Lecture for teachers
A chemical bond (hereinafter referred to as a bond) can be defined as the interaction of two or more atoms, as a result of which a chemically stable polyatomic microsystem (molecule, crystal, complex, etc.) is formed.
The doctrine of bonding occupies a central place in modern chemistry, since chemistry as such begins where the isolated atom ends and the molecule begins. In essence, all properties of substances are determined by the characteristics of the bonds in them. The main difference between a chemical bond and other types of interactions between atoms is that its formation is determined by a change in the state of the electrons in the molecule compared to the original atoms.
Communication theory should provide answers to a number of questions. Why are molecules formed? Why do some atoms interact while others do not? Why do atoms combine in certain ratios? Why are atoms arranged in a certain way in space? And finally, it is necessary to calculate the bond energy, its length and other quantitative characteristics. The correspondence of theoretical concepts to experimental data should be considered as a criterion for the truth of the theory.
There are two main methods for describing communication that allow you to answer the questions posed. These are the methods of valence bonds (BC) and molecular orbitals (MO). The first one is more visual and simple. The second is more strict and universal. Due to greater clarity, the focus here will be on the BC method.
Quantum mechanics allows us to describe the connection based on the most general laws. Although there are five types of bonds (covalent, ionic, metallic, hydrogen and intermolecular interaction bonds), the bond is uniform in nature, and the differences between its types are relative. The essence of communication is in Coulomb interaction, in the unity of opposites - attraction and repulsion. The division of communication into types and the difference in methods of describing it indicate not the diversity of communication, but rather the lack of knowledge about it at the present stage of development of science.
This lecture will cover topics such as chemical bond energy, the quantum mechanical model of covalent bonds, exchange and donor-acceptor mechanisms of covalent bond formation, atomic excitation, bond multiplicity, hybridization of atomic orbitals, electronegativity of elements and covalent bond polarity , concept of the molecular orbital method, chemical bonding in crystals.
Chemical bond energy
According to the principle of least energy, the internal energy of a molecule should decrease compared to the sum of the internal energies of the atoms that form it. The internal energy of a molecule includes the sum of the interaction energies of each electron with each nucleus, each electron with each other electron, and each nucleus with each other nucleus. Attraction must prevail over repulsion.
The most important characteristic of a bond is energy, which determines its strength. A measure of the strength of a bond can be both the amount of energy spent on breaking it (bond dissociation energy) and the value that, when summed over all bonds, gives the energy of formation of a molecule from elementary atoms. The energy of breaking a bond is always positive. The energy of bond formation is the same in magnitude, but has a negative sign.
For a diatomic molecule, the binding energy is numerically equal to the energy of dissociation of the molecule into atoms and the energy of formation of the molecule from atoms. For example, the binding energy in a HBr molecule is equal to the amount of energy released in the process H + Br = HBr. It is obvious that the binding energy of HBr is greater than the amount of energy released during the formation of HBr from gaseous molecular hydrogen and liquid bromine:
1/2Н 2 (g.) + 1/2Вr 2 (l.) = НBr (g.),
on the energy value of evaporation of 1/2 mol Br 2 and on the energy value of decomposition of 1/2 mol H 2 and 1/2 mol Br 2 into free atoms.
Quantum mechanical model of covalent bonding using the valence bond method using the example of a hydrogen molecule
In 1927, the Schrödinger equation was solved for the hydrogen molecule by German physicists W. Heitler and F. London. This was the first successful attempt to apply quantum mechanics to solve communication problems. Their work laid the foundations for the method of valence bonds, or valence schemes (VS).
The calculation results can be presented graphically in the form of dependences of the interaction forces between atoms (Fig. 1, a) and the energy of the system (Fig. 1, b) on the distance between the nuclei of hydrogen atoms. We will place the nucleus of one of the hydrogen atoms at the origin of coordinates, and the nucleus of the second will be brought closer to the nucleus of the first hydrogen atom along the abscissa axis. If the electron spins are antiparallel, the attractive forces (see Fig. 1, a, curve I) and repulsive forces (curve II) will increase. The resultant of these forces is represented by curve III. At first, the forces of attraction predominate, then the forces of repulsion. When the distance between the nuclei becomes equal to r 0 = 0.074 nm, the attractive force is balanced by the repulsive force. The balance of forces corresponds to the minimum energy of the system (see Fig. 1, b, curve IV) and, therefore, the most stable state. The depth of the “potential well” represents the bond energy E 0 H–H in the H 2 molecule at absolute zero. It is 458 kJ/mol. However, at real temperatures, bond breaking requires slightly less energy E H–H, which at 298 K (25 ° C) is equal to 435 kJ/mol. The difference between these energies in the H2 molecule is the energy of vibrations of hydrogen atoms (E coll = E 0 H–H – E H–H = 458 – 435 = 23 kJ/mol).
Rice. 1. Dependence of the forces of interaction between atoms (a) and the energy of the system (b)
on the distance between the nuclei of atoms in the H 2 molecule
When two hydrogen atoms containing electrons with parallel spins approach each other, the energy of the system constantly increases (see Fig. 1, b, curve V) and a bond is not formed.
Thus, the quantum mechanical calculation provided a quantitative explanation of the connection. If a pair of electrons has opposite spins, the electrons move in the field of both nuclei. Between the nuclei there appears an area with a high density of electron cloud - an excess negative charge that attracts positively charged nuclei. From the quantum mechanical calculation follow the provisions that are the basis of the VS method:
1. The reason for the connection is the electrostatic interaction of nuclei and electrons.
2. The bond is formed by an electron pair with antiparallel spins.
3. Bond saturation is due to the formation of electron pairs.
4. The strength of the connection is proportional to the degree of overlap of the electron clouds.
5. The directionality of the connection is due to the overlap of electron clouds in the region of maximum electron density.
Exchange mechanism of covalent bond formation using the BC method. Directionality and saturation of covalent bonds
One of the most important concepts of the BC method is valence. The numerical value of valence in the BC method is determined by the number of covalent bonds that an atom forms with other atoms.
The mechanism considered for the H2 molecule for the formation of a bond by a pair of electrons with antiparallel spins, which belonged to different atoms before the formation of the bond, is called exchange. If only the exchange mechanism is taken into account, the valence of an atom is determined by the number of its unpaired electrons.
For molecules more complex than H2, the principles of calculation remain unchanged. The formation of a bond is caused by the interaction of a pair of electrons with opposite spins, but with wave functions of the same sign, which are summed. The result of this is an increase in electron density in the region of overlapping electron clouds and contraction of nuclei. Let's look at examples.
In the fluorine molecule, the F2 bond is formed by 2p orbitals of fluorine atoms:
The highest density of the electron cloud is near the 2p orbital in the direction of the symmetry axis. If the unpaired electrons of fluorine atoms are in 2p x orbitals, the bond occurs in the direction of the x axis (Fig. 2). The 2p y and 2p z orbitals contain lone pairs of electrons that are not involved in the formation of bonds (shaded in Fig. 2). In what follows we will not depict such orbitals.
Rice. 2. Formation of the F 2 molecule
In the hydrogen fluoride molecule HF, the bond is formed by the 1s orbital of the hydrogen atom and the 2p x orbital of the fluorine atom:
The direction of the bond in this molecule is determined by the orientation of the 2px orbital of the fluorine atom (Fig. 3). The overlap occurs in the direction of the x axis of symmetry. Any other overlap option is energetically less favorable.
Rice. 3. Formation of the HF molecule
More complex d- and f-orbitals are also characterized by the directions of maximum electron density along their symmetry axes.
Thus, directionality is one of the main properties of a covalent bond.
The direction of the bond is well illustrated by the example of the hydrogen sulfide molecule H 2 S:
Since the symmetry axes of the valence 3p orbitals of the sulfur atom are mutually perpendicular, it should be expected that the H 2 S molecule should have a corner structure with an angle between the S–H bonds of 90° (Fig. 4). Indeed, the angle is close to the calculated one and is equal to 92°.
Rice. 4. Formation of the H 2 S molecule
Obviously, the number of covalent bonds cannot exceed the number of electron pairs forming the bonds. However, saturation as a property of a covalent bond also means that if an atom has a certain number of unpaired electrons, then all of them must participate in the formation of covalent bonds.
This property is explained by the principle of least energy. With each additional bond formed, additional energy is released. Therefore, all valence possibilities are fully realized.
Indeed, the stable molecule is H 2 S, not HS, where there is an unrealized bond (the unpaired electron is designated by a dot). Particles containing unpaired electrons are called free radicals. They are extremely reactive and react to form compounds containing saturated bonds.
Excitation of atoms
Let's consider the valence possibilities according to the exchange mechanism of some elements of the 2nd and 3rd periods of the periodic table.
The beryllium atom at the outer quantum level contains two paired 2s electrons. There are no unpaired electrons, so beryllium must have zero valence. However, in compounds it is divalent. This can be explained by the excitation of the atom, which consists in the transition of one of the two 2s electrons to the 2p sublevel:
In this case, excitation energy E* is expended, corresponding to the difference between the energies of the 2p and 2s sublevels.
When a boron atom is excited, its valence increases from 1 to 3:
and the carbon atom has from 2 to 4:
At first glance, it may seem that excitation contradicts the principle of least energy. However, as a result of excitation, new, additional connections arise, due to which energy is released. If this additional energy released is greater than that expended on excitation, the principle of least energy is ultimately satisfied. For example, in a CH4 methane molecule, the average C–H bond energy is 413 kJ/mol. The energy expended for excitation is E* = 402 kJ/mol. The energy gain due to the formation of two additional bonds will be:
D E = E additional light – E* = 2,413 – 402 = 424 kJ/mol.
If the principle of least energy is not respected, i.e. E add.st.< Е*, то возбуждение не происходит. Так, энергетически невыгодным оказывается возбуждение атомов элементов 2-го периода за счет перехода электронов со второго на третий квантовый уровень.
For example, oxygen is only divalent for this reason. However, the electronic analogue of oxygen - sulfur - has greater valence capabilities, since the third quantum level has a 3d sublevel, and the energy difference between the 3s, 3p and 3d sublevels is incomparably smaller than between the second and third quantum levels of the oxygen atom:
For the same reason, the elements of the 3rd period - phosphorus and chlorine - exhibit variable valence, in contrast to their electronic analogues in the 2nd period - nitrogen and fluorine. Excitation to the corresponding sublevel can explain the formation of chemical compounds of group VIIIa elements of the 3rd and subsequent periods. No chemical compounds were found in helium and neon (1st and 2nd periods), which have a completed external quantum level, and they are the only truly inert gases.
Donor-acceptor mechanism of covalent bond formation
A pair of electrons with antiparallel spins forming a bond can be obtained not only by the exchange mechanism, which involves the participation of electrons from both atoms, but also by another mechanism, called donor-acceptor: one atom (donor) provides a lone pair of electrons for the formation of the bond, and the other (acceptor) – vacant quantum cell:
The result for both mechanisms is the same. Often bond formation can be explained by both mechanisms. For example, an HF molecule can be obtained not only in the gas phase from atoms according to the exchange mechanism, as shown above (see Fig. 3), but also in an aqueous solution from H + and F – ions according to the donor-acceptor mechanism:
There is no doubt that molecules produced by different mechanisms are indistinguishable; connections are completely equivalent. Therefore, it is more correct not to distinguish the donor-acceptor interaction as a special type of bond, but to consider it only a special mechanism for the formation of a covalent bond.
When they want to emphasize the mechanism of bond formation precisely according to the donor-acceptor mechanism, it is denoted in structural formulas by an arrow from the donor to the acceptor (D® A). In other cases, such a connection is not isolated and is indicated by a dash, as in the exchange mechanism: D–A.
Bonds in the ammonium ion formed by the reaction: NH 3 + H + = NH 4 +,
are expressed by the following scheme:
The structural formula of NH 4 + can be represented as
.
The second form of notation is preferable, since it reflects the experimentally established equivalence of all four connections.
The formation of a chemical bond by the donor-acceptor mechanism expands the valence capabilities of atoms: valence is determined not only by the number of unpaired electrons, but also by the number of lone electron pairs and vacant quantum cells involved in the formation of bonds. So, in the example given, the valency of nitrogen is four.
The donor-acceptor mechanism is successfully used to describe the bonding in complex compounds using the BC method.
Multiplicity of communication. s- and p -Connections
The connection between two atoms can be carried out not only by one, but also by several electron pairs. It is the number of these electron pairs that determines the multiplicity in the BC method - one of the properties of a covalent bond. For example, in the ethane molecule C 2 H 6 the bond between the carbon atoms is single (single), in the ethylene molecule C 2 H 4 it is double, and in the acetylene molecule C 2 H 2 it is triple. Some characteristics of these molecules are given in table. 1.
Table 1
Changes in bond parameters between C atoms depending on its multiplicity
As the bond multiplicity increases, as one would expect, its length decreases. The bond multiplicity increases discretely, that is, by an integer number of times, therefore, if all bonds were the same, the energy would also increase by a corresponding number of times. However, as can be seen from table. 1, the binding energy increases less rapidly than the multiplicity. Consequently, the connections are unequal. This can be explained by differences in the geometric ways in which the orbitals overlap. Let's look at these differences.
A bond formed by overlapping electron clouds along an axis passing through the nuclei of atoms is called s-bond.
If the s-orbital is involved in the bond, then only s - connection (Fig. 5, a, b, c). This is where it got its name, since the Greek letter s is synonymous with the Latin s.
When p-orbitals (Fig. 5, b, d, e) and d-orbitals (Fig. 5, c, e, f) participate in the formation of a bond, s-type overlap occurs in the direction of the highest density of electron clouds, which is the most energetically favorable. Therefore, when forming a connection, this method is always implemented first. Therefore, if the connection is single, then this is mandatory s - connection, if multiple, then one of the connections is certainly s-connection.
Rice. 5. Examples of s-bonds
However, from geometric considerations it is clear that between two atoms there can be only one s -connection. In multiple bonds, the second and third bonds must be formed by a different geometric method of overlapping electron clouds.
The bond formed by the overlap of electron clouds on either side of an axis passing through the nuclei of atoms is called p-bond. Examples p - connections are shown in Fig. 6. Such overlap is energetically less favorable than s -type. It is carried out by the peripheral parts of electron clouds with lower electron density. Increasing the multiplicity of the connection means the formation p -bonds that have lower energy compared to s - communication. This is the reason for the nonlinear increase in binding energy in comparison with the increase in multiplicity.
Rice. 6. Examples of p-bonds
Let's consider the formation of bonds in the N 2 molecule. As is known, molecular nitrogen is chemically very inert. The reason for this is the formation of a very strong NєN triple bond:
A diagram of the overlap of electron clouds is shown in Fig. 7. One of the bonds (2рх–2рх) is formed according to the s-type. The other two (2рz–2рz, 2рy–2рy) are p-type. In order not to clutter the figure, the image of the overlap of 2py clouds is shown separately (Fig. 7, b). To get the general picture, Fig. 7, a and 7, b should be combined.
At first glance it may seem that s -bond, limiting the approach of atoms, does not allow the orbitals to overlap p -type. However, the image of the orbital includes only a certain fraction (90%) of the electron cloud. The overlap occurs with a peripheral region located outside such an image. If we imagine orbitals that include a large fraction of the electron cloud (for example, 95%), then their overlap becomes obvious (see dashed lines in Fig. 7, a).
Rice. 7. Formation of the N 2 molecule
To be continued
V.I. Elfimov,
professor of Moscow
State Open University
Link length - internuclear distance. The shorter this distance, the stronger the chemical bond. The length of a bond depends on the radii of the atoms forming it: the smaller the atoms, the shorter the bond between them. For example, the H-O bond length is shorter than the H-N bond length (due to less oxygen atom exchange).
An ionic bond is an extreme case of a polar covalent bond.
Metal connection.
The prerequisite for the formation of this type of connection is:
1) the presence of a relatively small number of electrons at the outer levels of atoms;
2) the presence of empty (vacant orbitals) on the outer levels of metal atoms
3) relatively low ionization energy.
Let's consider the formation of a metal bond using sodium as an example. The valence electron of sodium, which is located on the 3s sublevel, can relatively easily move through the empty orbitals of the outer layer: along 3p and 3d. When atoms come closer together as a result of the formation of a crystal lattice, the valence orbitals of neighboring atoms overlap, due to which electrons move freely from one orbital to another, establishing a bond between ALL atoms of the metal crystal.
At the nodes of the crystal lattice there are positively charged metal ions and atoms, and between them there are electrons that can move freely throughout the crystal lattice. These electrons become common to all atoms and ions of the metal and are called "electron gas". The bond between all positively charged metal ions and free electrons in the metal crystal lattice is called metal bond.
The presence of a metallic bond determines the physical properties of metals and alloys: hardness, electrical conductivity, thermal conductivity, malleability, ductility, metallic luster. Free electrons can carry heat and electricity, so they are the reason for the main physical properties that distinguish metals from non-metals - high electrical and thermal conductivity.
Hydrogen bond.
Hydrogen bond occurs between molecules that contain hydrogen and atoms with high EO (oxygen, fluorine, nitrogen). Covalent bonds H-O, H-F, H-N are highly polar, due to which an excess positive charge accumulates on the hydrogen atom, and an excess negative charge on the opposite poles. Between oppositely charged poles, forces of electrostatic attraction arise - hydrogen bonds.
Hydrogen bonds can be either intermolecular or intramolecular. The energy of a hydrogen bond is approximately ten times less than the energy of a conventional covalent bond, but nevertheless, hydrogen bonds play an important role in many physicochemical and biological processes. In particular, DNA molecules are double helices in which two chains of nucleotides are linked by hydrogen bonds. Intermolecular hydrogen bonds between water and hydrogen fluoride molecules can be depicted (by dots) as follows:
Substances with hydrogen bonds have molecular crystal lattices. The presence of a hydrogen bond leads to the formation of molecular associates and, as a consequence, to an increase in the melting and boiling points.
In addition to the listed main types of chemical bonds, there are also universal forces of interaction between any molecules that do not lead to the breaking or formation of new chemical bonds. These interactions are called van der Waals forces. They determine the attraction of molecules of a given substance (or various substances) to each other in liquid and solid states of aggregation.
Different types of chemical bonds determine the existence of different types of crystal lattices (table).
Substances consisting of molecules have molecular structure. These substances include all gases, liquids, as well as solids with a molecular crystal lattice, such as iodine. Solids with an atomic, ionic or metal lattice have non-molecular structure, they have no molecules.
Table
| Feature of the crystal lattice | Lattice type | |||
| Molecular | Ionic | Nuclear | Metal | |
| Particles at lattice nodes | Molecules | Cations and anions | Atoms | Metal cations and atoms |
| The nature of the connection between particles | Intermolecular interaction forces (including hydrogen bonds) | Ionic bonds | Covalent bonds | Metal connection |
| Bond strength | Weak | Durable | Very durable | Various strengths |
| Distinctive physical properties of substances | Low-melting or sublimating, low hardness, many soluble in water | Refractory, hard, brittle, many soluble in water. Solutions and melts conduct electric current | Very refractory, very hard, practically insoluble in water | High electrical and thermal conductivity, metallic luster, ductility. |
| Examples of substances | Simple substances - non-metals (in solid state): Cl 2, F 2, Br 2, O 2, O 3, P 4, sulfur, iodine (except silicon, diamond, graphite); complex substances consisting of non-metal atoms (except ammonium salts): water, dry ice, acids, non-metal halides: PCl 3, SiF 4, CBr 4, SF 6, organic substances: hydrocarbons, alcohols, phenols, aldehydes, etc. | Salts: sodium chloride, barium nitrate, etc.; alkalis: potassium hydroxide, calcium hydroxide, ammonium salts: NH 4 Cl, NH 4 NO 3, etc., metal oxides, nitrides, hydrides, etc. (compounds of metals with non-metals) | Diamond, graphite, silicon, boron, germanium, silicon oxide (IV) - silica, SiC (carborundum), black phosphorus (P). | Copper, potassium, zinc, iron and other metals |
| Comparison of substances by melting and boiling points. | ||||
| Due to weak intermolecular interaction forces, such substances have the lowest melting and boiling points. Moreover, the greater the molecular weight of the substance, the higher the t 0 pl. it has. Exceptions are substances whose molecules can form hydrogen bonds. For example, HF has a higher t0 pl. than HCl. | Substances have high t 0 pl., but lower than substances with an atomic lattice. The higher the charges of the ions that are located in the lattice sites and the shorter the distance between them, the higher the melting point of the substance. For example, t 0 pl. CaF 2 is higher than t 0 pl. KF. | They have the highest t 0 pl. The stronger the bond between the atoms in the lattice, the higher the t 0 pl. has substance. For example, Si has a lower t0 pl. than C. | Metals have different t0 pl.: from -37 0 C for mercury to 3360 0 C for tungsten. |
Chemical bond
All interactions leading to the combination of chemical particles (atoms, molecules, ions, etc.) into substances are divided into chemical bonds and intermolecular bonds (intermolecular interactions).
Chemical bonds- bonds directly between atoms. There are ionic, covalent and metallic bonds.
Intermolecular bonds- connections between molecules. These are hydrogen bonds, ion-dipole bonds (due to the formation of this bond, for example, the formation of a hydration shell of ions occurs), dipole-dipole (due to the formation of this bond, molecules of polar substances are combined, for example, in liquid acetone), etc.
Ionic bond- a chemical bond formed due to the electrostatic attraction of oppositely charged ions. In binary compounds (compounds of two elements), it is formed when the sizes of the bonded atoms are very different from each other: some atoms are large, others are small - that is, some atoms easily give up electrons, while others tend to accept them (usually these are atoms of the elements that form typical metals and atoms of elements forming typical nonmetals); the electronegativity of such atoms is also very different.
Ionic bonding is non-directional and non-saturable.
Covalent bond- a chemical bond that occurs due to the formation of a common pair of electrons. A covalent bond is formed between small atoms with the same or similar radii. A necessary condition is the presence of unpaired electrons in both bonded atoms (exchange mechanism) or a lone pair in one atom and a free orbital in the other (donor-acceptor mechanism):
| A) | H· + ·H H:H | H-H | H 2 | (one shared pair of electrons; H is monovalent); |
| b) | NN | N 2 | (three shared pairs of electrons; N is trivalent); | |
| V) | H-F | HF | (one shared pair of electrons; H and F are monovalent); | |
| G) | NH4+ | (four shared pairs of electrons; N is tetravalent) |
- Based on the number of shared electron pairs, covalent bonds are divided into
- simple (single)- one pair of electrons,
- double- two pairs of electrons,
- triples- three pairs of electrons.
Double and triple bonds are called multiple bonds.
According to the distribution of electron density between the bonded atoms, a covalent bond is divided into non-polar And polar. A non-polar bond is formed between identical atoms, a polar one - between different ones.
Electronegativity- a measure of the ability of an atom in a substance to attract common electron pairs.
The electron pairs of polar bonds are shifted towards more electronegative elements. The displacement of electron pairs itself is called bond polarization. The partial (excess) charges formed during polarization are designated + and -, for example: .
Based on the nature of the overlap of electron clouds ("orbitals"), a covalent bond is divided into -bond and -bond.
-A bond is formed due to the direct overlap of electron clouds (along the straight line connecting the atomic nuclei), -a bond is formed due to lateral overlap (on both sides of the plane in which the atomic nuclei lie).
A covalent bond is directional and saturable, as well as polarizable.
The hybridization model is used to explain and predict the mutual direction of covalent bonds.
Hybridization of atomic orbitals and electron clouds- the supposed alignment of atomic orbitals in energy, and electron clouds in shape when an atom forms covalent bonds.
The three most common types of hybridization are: sp-, sp 2 and sp 3 -hybridization. For example:
sp-hybridization - in molecules C 2 H 2, BeH 2, CO 2 (linear structure);
sp 2-hybridization - in molecules C 2 H 4, C 6 H 6, BF 3 (flat triangular shape);
sp 3-hybridization - in molecules CCl 4, SiH 4, CH 4 (tetrahedral form); NH 3 (pyramidal shape); H 2 O (angular shape).
Metal connection- a chemical bond formed by sharing the valence electrons of all bonded atoms of a metal crystal. As a result, a single electron cloud of the crystal is formed, which easily moves under the influence of electrical voltage - hence the high electrical conductivity of metals.
A metallic bond is formed when the atoms being bonded are large and therefore tend to give up electrons. Simple substances with a metallic bond are metals (Na, Ba, Al, Cu, Au, etc.), complex substances are intermetallic compounds (AlCr 2, Ca 2 Cu, Cu 5 Zn 8, etc.).
The metal bond does not have directionality or saturation. It is also preserved in metal melts.
Hydrogen bond- an intermolecular bond formed due to the partial acceptance of a pair of electrons from a highly electronegative atom by a hydrogen atom with a large positive partial charge. It is formed in cases where one molecule contains an atom with a lone pair of electrons and high electronegativity (F, O, N), and the other contains a hydrogen atom bound by a highly polar bond to one of such atoms. Examples of intermolecular hydrogen bonds:
H—O—H OH 2 , H—O—H NH 3 , H—O—H F—H, H—F H—F.
Intramolecular hydrogen bonds exist in the molecules of polypeptides, nucleic acids, proteins, etc.
A measure of the strength of any bond is the bond energy.
Communication energy- the energy required to break a given chemical bond in 1 mole of a substance. The unit of measurement is 1 kJ/mol.
The energies of ionic and covalent bonds are of the same order, the energy of hydrogen bonds is an order of magnitude less.
The energy of a covalent bond depends on the size of the bonded atoms (bond length) and on the multiplicity of the bond. The smaller the atoms and the greater the bond multiplicity, the greater its energy.
The ionic bond energy depends on the size of the ions and their charges. The smaller the ions and the greater their charge, the greater the binding energy.
Structure of matter
According to the type of structure, all substances are divided into molecular And non-molecular. Among organic substances, molecular substances predominate, among inorganic substances, non-molecular substances predominate.
Based on the type of chemical bond, substances are divided into substances with covalent bonds, substances with ionic bonds (ionic substances) and substances with metallic bonds (metals).
Substances with covalent bonds can be molecular or non-molecular. This significantly affects their physical properties.
Molecular substances consist of molecules connected to each other by weak intermolecular bonds, these include: H 2, O 2, N 2, Cl 2, Br 2, S 8, P 4 and other simple substances; CO 2, SO 2, N 2 O 5, H 2 O, HCl, HF, NH 3, CH 4, C 2 H 5 OH, organic polymers and many other substances. These substances do not have high strength, have low melting and boiling points, do not conduct electricity, and some of them are soluble in water or other solvents.
Non-molecular substances with covalent bonds or atomic substances (diamond, graphite, Si, SiO 2, SiC and others) form very strong crystals (with the exception of layered graphite), they are insoluble in water and other solvents, have high melting and boiling points, most of them they do not conduct electric current (except for graphite, which is electrically conductive, and semiconductors - silicon, germanium, etc.)
All ionic substances are naturally non-molecular. These are solid, refractory substances, solutions and melts of which conduct electric current. Many of them are soluble in water. It should be noted that in ionic substances, the crystals of which consist of complex ions, there are also covalent bonds, for example: (Na +) 2 (SO 4 2-), (K +) 3 (PO 4 3-), (NH 4 + )(NO 3-), etc. The atoms that make up complex ions are connected by covalent bonds.
Metals (substances with metallic bonds) very diverse in their physical properties. Among them there are liquid (Hg), very soft (Na, K) and very hard metals (W, Nb).
The characteristic physical properties of metals are their high electrical conductivity (unlike semiconductors, it decreases with increasing temperature), high heat capacity and ductility (for pure metals).
In the solid state, almost all substances are composed of crystals. Based on the type of structure and type of chemical bond, crystals (“crystal lattices”) are divided into atomic(crystals of non-molecular substances with covalent bonds), ionic(crystals of ionic substances), molecular(crystals of molecular substances with covalent bonds) and metal(crystals of substances with a metallic bond).
Tasks and tests on the topic "Topic 10. "Chemical bonding. Structure of matter."
- Types of chemical bond - Structure of matter grade 8–9
Lessons: 2 Assignments: 9 Tests: 1
- Assignments: 9 Tests: 1
After working through this topic, you should understand the following concepts: chemical bond, intermolecular bond, ionic bond, covalent bond, metallic bond, hydrogen bond, simple bond, double bond, triple bond, multiple bonds, non-polar bond, polar bond, electronegativity, bond polarization , - and -bond, hybridization of atomic orbitals, binding energy.
You must know the classification of substances by type of structure, by type of chemical bond, the dependence of the properties of simple and complex substances on the type of chemical bond and the type of “crystal lattice”.
You must be able to: determine the type of chemical bond in a substance, the type of hybridization, draw up diagrams of bond formation, use the concept of electronegativity, a number of electronegativity; know how electronegativity changes in chemical elements of the same period and one group to determine the polarity of a covalent bond.
After making sure that everything you need has been learned, proceed to completing the tasks. We wish you success.
Recommended reading:
- O. S. Gabrielyan, G. G. Lysova. Chemistry 11th grade. M., Bustard, 2002.
- G. E. Rudzitis, F. G. Feldman. Chemistry 11th grade. M., Education, 2001.
Why can atoms combine with each other and form molecules? What is the reason for the possible existence of substances that contain atoms of completely different chemical elements? These are global questions affecting the fundamental concepts of modern physical and chemical science. You can answer them by having an idea of the electronic structure of atoms and knowing the characteristics of the covalent bond, which is the basic basis for most classes of compounds. The purpose of our article is to become familiar with the mechanisms of formation of various types of chemical bonds and compounds containing them in their molecules.
Electronic structure of the atom
Electrically neutral particles of matter, which are its structural elements, have a structure that mirrors the structure of the Solar system. Just as the planets revolve around the central star - the Sun, so the electrons in an atom move around a positively charged nucleus. To characterize a covalent bond, the electrons located at the last energy level and furthest from the nucleus will be significant. Since their connection with the center of their own atom is minimal, they can easily be attracted by the nuclei of other atoms. This is very important for the occurrence of interatomic interactions leading to the formation of molecules. Why is the molecular form the main type of existence of matter on our planet? Let's figure it out.
Basic property of atoms
The ability of electrically neutral particles to interact, leading to a gain in energy, is their most important feature. Indeed, under normal conditions, the molecular state of a substance is more stable than the atomic state. The basic principles of modern atomic-molecular science explain both the principles of molecular formation and the characteristics of covalent bonds. Let us recall that there can be from 1 to 8 electrons per atom; in the latter case, the layer will be complete, and therefore very stable. The atoms of noble gases: argon, krypton, xenon - inert elements that complete each period in D.I. Mendeleev’s system - have this structure of the external level. The exception here would be helium, which has not 8, but only 2 electrons at the last level. The reason is simple: in the first period there are only two elements, the atoms of which have a single electron layer. All other chemical elements have from 1 to 7 electrons on the last, incomplete layer. In the process of interaction with each other, the atoms will tend to be filled with electrons to the octet and restore the configuration of the atom of the inert element. This state can be achieved in two ways: by losing one’s own or accepting someone else’s negatively charged particles. These forms of interaction explain how to determine which bond - ionic or covalent - will arise between the atoms entering the reaction.
Mechanisms of formation of a stable electronic configuration
Let's imagine that two simple substances enter into a compound reaction: sodium metal and chlorine gas. A substance of the salt class is formed - sodium chloride. It has an ionic type of chemical bond. Why and how did it arise? Let us again turn to the structure of the atoms of the starting substances. Sodium has only one electron in the last layer, weakly bound to the nucleus due to the large radius of the atom. The ionization energy of all alkali metals, which includes sodium, is low. Therefore, the electron of the outer level leaves the energy level, is attracted by the nucleus of the chlorine atom and remains in its space. This sets a precedent for the Cl atom to become a negatively charged ion. Now we are no longer dealing with electrically neutral particles, but with charged sodium cations and chlorine anions. In accordance with the laws of physics, electrostatic attraction forces arise between them, and the compound forms an ionic crystal lattice. The mechanism of formation of an ionic type of chemical bond that we have considered will help to more clearly clarify the specifics and main characteristics of a covalent bond.
Common electron pairs
If an ionic bond occurs between atoms of elements that differ greatly in electronegativity, i.e., metals and nonmetals, then the covalent type appears during the interaction of atoms of both the same and different nonmetallic elements. In the first case, it is customary to talk about a nonpolar, and in the other, about a polar form of a covalent bond. The mechanism of their formation is common: each of the atoms partially gives up electrons for common use, which are combined in pairs. But the spatial arrangement of electron pairs relative to the atomic nuclei will be different. On this basis, types of covalent bonds are distinguished - non-polar and polar. Most often, in chemical compounds consisting of atoms of non-metallic elements, there are pairs consisting of electrons with opposite spins, i.e., rotating around their nuclei in opposite directions. Since the movement of negatively charged particles in space leads to the formation of electron clouds, which ultimately ends in their mutual overlap. What are the consequences of this process for atoms and what does it lead to?
Physical properties of covalent bond
It turns out that a two-electron cloud with a high density appears between the centers of two interacting atoms. The electrostatic forces of attraction between the negatively charged cloud itself and the nuclei of atoms increase. A portion of energy is released and the distances between atomic centers decrease. For example, at the beginning of the formation of the H 2 molecule, the distance between the nuclei of hydrogen atoms is 1.06 A, after the clouds overlap and the formation of a common electron pair - 0.74 A. Examples of covalent bonds formed according to the mechanism described above can be found among both simple and among complex inorganic substances. Its main distinguishing feature is the presence of common electron pairs. As a result, after the emergence of a covalent bond between atoms, for example, hydrogen, each of them acquires the electronic configuration of inert helium, and the resulting molecule has a stable structure.
Spatial shape of the molecule
Another very important physical property of a covalent bond is directionality. It depends on the spatial configuration of the molecule of the substance. For example, when two electrons overlap with a spherical cloud shape, the appearance of the molecule is linear (hydrogen chloride or hydrogen bromide). The shape of the water molecules in which the s- and p-clouds hybridize is angular, and the very strong particles of nitrogen gas have the shape of a pyramid.
The structure of simple substances - nonmetals
Having found out what kind of bond is called covalent, what characteristics it has, now is the time to understand its varieties. If atoms of the same non-metal - chlorine, nitrogen, oxygen, bromine, etc. - interact with each other, then the corresponding simple substances are formed. Their common electron pairs are located at the same distance from the centers of the atoms, without moving. Compounds with a non-polar type of covalent bond have the following characteristics: low boiling and melting points, insolubility in water, dielectric properties. Next, we will find out which substances are characterized by a covalent bond, in which a displacement of common electron pairs occurs.
Electronegativity and its effect on the type of chemical bond
The property of a certain element to attract electrons to itself from an atom of another element in chemistry is called electronegativity. The scale of values for this parameter, proposed by L. Pauling, can be found in all textbooks on inorganic and general chemistry. Fluorine has its highest value - 4.1 eV, other active non-metals have a smaller value, and the lowest value is characteristic of alkali metals. If elements that differ in their electronegativity react with each other, then inevitably one, more active, will attract negatively charged particles of the atom of a more passive element to its nucleus. Thus, the physical properties of a covalent bond directly depend on the ability of the elements to donate electrons for common use. The common pairs formed in this case are no longer located symmetrically relative to the nuclei, but are shifted towards the more active element.
Features of connections with polar coupling
Substances in whose molecules the shared electron pairs are asymmetrical with respect to the atomic nuclei include hydrogen halides, acids, compounds of chalcogens with hydrogen, and acid oxides. These are sulfate and nitrate acids, oxides of sulfur and phosphorus, hydrogen sulfide, etc. For example, a hydrogen chloride molecule contains one common electron pair formed by unpaired electrons of hydrogen and chlorine. It is shifted closer to the center of the Cl atom, which is a more electronegative element. All substances with polar bonds in aqueous solutions dissociate into ions and conduct electric current. The compounds we have given also have higher melting and boiling points compared to simple non-metallic substances.
Methods for breaking chemical bonds
In organic chemistry, saturated hydrocarbons and halogens follow a radical mechanism. A mixture of methane and chlorine reacts in light and at ordinary temperatures in such a way that chlorine molecules begin to split into particles carrying unpaired electrons. In other words, the destruction of the common electron pair and the formation of very active radicals -Cl are observed. They are able to influence methane molecules in such a way that they break the covalent bond between carbon and hydrogen atoms. An active species -H is formed, and the free valency of the carbon atom accepts a chlorine radical, and the first reaction product is chloromethane. This mechanism of molecular breakdown is called homolytic. If the common pair of electrons is completely transferred to one of the atoms, then they speak of a heterolytic mechanism, characteristic of reactions taking place in aqueous solutions. In this case, polar water molecules will increase the rate of destruction of the chemical bonds of the soluble compound.
Double and triple bonds
The vast majority of organic substances and some inorganic compounds contain not one, but several common electron pairs in their molecules. The multiplicity of covalent bonds reduces the distance between atoms and increases the stability of compounds. They are usually referred to as chemically resistant. For example, a nitrogen molecule has three pairs of electrons; they are designated in the structural formula by three dashes and determine its strength. The simple substance nitrogen is chemically inert and can only react with other compounds, such as hydrogen, oxygen or metals, when heated or under elevated pressure, or in the presence of catalysts.
Double and triple bonds are inherent in such classes of organic compounds as unsaturated diene hydrocarbons, as well as substances of the ethylene or acetylene series. Multiple bonds determine the basic chemical properties: addition and polymerization reactions that occur at the places where they are broken.
In our article, we gave a general description of covalent bonds and examined its main types.
Electronegativity is the ability of atoms to displace electrons in their direction when forming a chemical bond. This concept was introduced by the American chemist L. Pauling (1932). Electronegativity characterizes the ability of an atom of a given element to attract a common electron pair in a molecule. Electronegativity values determined by various methods differ from each other. In educational practice, they most often use relative rather than absolute values of electronegativity. The most common is a scale in which the electronegativity of all elements is compared with the electronegativity of lithium, taken as one.
Among the elements of groups IA - VIIA:
The patterns of changes in electronegativity among d-block elements are much more complex.electronegativity, as a rule, increases in periods (“from left to right”) with increasing atomic number, and decreases in groups (“from top to bottom”).
These include: hydrogen, carbon, nitrogen, phosphorus, oxygen, sulfur, selenium, fluorine, chlorine, bromine and iodine. According to a number of characteristics, a special group of noble gases (helium-radon) is also classified as nonmetals.Elements with high electronegativity, the atoms of which have high electron affinity and high ionization energy, i.e., prone to the addition of an electron or the displacement of a pair of bonding electrons in their direction, are called nonmetals.
Metals include most of the elements of the Periodic Table.
This set of physical properties that distinguish metals from non-metals is explained by the special type of bond that exists in metals. All metals have a clearly defined crystal lattice. Along with atoms, its nodes contain metal cations, i.e. atoms that have lost their electrons. These electrons form a socialized electron cloud, the so-called electron gas. These electrons are in the force field of many nuclei. This bond is called metallic. The free migration of electrons throughout the volume of the crystal determines the special physical properties of metals.Metals are characterized by low electronegativity, i.e., low ionization energy and electron affinity. Metal atoms either donate electrons to nonmetal atoms or mix pairs of bonding electrons from themselves. Metals have a characteristic luster, high electrical conductivity and good thermal conductivity. They are mostly durable and malleable.
Metals include all d and f elements. If from the Periodic Table you mentally select only blocks of s- and p-elements, i.e., elements of group A and draw a diagonal from the upper left corner to the lower right corner, then it turns out that non-metallic elements are located on the right side of this diagonal, and metallic ones - in the left. Adjacent to the diagonal are elements that cannot be unambiguously classified as either metals or non-metals. These intermediate elements include: boron, silicon, germanium, arsenic, antimony, selenium, polonium and astatine.
Ideas about covalent and ionic bonds played an important role in the development of ideas about the structure of matter, however, the creation of new physical and chemical methods for studying the fine structure of matter and their use showed that the phenomenon of chemical bonding is much more complex. It is currently believed that any heteroatomic bond is both covalent and ionic, but in different proportions. Thus, the concept of covalent and ionic components of a heteroatomic bond is introduced. The greater the difference in electronegativity of the bonding atoms, the greater the polarity of the bond. When the difference is more than two units, the ionic component is almost always predominant. Let's compare two oxides: sodium oxide Na 2 O and chlorine oxide (VII) Cl 2 O 7. In sodium oxide, the partial charge on the oxygen atom is -0.81, and in chlorine oxide -0.02. This effectively means that the Na-O bond is 81% ionic and 19% covalent. The ionic component of the Cl-O bond is only 2%.
List of used literature
- Popkov V. A., Puzakov S. A. General chemistry: textbook. - M.: GEOTAR-Media, 2010. - 976 pp.: ISBN 978-5-9704-1570-2. [With. 35-37]
- Volkov, A.I., Zharsky, I.M. Big chemical reference book / A.I. Volkov, I.M. Zharsky. - Mn.: Modern School, 2005. - 608 with ISBN 985-6751-04-7.