure do not exist in the resulting molecule. Atoms in a molecule retain only the electronic configurations of their inner electron shells, which do not overlap when bonds are formed.
The ability of an atom to attach or replace a certain number of other atoms to form chemical bonds is called valency. According to the VS method, each atom donates one unpaired electron to form a common electron pair (covalent bond). The quantitative measure of valence in the exchange mechanism of the VS method is the number of unpaired electrons in an atom in the ground or excited state. These include unpaired electrons in the outer shells of atoms s- and R-elements, outer and pre-outer shells d- elements.
When a chemical bond is formed, an atom can go into an excited state as a result of the separation of a pair or pairs of electrons and the transition of one (or several electrons equal to the number of disconnected pairs) to a free orbital of the same shell. For example, the electronic configuration of calcium in the ground state is written as 4s 2 . In accordance with the exchange mechanism of the VS method, its valence is equal to zero, i.e. for Sa (…4s 2) valency B=0. At the calcium atom in the fourth shell (n=4) there are vacant R- orbitals. When an atom is excited, electrons are depaired and one of 4s- electrons goes to free 4s- orbital. The valency of calcium in the excited state is equal to two, i.e. when steaming, the valence increases by two units.
| 4s | 4p | 4s | 4p | |||||||||||||||
| Ca | ↓ | → | Ca* | B*=2 | ||||||||||||||
Unlike oxygen and fluorine, whose electron pairs cannot be separated, because there are no other vacant orbitals on the second level, the electron pairs of sulfur and chlorine can be steamed out, because the third level has vacant 3d orbitals. Accordingly, sulfur, in addition to the valency of the ground state I and II,
| 3s | 3p | 3d | ||||||||
| ↓ | ↓ |
it also has valencies IV and VI in excited states:
| 3s | 3p | 3d | ||||||||
| ↓ | ||||||||||
Spatial structure of molecules .
As shown earlier, a covalent chemical bond is characterized by directionality, which is due to certain orientations of the AO in space.
The bond formed by the overlapping of the AO along the line connecting the nuclei of the connecting atoms is called σ bond. Examples of the formation of σ-bonds are the overlaps of s-orbitals, s- and p-orbitals, p-orbitals, d-orbitals, as well as d- and s-orbitals, d- and p-orbitals, etc. Some of the examples of σ-bonds are given below.
It can be seen that in the case of σ bonds, the region of maximum electron density is located on the line connecting the nuclei of atoms.
A bond formed by overlapping AO on both sides of the line connecting the nuclei of atoms (lateral overlap) is called π-bond. A π bond can be formed by overlapping p-p, p-d, f-p, f-d, and f-f orbitals. Below are examples of the formation of π-bonds.
Since the degree of orbital overlap during the formation of π bonds is small compared to σ bonds, the energy of these bonds is much lower.
When a π-bond is superimposed on a σ-bond, a double bond is formed, for example, in the molecules of oxygen, ethylene, carbon dioxide:
O=O, C=C, O=C=O.
When two π-bonds are superimposed on one σ-bond, a triple bond arises, for example, in the molecules of nitrogen, acetylene, hydrocyanic acid :
The higher the bond multiplicity, the greater its energy and the shorter the bond length.
Some forms of compounds cannot be explained in terms of their formation from excited or unexcited atoms. So, in a methane molecule, all C-H bonds are equivalent, which contradicts the set of orbitals for excited and unexcited forms of the carbon atom. A consistent substantiation of this and other facts was found within the concept of AO hybridization.
Hybridization- this is a mixture of atomic orbitals of different energy and shape, leading to the formation of the same number of hybrid orbitals of the same energy and shape. The equivalence of hybrid orbitals causes not only the formation of bonds of equal energy, but also the same bond angles between the bonds formed by these orbitals. It should be emphasized that hybrid AOs are formed at the same atom having different orbitals, and the objects of hybridization are orbitals with close energy values.
In the case of methane, hybridization results from the mixing of one s- and three p-orbitals in the excited state of the carbon atom, the so-called sp 3 hybridization.
| 2p | 2p | ||||||||||||
| 2s | → | 2s | → | ||||||||||
| ↓ |
The formation of hybrid orbitals determines the energetic advantage of the chemical compounds formed through these orbitals. This is due to two factors.
First, hybrid orbitals are asymmetric, which causes a large degree of overlap in the formation of chemical bonds and their greater strength.
Secondly, the bond angles between hybrid orbitals are larger than non-hybrid ones, which causes a lower degree of repulsion between the bond electrons formed by these orbitals and makes molecular systems more stable.
With sp 3 hybridization, the longitudinal symmetry axes of the hybrid orbitals are at an angle of 109º28" with respect to each other - corresponding to their direction to the corners of the tetrahedron, the center of which is the atomic nucleus.
If the objects of hybridization are one s and two p-orbitals, then this type of hybridization is called sp 2 - hybridization, and the angles between the longitudinal axes of these orbitals are 120ºС and correspond to the minimum repulsion between valence electrons.
When mixing one s- and one p-orbital, sp-hybridization takes place. In this case, the bond angle between the hybrid orbitals is 180°C.
The spatial structure of molecules is determined by the number of atoms in the molecule, the hybridization of orbitals, and the number of unpaired electrons on them responsible for the formation of bonds.
A molecule formed by two atoms is linear. If there are two unpaired p-electrons on the outer shell of an atom, then when their AO orbitals overlap with other atoms, corner molecules are formed. These atoms include the atoms of p-elements of group VI (O, S, Se, Te), the electronic configuration of the outer shells of which is given below.
| ns | np | ||||
| ↓ | ↓ |
Two p-orbitals with unpaired electrons are located perpendicular to each other, so the angle in the H 2 S, H 2 Se and H 2 Te molecules is close to 90˚. Due to the repulsion of electrons, the valence angle between the bonds in the H 2 S molecule is somewhat higher than 90˚. For water molecules, the angle between bonds is much larger and equal to 105˚. Such a structure can be explained if we assume that this is the sp 2 hybridization of oxygen AO during the formation of water. In this case, two hybrid orbitals are overlapped by hydrogen s-orbitals. The repulsion of the valence electrons of the H-O bonds from the lone pairs of oxygen electrons causes a decrease in the bond angle from 120 ° to 105 °.
Evolution of the valence bond method
First approximate solution Schrödinger equations for one of the simplest molecules, the hydrogen molecule, was produced in 1927. V. Geytler and F. London. These authors first considered a system of two hydrogen atoms located at a great distance from each other. Under this condition, only the interaction of each electron with its “own” nucleus can be taken into account, and all other interactions (mutual repulsion of nuclei, attraction of each electron to a “foreign” nucleus, interaction between electrons) can be neglected. Then it becomes possible to express the dependence of the wave function of the system under consideration on the coordinates and thereby determine the density of the general electron cloud (electron density) at any point in space.
Further Geytler and London assumed that the dependence of the wave function on the coordinates found by them is also preserved when the hydrogen atoms approach each other. In this case, however, it is also necessary to take into account those interactions (between nuclei, between electrons, etc.) that could be neglected at a considerable distance of atoms from each other. These additional interactions are considered as some corrections ("perturbations") to the initial state of electrons in free hydrogen atoms.
As a result, equations were obtained that make it possible to find the dependence of the potential energy E system consisting of two hydrogen atoms, on the distance r between the nuclei of these atoms. It turned out that the results of the calculation depend on whether the spins of the interacting electrons are the same or opposite in sign. With the same direction of spins, the approach of atoms leads to a continuous increase in the energy of the system. In the latter case, the approach of the atoms requires an expenditure of energy, so that such a process turns out to be energetically unfavorable and no chemical bond arises between the atoms. With oppositely directed spins, the approach of atoms up to a certain distance r is accompanied by a decrease in the energy of the system. At r = r0 the system has the lowest potential energy, i.e. is in the most stable state; further approach of the atoms again leads to an increase in energy. But this also means that in the case of oppositely directed electron spins, a molecule is formed H 2- a stable system of two hydrogen atoms located at a certain distance from each other.
The formation of a chemical bond between hydrogen atoms is the result of the interpenetration ("overlapping") of electron clouds, which occurs when interacting atoms approach each other. As a result of such interpenetration, the density of the negative electric charge in the internuclear space increases. Positively charged atomic nuclei are attracted to the region of overlapping electron clouds. This attraction prevails over the mutual repulsion of like-charged electrons, so that a stable molecule is formed as a result.
Thus, the study made it possible to conclude that the chemical bond in the hydrogen molecule is carried out by the formation of a pair of electrons with oppositely directed spins belonging to both atoms. The theory of chemical bond developed on this basis and for more complex molecules was called valence bond method. The important point is that whenever a chemical bond is formed, the spins of a pair of electrons must be antiparallel. This is in line with Pauli principle and emphasizes that when a chemical bond is formed, electrons pass into a new quantum state.
The presence of paired electrons is an "indicator" of the presence of a chemical bond, but not the cause of its formation. The study of the cause of the formation of a chemical bond has so far shown that the energy of a system of two atoms decreases when the electrons are more likely to be in the internuclear space (as if “delayed” in this region). Such a delay leads to a decrease in their kinetic energy, as a result, the negative component of the total energy of the molecule prevails, the molecule becomes stable or, as they say, a chemical bond is formed.
The method of valence bonds gave a theoretical explanation of the most important properties of a covalent bond, made it possible to understand the structure of a large number of molecules. Although this method did not turn out to be universal and in some cases is not able to correctly describe the structure and properties of molecules, nevertheless it played a big role in the development of the quantum mechanical theory of chemical bonding and has not lost its significance to date in a qualitative understanding of the nature of chemical bonding.
Basic provisions of the method of valence bonds
The method of valence bonds describes the mechanism of occurrence of a covalent bond and is based on the following basic principles:
- A chemical bond between two atoms occurs through one or more shared electron pairs.
Both electrons of a common electron pair are held simultaneously by two nuclei, which is energetically more favorable than the presence of each electron in the field of “its own” nucleus.
Such a chemical bond is two-center.
For example, depict the formation of a molecule F2 with the help of quantum cells of the external energy level (the electronic formula of the atom F: 1s 2 2s 2 2p 5):
Paired electrons of the outer level of an atom must be separated (steered) to form chemical bonds with other atoms. The atom will go into a new valence state. The energy expended on such a process of excitation of an atom is compensated by the energy released during the formation of a chemical bond (it should be remembered that the possibilities of excitation of atoms are limited by the number of free orbitals in the corresponding energy sublevels).
- The covalent bond has the property of saturation, as a result of which the molecules have a well-defined composition.
For example, during the formation of a methane molecule CH 4 each of the four unpaired electrons of the excited carbon atom connected with the electron of the hydrogen atom, 4 covalent bonds were formed; more electron pairs in this case cannot be formed, molecules CH 5, CH 6 etc. does not exist.
(Note: the interaction of valence-saturated compounds with each other is possible with the formation of one or more additional donor-acceptor bonds according to a special mechanism).
- The covalent bond is directed in space, which determines the spatial structure of molecules (directivity property).
Depending on which electrons are bonding - s-, p-, d- or f- electrons, bond energies, bond lengths, as well as their direction in space are significantly different.
Electron clouds have a different shape, so their mutual overlap is carried out in several ways: σ- (sigma), π- (pi) and δ (delta)-bonds.
If the overlap of electron clouds occurs along the line connecting the nuclei - this is σ- connection; if clouds overlap outside this line, there are π- and δ -connections.
If one common electron pair has arisen between atoms (usually σ- connection), such a connection is called single, if two or more, then multiple: double, triple.
For example, the formation of a nitrogen molecule N 2 carried out by three common electron pairs. For each nitrogen atom, 3 unpaired bonds are involved in the formation of bonds R-electron directed in three-dimensional space at an angle of 90 0 to each other and oriented respectively along the axes x, y, z(these are the properties R- sublevel and R-orbitals dictated by the magnetic quantum number).
Two nitrogen atoms joining into a molecule N 2, can form one σ- communication (clouds oriented along the axis overlap) X) and two π- connections (clouds oriented along the axes of at and z).
Hybridization of atomic orbitals
The structure of molecules depends primarily on the type and properties of those orbitals that atoms provide for the formation of chemical bonds. But, in addition to this factor, the phenomenon of hybridization of orbitals affects the spatial structure of molecules.
hybridization called the formation of new orbitals of equal shape and energy from orbitals of different types. Mixed, hybrid orbitals in the diagrams are conventionally depicted:
sp hybridization
From one s-orbitals and one R-orbitals form two hybrid, mixed orbitals sp-type, directed with respect to each other by 180°.
For example: molecules have a linear shape Ven 2 and SnCl 2 With sp-hybridization of the atom of beryllium and tin, respectively.
sp 2 hybridization
From one s-orbitals and two R three orbitals are formed sp 2 hybrid orbitals located in the same plane at an angle of 120° to each other.
Mutual orientation of three sp 2-hybrid orbitals - trigonal. concept sp 2-hybridizations are used to describe planar trigonal molecules.
For example: aluminum fluoride molecule A1F3. The excitation of the aluminum atom is accompanied by steaming s2- electrons of the outer level on p-sublevel. Therefore, the electronic configuration of the outer level of an aluminum atom in an excited state is 3s 1 3p 2. The orbitals of the aluminum atom populated with electrons hybridize and orient themselves in the same plane at an angle of 120° to each other. Each of the three hybrid electron clouds sp 2-orbitals overlap with electron clouds p-orbitals of three fluorine atoms.
sp 3 hybridization
sp 3-hybridization occurs when combined s-orbital and three R-orbitals; four sp 3-hybrid orbitals, oriented no longer in one plane, but in the volume of the tetrahedron and directed from the center of the tetrahedron to its 4 vertices; the bond angle between two chemical bonds is 109°28".
For example: structure of the methane molecule CH 4. An excited carbon atom has four unpaired electrons: one s- and three R- electron. It would seem that the four chemical bonds formed by them with s- electrons of four hydrogen atoms must be unequal. However, it has been experimentally established that all 4 bonds in the molecule CH 4 completely identical in length and energy, and the angles between the bonds are 109 ° 28 ". Therefore, in the molecule CH 4 occurs sp 3-hybridization.
More complex cases of hybridization involving d-electrons, (for example, sp 3 d 2- hybridization).
The phenomenon of hybridization, i.e. mixing, equalization of electron density, energetically beneficial for the atom, since the hybrid orbitals have a deeper overlap and stronger chemical bonds are formed. A small amount of energy spent on excitation of the atom and hybridization of orbitals is more than offset by the energy released when chemical bonds occur. Bond angles are dictated by considerations of maximum symmetry and stability.
On hybrid orbitals, as well as on ordinary orbitals, not only one electron can be located, but also two. For example, four sp 3-hybrid orbitals of the oxygen atom O are such that two of them contain a pair of electrons, and two - one unpaired electron. From the modern standpoint, the structure of the water molecule is considered taking into account the hybridization of the atomic orbitals O and tetrahedral structure of the molecule H 2 O generally.
Valence by the exchange mechanism of the method
The ability of an atom to attach or replace a certain number of other atoms to form chemical bonds is called valence. According to the exchange mechanism of the valence bond method, each atom donates one unpaired electron to form a common electron pair (covalent bond). The quantitative measure of valence in the exchange mechanism of the method of valence bonds is the number of unpaired electrons in an atom in the ground or excited state of the atom. These are the unpaired electrons in the outer shells of s- and p- elements, outer and pre-outer shells d- elements, outer, pre-outer and pre-outer shells of f-elements.
When a chemical bond is formed, an atom can go into an excited state as a result of the separation of a pair (or pairs) of electrons and the transition of one electron (or several electrons equal to the number of disconnected pairs) into a free orbital of the same shell.
For example: The electronic configuration of calcium in the ground state is written as:
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2
In accordance with the exchange mechanism of the method of valence bonds, its valency is zero B=0. At the calcium atom in the fourth shell ( n=4) there are vacant R- orbitals. When an atom is excited, electrons are depaired and one of 4s- electrons goes into free 4p-orbital. The valency of calcium in the excited state is equal to two, i.e. when steaming, the valency increases by two units:
Unlike oxygen and fluorine, whose electron pairs cannot be separated, since there are no vacant orbitals in the second shell, the electron pairs of sulfur and chlorine atoms can be paired into vacant orbitals 3d-subshells, respectively, sulfur, in addition to the valency of the ground state 1 and 2, also has valencies 4 and 6 in the excited state, and chlorine, in addition to valence 1 in the ground state, has valences 3, 5 and 7 in the excited state.
Electronic configurations of atoms of some elements in the ground and excited states
| Element | Basic state | excited state | ||||||||
| Electronic configuration |
Orbital filling | Valence | Electronic configuration |
Orbital filling | Valence | |||||
| s | p | d | s | p | d | |||||
| Hydrogen | 1s 1 |  |
1 | |||||||
| Helium | 1s2 |  |
0 | |||||||
| Beryllium | 2s 2 |  |
 |
0 | 2s 1 2p 1 |  |
 |
2 | ||
| Carbon | 2s 2 2p 2 | |
 |
1,2 | 2s 1 2p 3 |  |
 |
1,2,4 | ||
| Oxygen | 2s 2 2p 4 | |
 |
1,2 | ||||||
| Fluorine | 2s 2 2p 5 | |
 |
1 | ||||||
| Sulfur | 3s 2 3p 4 |  |
 |
 |
1,2 | 3s 1 3p 3 3d 2 |  |
 |
 |
1,2,4,6 |
| Chlorine | 3s 2 3p 5 | |
 |
|
1 | 3s 1 3p 3 3d 3 | |
|
 |
1,3,5,7 |
Most atoms d- and f-elements on the outer shells in the ground state there are no unpaired electrons, therefore their valency in the ground state is zero, despite the fact that on the pre-outer d- and f subshells contain unpaired electrons. The latter cannot form electron pairs with the electrons of other atoms, since they are closed by the electrons of the outer shell. When an atom is excited, the paired electrons of the outer shell enter into a chemical bond and open the inner electron shells.
For example: the valency of iron in the ground state is zero:
In the excited state, disconnection occurs 4s-pairs of electrons:
The valency of iron in an excited state is determined not only 4s-, 4p-, but also 3d- unpaired electrons. However, a couple 3d-electrons cannot be separated because there are no vacant orbitals in the third shell, so the maximum valency of iron is six.
In osmium, when excited, not only external 6s-electrons, but also preexternal 5d-electrons, because in the fifth shell there is also 5f-subshell with free orbitals, so the maximum valency of osmium is eight:
1. covalent chemical bond form two electrons with opposite spins belonging to two atoms. For example, when two hydrogen atoms approach, their electron orbitals partially overlap and a common pair of electrons is formed
H× + × H = H : H
A covalent bond can also be formed by the donor-acceptor mechanism. Mechanism of Education covalent bonds due to an electron pair of one atom (donor) and another atom (acceptor), which provides a free orbital for this pair, is called donor-acceptor.
Let us consider as an example the mechanism of formation of the ammonium ion NH 4 + . In the NH 3 molecule, three shared electron pairs form three N-H bonds, the fourth pair of external electrons is unshared, it can give a bond with a hydrogen ion, resulting in an ammonium ion NH 4 +:
Thus, the NH 4 + ion has four covalent bonds, and all four N-H bonds are equivalent, that is, the electron density is evenly distributed between them.
2. When a covalent chemical bond is formed, the wave functions of electrons (electronic orbitals) overlap, and the bond will be the stronger, the greater this overlap.
3. The covalent chemical bond is located in the direction in which the possibility of overlapping the wave functions of the electrons forming the bond will be greatest.
4. Valence an atom in the normal (unexcited) state is determined by:
- the number of unpaired electrons participating in the formation of common electron pairs with electrons of other atoms;
- the presence of a donor ability (due to one unshared electron pair).
In an excited state, the valence of an atom is determined by:
- the number of unpaired electrons;
- the number of vacant orbitals capable of accepting donor electron pairs.
Thus, valency is expressed in small integers and has no sign. The measure of valence is the number chemical bonds by which this atom is connected to others.
The valence electrons primarily include the electrons of the outer levels, but for the elements of the secondary subgroups, they also include the electrons of the penultimate (anterior) levels.
Consider the electronic configuration of the boron atom:
where B* is the boron atom in the excited state.
The boron atom is monovalent in the ground state. The boron atom in an excited state has three unpaired electrons and can form compounds where it will be trivalent. The energy spent on the transition of an atom to an excited state within one energy level, as a rule, is compensated in excess by the energy released during the formation of additional bonds.
Due to the presence of a free orbital in the boron atom, boron is one of the strongest acceptors of unshared electron pairs. For example:
As a result, a complex ion is formed - having four covalent s-bonds.
Let's imagine the electron distribution scheme in the nitrogen atom:
Since nitrogen has three unpaired electrons, its valency is three. The transition of the nitrogen atom to an excited state is impossible, since the second energy layer does not contain a d-orbital. The nitrogen atom can donate an unshared electron pair of outer electrons to an atom that has a free orbital (acceptor). For example, in the ammonium ion, nitrogen is tetravalent (see point 1).
Schemes of overlapping atomic orbitals during the formation of bonds in H 2 O, NH 3 , CH 4 molecules
The water molecule consists of an oxygen atom and two hydrogen atoms. Two unpaired p-electrons of an oxygen atom occupy two orbitals, which are located at an angle of 90 o to each other. When a water molecule is formed, the orbital of each p-electron overlaps with the orbital s - an electron of the hydrogen atom (Fig. 7.1).
The angle between the bonds should be close to the angle between the p-electron clouds, i.e. to 90 about. It has been experimentally found that the angle between bonds in a water molecule is 104.5 o. This is due to the fact that the electrons are more drawn to the oxygen atom, since the O–H bond is a polar covalent bond. Thus, there is a repulsion of positive charges arising from hydrogen atoms, which leads to an increase in the angle between bonds.
The formation of an ammonia molecule involves three unpaired p-electrons of the nitrogen atom, the orbitals of which are located in three mutually perpendicular directions, and s-electrons of three hydrogen atoms (Fig. 7.2).
Three N–H bonds in the ammonia molecule should be at angles to each other close to 90°. It was experimentally found that the angle between the bonds in the ammonia molecule is 107.3 o, this is due to the same reason as in the case of the water molecule. In addition, we do not take into account the participation of 2s electrons in the formation of chemical bonds.
When a methane molecule is formed, the carbon atom goes into an excited state, that is, it has three unpaired p-electrons and one s-electron.
Arguing in the same way as in the previous cases, we can assume that the carbon atom will form three bonds directed at an angle of 90 o to each other and a bond directed arbitrarily, since it is formed by an s-electron, and the s-electron has spherical symmetry.
Since the p-orbitals are more elongated from the nucleus, the s-orbital, they overlap more with the orbitals of other atoms, which means that the bonds formed by the p-electrons should be stronger. But it is known from the experiment that all bonds in the methane molecule are equivalent and are directed to the vertices of the tetrahedron (the angle between the bonds is 109.5 o.
This phenomenon is explained by the concept of hybridization of wave functions introduced by Pauling and Slater. Hybridization of valence orbitals is their alignment in shape and energy. The concept of hybridization is used when electrons belonging to different types of orbitals participate in the formation of bonds in a molecule. The hybrid orbital is asymmetric and strongly elongated on one side of the nucleus.
Let us consider the electronic structure of the methane molecule, but from the point of view of the hybridization method. Four unpaired electrons of a carbon atom interact with each other during the formation of a chemical bond with the electrons of another atom, giving four new equivalent hybrid clouds. Such hybridization is called sp 3 hybridization. Four absolutely identical sp 3 -hybrid orbitals of the carbon atom are located at an angle of 109.5 o to each other and are directed to the vertices of the tetrahedron, in the center of which the carbon atom is located (Fig. 7.3).
The question arises - is it possible to explain the formation of a chemical bond between atoms in H 2 O and NH 3 molecules from the standpoint of hybridization of orbitals? The directionality of the bonds in these molecules can be explained using the concept of hybridization. This approach is even more accurate than the previous one. This is due to the fact that the hybrid orbital is strongly elongated in one direction from the nucleus, and the overlap of the hybrid orbitals with the electron orbitals of other atoms is stronger than the overlap of the s- and p-electron orbitals, which leads to the formation of a stronger bond, and, therefore, and more stable molecules.
Before proceeding to the consideration of the structure of H 2 O and NH 3 molecules using the hybridization model, we will compose an algorithm for determining the geometry of the molecule by this method:
- it is necessary to determine the presence of unshared electron pairs or unpaired electrons in the central atoms (by position in the periodic system);
- you should find the number of hybrid orbitals, which is equal to the sum of the number of s-bonds and the number of unshared electron pairs of the central atom;
- it is necessary to set the type of hybridization of orbitals:
Other types of hybridization of electron wave functions are also possible, for example, hybridization involving d-orbitals.
The formation of a chemical bond in H 2 O and NH 3 molecules can also be explained from the standpoint of sp 3 hybridization of the atomic orbitals of oxygen and nitrogen (Fig. 7.4).
While the carbon atom has all four hybrid orbitals occupied by bonding electron pairs (Fig. 7.3), the nitrogen atom has one of the four hybrid orbitals occupied by a lone electron pair (angle 107.3 o), and the oxygen atom has two orbitals (angle 104.5 o) (Fig. 7.4). This means that the repulsive action of unshared electron pairs affects the bond angles - when moving from methane molecules to ammonia and water molecules, the bond angle decreases.
Multiple bonds
-Connection - a chemical bond formed as a result of the overlap of electron orbitals along the line connecting the nuclei of atoms.
-Connection- a chemical bond formed as a result of the overlap of electron orbitals on both sides of the line connecting the nuclei of atoms.
The method of imposing valence schemes. Delocalized -bond
Let us consider the method of imposing valence schemes on the example of hydrogen azide HN 3 . In the HN 3 molecule, nitrogen atoms are bonded to each other, and one of them is bonded to hydrogen.
The central nitrogen atom can be trivalent due to unpaired electrons, but in both cases two unpaired electrons remain in the molecule, which makes the schemes unlikely.
Let us transfer one of the s-electrons of the central atom to the “upper” nitrogen atom:
An equiprobable scheme will be obtained if we move one of the s-electrons of the central atom to another nitrogen atom:
The concept of the method of molecular orbitals (MO)
The method of valence bonds (BC), despite its clarity, is not universal. It is satisfactorily applicable only to the description of the connection in the compounds of elements of periods I and II. The reason is that the chemical bond is presented as the result of the interaction of only a valence pair of electrons. The interaction of other electrons is not taken into account. With an increase in the number of electrons, this unaccounted for interaction becomes more and more pronounced, which leads to a discrepancy between the theoretical concepts of the VS method and experimental data.
There are already examples of such a discrepancy in periods I and II. The existence of the molecular hydrogen ion H^, consisting of two nuclei and a single electron, has been established. The bond, therefore, is formed by a single electron, and not by a pair, as the BC method suggests. The 0 2 oxygen molecule, as shown by magnetic measurements, has two unpaired electrons. However, according to the VS method, the 0 2 molecule must contain a double bond and there cannot be unpaired electrons.
These and many other "paradoxes" inexplicable from the standpoint of the VS method are easily explained by the MO method. A molecule in the MO method is considered as a single system of electrons and nuclei, in which each electron moves in the field of other electrons and nuclei. There are no atoms in a molecule; electrons do not belong to individual nuclei, but to the molecule as a whole. Quantum-mechanical regularities of the atom are transferred to the molecule: in the molecule there are MOs, molecular levels and sublevels characterized by molecular quantum numbers; the principle of least energy, the Pauli principle and Hund's rule are observed.
Molecular orbitals, unlike atomic orbitals (AO), are multicenter, and therefore have a more complex shape. By analogy with AO ( s, p, d, f) MO are denoted by the corresponding Greek letters a, ts, 5,
In the simplest approximation molecular orbitals can be represented as structures from the orbitals of the original atoms, obtained as linear combinations of atomic orbitals(MO L KAO). In this case, the initial AO should be close in energy and have the same symmetry with respect to the axis of the molecule. The number of formed MOs is equal to the number of initial AOs.
Consider the formation of a hydrogen molecule H 2 . From two Is-AOs of two H atoms, two MO molecules of H 2 are formed. The wave function of one of them is a linear combination of AOs with the same signs belonging to the atoms H (A) and H (B):
where Cj and C 2 are normalization factors (for a homonuclear molecule
“1= SUS The summation of wave functions with the same signs leads to an increase in the electron density between the nuclei and the formation of a chemical bond. Such an orbital is called binding(o l5). It has less energy than the original AO. The binding electron density is located along the line connecting the nuclei, so such an orbital is denoted by sr b.
where C 3 and C 4 are normalization factors (for a homonuclear molecule C 3 = C 4).
The wave function of the second MO is a linear combination of AO with opposite signs:
The electron density between the nuclei decreases and at a certain point it is equal to zero. Such an orbital has a higher energy than the original AO. It does not provide communication and is called loosening(a ls *). The forms of the st, 5 *- and o, 5 -orbitals are shown in Fig. 3.19.
Rice. 3.19.
According to the principle of least energy, the bonding st^-orbital is filled first. According to the Pauli exclusion, it can have two electrons with opposite spins. Therefore, both electrons in the H 2 molecule are located in the bonding orbital, while the loosening orbital remains vacant. The energy diagram of the H 2 molecule is shown in fig. 3.20, a. MOs are presented in the center of the diagram, and AOs of the initial atoms are shown along the edges.
The difference between the energy of the initial AO and the binding MO (binding energy Eb) is approximately equal to the difference between the energy of the loosening MO and the initial AO (loosening energy E). The question of the stability of a molecule is reduced to the energy balance of all bonding and loosening electrons: the number of electrons in bonding orbitals must be greater than in loosening ones.
Rice. 3.20. Energy diagrams of molecules and ions: a -H 2; b -H 2 +; c -H 2
The formation of a single bond corresponds to an excess pair of electrons in the bonding orbitals. The bond multiplicity (CC) in the MO method is defined as the number of excess pairs of electrons in bonding orbitals (lb) compared to loosening orbitals
The electronic formula of the H 2 molecule is written as follows: 2H->H 2 [(a lj) 2].
The explanation of the formation of the molecular ion H 2 using the MO method does not cause difficulties: the only electron goes to the a b orbital (see Fig. 3.20, b). Multiplicity of communication Y 2 . Electronic formula: H + + H -» HJ [(1s)‘].
The existence of another molecular hydrogen ion H^ is also possible: H“ + H^H2[(a ls) 2 (a* li) 1 ]. The third electron enters the loosening orbital (see Fig. 3.20, in). Communication multiplicity
An isoelectronic H 2 ion should be an He 2 ion with KS = 1 / 2: Not + + He -> Not 2 + [(a b R (a 1 *) 1]. Indeed, such an ion exists.
It is known that helium does not form diatomic molecules. This is in accordance with the representations of the MO method. In a hypothetical He 2 molecule, the number of electrons on the bonding and loosening MOs would be the same: 2He -> He 2 [(o b) 2 (st*) 2]. There is no gain in energy during its formation. The formal multiplicity of the connection is equal to zero.
A similar approach is used to analyze diatomic homonuclear molecules and ions of other s-elements. In period II, for example, there is a molecule Li 2 [(aj.) 2 (a 1 *) 2 (a 2i) 2], but there is no
For atoms of elements of the II period, valence, except 2s-, are 2p x ~, 2p y - and 2/> z-orbitals. They have the same energy, i.e. /orbitals are threefold degenerate. When chemical bonds are formed, the degeneracy is partially lifted. If in a molecule of type A 2 we choose for the x axis a line passing through the nuclei of atoms, then the combination of 2/^-orbitals leads to the formation of two a-orbitals: the bonding and loosening with 2r? Combination 2/? r-orbitals leads to the formation of bonding and antibonding orbitals of the l-type: l 2p and l* 2p. Similar MO (p 2p and n * 2p) are formed by the combination of 2/urbitals. They differ from l 2r and l 2r only position in space - rotation by 90 ° around the x-axis. The forms of binding and loosening MOs formed from 2/nAO are shown in Figs. 3.21.
Rice. 3.21. The combination of 2p orbitals in a molecule of type a - su 2p1 b - k 2p
Orbitals p 2p and p 2p have the same energy, that is, they are doubly degenerate. Since the a-bond is stronger than the l-bond, the bonding a-orbitals have less energy than the bonding l-orbitals (Fig. 3.22, b). However, this is true only for diatomic homonuclear molecules of the end of the period. For them, the increase in energy and, accordingly, the filling of the MO follow in this order:
Electrons on a 2s- and orbitals repel each other. With energetic closeness is- and 2/orbitals characteristic of the beginning of the period, this leads to the fact that p 2p- and p 2 Orbitals become energetically more favorable than from 2r-orbital. The order of filling in the MO changes accordingly:
This is observed for diatomic homonuclear molecules of the beginning of the period up to N 2 (Fig. 3.22, a).
Rice. 3.22. Energy diagrams of compounds of type A 2: a - N 2; b - 0 2 (1s- and 2s-A0 and g 1s - ct * s - cr 2s -, cj 2s - MO not shown)
Consider the formation of the N 2 molecule. Six 2p electrons of two nitrogen atoms go to three bonding orbitals (Fig. 3.22, a). The electronic formula of the molecule is:
There are no unpaired electrons in the molecule. multiplicity
connection COP =- = 3.
The last electrons are located in bonding orbitals. The removal of an electron from them leads to a decrease in the multiplicity and binding energy. Therefore, the binding energy in the N 2 ion is less than in the N 2 molecule.
In the 0 2 molecule (Fig. 3.22, b) the last two electrons fill the antibonding orbitals:
and are arranged according to Hund's rule, one in two orbitals. Therefore, there are two unpaired electrons in the 0 2 molecule. multiplicity
bonds KS = --- = 2. The detachment of an electron from a molecule occurs with
loosening orbitals, the multiplicity and binding energy increase in this case. Therefore, the OJ ion is stronger than the molecule 0 2
In diatomic heteronuclear molecules, AO overlap is possible only if their energies are close. Nevertheless, the energies of the interacting AOs are different, and their relative contribution to the MO formation is also different, which is reflected by the normalization factors: Cj f From 2 in equation (3.11) and C 3 f From 4 in equation (4.12). The binding MO is closer in energy to the AO of a more electronegative atom, and the loosening MO is closer to the AO of a less electronegative atom (Fig. 3.23). The polarity of the bond in the molecule is determined by the difference in the energy of the initial AO, and the value b is proportional to the polar component of the bond, and the value a- covalent.
Rice. 3.23.
Consider, as an example, the molecule of carbon monoxide CO. Using the HS method, the bond can be represented as a triple bond, which consists of two bonds formed by the exchange mechanism and one by the donor-acceptor mechanism:
The CO molecule is isoelectronic to the nitrogen molecule N 2, therefore, the form of writing the electronic formulas of these molecules according to the MO method is the same:
Rice. 3.24.
The energy diagrams of CO and N 2 are also similar to each other (Fig. 3.24 and Fig. 3.22, a).
The difference is that the orbitals of oxygen, a more electronegative atom, make a greater contribution to the bonding orbitals, and the orbitals of carbon, a less electronegative atom, contribute to the loosening orbitals.
The multiplicity of communication, as well as according to the VS method, is equal to three. But, using MO, there is no need to resort to explaining the bonds with the help of two mechanisms: exchange and donor-acceptor.
The presence of two methods for describing the chemical bond, VS and MO, indicates the imperfection of modern ideas about the chemical bond. Where both of these methods are applicable, they give similar results. However, there are specific questions that only one of them can answer. Nevertheless, it is fair to believe that these quantum mechanical methods do not exclude, but complement each other.
