.
You know that atoms can combine with each other to form both simple and complex substances. In this case, various types of chemical bonds are formed: ionic, covalent (non-polar and polar), metallic and hydrogen. One of the most essential properties of the atoms of elements, which determine what kind of bond is formed between them - ionic or covalent, - is the electronegativity, i.e. the ability of atoms in a compound to attract electrons to itself.
A conditional quantitative assessment of electronegativity is given by the scale of relative electronegativity.
In periods, there is a general tendency for the growth of the electronegativity of the elements, and in groups - their decline. Electronegativity elements are arranged in a row, on the basis of which it is possible to compare the electronegativity of elements in different periods.
The type of chemical bond depends on how large the difference in the electronegativity values of the connecting atoms of the elements is. The more the atoms of the elements forming the bond differ in electronegativity, the more polar the chemical bond is. It is impossible to draw a sharp boundary between the types of chemical bonds. In most compounds, the type of chemical bond is intermediate; for example, a highly polar covalent chemical bond is close to an ionic bond. Depending on which of the limiting cases is closer in nature to the chemical bond, it is referred to as either an ionic or a covalent polar bond.
Ionic bond.An ionic bond is formed by the interaction of atoms that differ sharply from each other in electronegativity. For example, typical metals lithium (Li), sodium (Na), potassium (K), calcium (Ca), strontium (Sr), barium (Ba) form an ionic bond with typical non-metals, mainly halogens.
In addition to alkali metal halides, ionic bonds are also formed in compounds such as alkalis and salts. For example, in sodium hydroxide (NaOH) and sodium sulfate (Na 2 SO 4), ionic bonds exist only between sodium and oxygen atoms (the rest of the bonds are covalent polar).
Covalent non-polar bond.When atoms interact with the same electronegativity, molecules are formed with a covalent non-polar bond. Such a bond exists in the molecules of the following simple substances: H 2 , F 2 , Cl 2 , O 2 , N 2 . Chemical bonds in these gases are formed through common electron pairs, i.e. when the corresponding electron clouds overlap, due to the electron-nuclear interaction, which occurs when the atoms approach each other.
When compiling the electronic formulas of substances, it should be remembered that each common electron pair is a conditional image of an increased electron density resulting from the overlap of the corresponding electron clouds.
covalent polar bond.During the interaction of atoms, the values of the electronegativity of which differ, but not sharply, there is a shift of the common electron pair to a more electronegative atom. This is the most common type of chemical bond found in both inorganic and organic compounds.
Covalent bonds fully include those bonds that are formed by the donor-acceptor mechanism, for example, in hydronium and ammonium ions.
Metal connection.
The bond that is formed as a result of the interaction of relatively free electrons with metal ions is called a metallic bond. This type of bond is typical for simple substances - metals.
The essence of the process of formation of a metallic bond is as follows: metal atoms easily give up valence electrons and turn into positively charged ions. Relatively free electrons, detached from the atom, move between positive metal ions. A metallic bond arises between them, i.e., the electrons, as it were, cement the positive ions of the crystal lattice of metals.
Hydrogen bond.
A bond that forms between the hydrogen atoms of one molecule and an atom of a strongly electronegative element(O, N, F) another molecule is called a hydrogen bond.
The question may arise: why exactly does hydrogen form such a specific chemical bond?
This is because the atomic radius of hydrogen is very small. In addition, when a single electron is displaced or completely donated, hydrogen acquires a relatively high positive charge, due to which the hydrogen of one molecule interacts with atoms of electronegative elements that have a partial negative charge that is part of other molecules (HF, H 2 O, NH 3) .
Let's look at some examples. Usually we represent the composition of water with the chemical formula H 2 O. However, this is not entirely accurate. It would be more correct to denote the composition of water by the formula (H 2 O) n, where n \u003d 2.3.4, etc. This is due to the fact that individual water molecules are interconnected through hydrogen bonds.
Hydrogen bonds are usually denoted by dots. It is much weaker than an ionic or covalent bond, but stronger than the usual intermolecular interaction.
The presence of hydrogen bonds explains the increase in the volume of water with decreasing temperature. This is due to the fact that as the temperature decreases, the molecules become stronger and therefore the density of their “packing” decreases.
When studying organic chemistry, the following question also arose: why are the boiling points of alcohols much higher than those of the corresponding hydrocarbons? This is explained by the fact that hydrogen bonds are also formed between alcohol molecules.
An increase in the boiling point of alcohols also occurs due to the enlargement of their molecules.
The hydrogen bond is also characteristic of many other organic compounds (phenols, carboxylic acids, etc.). From courses in organic chemistry and general biology, you know that the presence of a hydrogen bond explains the secondary structure of proteins, the structure of the double helix of DNA, i.e., the phenomenon of complementarity.
E.N.FRENKEL
Chemistry tutorial
A guide for those who do not know, but want to learn and understand chemistry
Part I. Elements of General Chemistry
(first level of difficulty)
Continuation. See in No. 13, 18, 23/2007;
6/2008
Chapter 4
In the previous chapters of this manual, there were discussions about the fact that matter is made of molecules, and molecules are made of atoms. Have you ever wondered why the atoms that make up a molecule do not fly apart in different directions? What holds atoms together in a molecule?
Holds them chemical bond .
In order to understand the nature of the chemical bond, it is enough to recall a simple physical experiment. Two balls hanging side by side on strings do not “react” to each other in any way. But if you give one ball a positive charge and the other a negative charge, they will be attracted to each other. Isn't this the force that attracts atoms to each other? Indeed, studies have shown that the chemical bond is electrical in nature.
Where do charges come from in neutral atoms?
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When describing the structure of atoms, it was shown that all atoms, with the exception of noble gas atoms, tend to gain or give away electrons. The reason is the formation of a stable eight-electron external level (as in noble gases). When receiving or giving off electrons, electric charges arise and, as a result, electrostatic interaction of particles. This is how ionic bond , i.e. connection between ions.
Ions are stable charged particles that are formed as a result of receiving or giving off electrons.
For example, an atom of an active metal and an active non-metal participate in the reaction:
In this process, a metal atom (sodium) gives up electrons:
a) Is such a particle stable?
b) How many electrons are left in the sodium atom?
c) Will this particle have a charge?
Thus, in this process, a stable particle was formed (8 electrons at the outer level), which has a charge, because the nucleus of the sodium atom still has a charge of +11, and the remaining electrons have a net charge of -10. Therefore, the charge of the sodium ion is +1. A summary of this process looks like this:
What happens to the sulfur atom? This atom accepts electrons until the outer level is completed:
A simple calculation shows that this particle has a charge:
Oppositely charged ions are attracted, resulting in an ionic bond and an "ionic molecule":
There are other ways to form ions, which will be discussed in Chapter 6.
Formally, this molecular composition is attributed to sodium sulfide, although a substance consisting of ions has approximately the following structure (Fig. 1):
In this way, substances consisting of ions do not contain individual molecules! In this case, we can only speak of a conditional "ionic molecule".
Task 4.1. Show how the transition of electrons occurs when an ionic bond occurs between atoms:
a) calcium and chlorine;
b) aluminum and oxygen.
Remember! A metal atom donates outer electrons; the nonmetal atom accepts the missing electrons.
Conclusion. An ionic bond, according to the mechanism described above, is formed between the atoms of active metals and active non-metals.
Studies, however, show that the complete transition of electrons from one atom to another does not always occur. Very often, a chemical bond is formed not by giving and receiving electrons, but as a result of the formation of common electron pairs*. Such a connection is called covalent .
Covalent bond occurs due to the formation of common electron pairs. This type of bond is formed, for example, between atoms of non-metals. So, it is known that the nitrogen molecule consists of two atoms - N 2. How does a covalent bond between these atoms arise? To answer this question, it is necessary to consider the structure of the nitrogen atom:
Question. How many electrons are missing before the completion of the outer level?
Answer: Three electrons are missing. Therefore, denoting each electron of the outer level with a dot, we get:
Question. Why are three electrons indicated by single dots?
Answer: The point is that we want to show the formation of common pairs of electrons. A pair is two electrons. Such a pair occurs, in particular, if each atom contributes one electron to form a pair. The nitrogen atom is three electrons short of completing its outer level. This means that he must “prepare” three single electrons for the formation of future pairs (Fig. 2).
Received electronic formula of the molecule nitrogen, which shows that each nitrogen atom now has eight electrons (six of them are circled in an oval plus 2 of their electrons); three common pairs of electrons appeared between the atoms (the intersection of the circles).
Each pair of electrons corresponds to one covalent bond. How many covalent bonds are there? Three. Each bond (each common pair of electrons) will be shown with a dash (string stroke):
All these formulas, however, do not give an answer to the question: what binds atoms during the formation of a covalent bond? The electronic formula shows that a common pair of electrons is located between atoms. In this region of space, an excess negative charge arises. And the nuclei of atoms, as you know, have a positive charge. Thus, the nuclei of both atoms are attracted to a common negative charge, which arose due to common electron pairs (more precisely, the intersection of electron clouds) (Fig. 3).
Can there be such a bond between different atoms? Maybe. Let the nitrogen atom interact with the hydrogen atoms:
The structure of the hydrogen atom shows that the atom has one electron. How many such atoms need to be taken so that the nitrogen atom "gets what it wants" - three electrons? Obviously three hydrogen atoms
(Fig. 4):
Cross in fig. 4 denotes the electrons of the hydrogen atom. The electronic formula of the ammonia molecule shows that the nitrogen atom has eight electrons, and each hydrogen atom has two electrons (and there cannot be more at the first energy level).
The graphic formula shows that the nitrogen atom has a valence of three (three dashes, or three valence strokes), and each hydrogen atom has a valence of one (one dash each).
Although both N 2 and NH 3 molecules contain the same nitrogen atom, the chemical bonds between the atoms differ from each other. In the nitrogen molecule N 2 chemical bonds form identical atoms, so the common pairs of electrons are in the middle between the atoms. Atoms remain neutral. This chemical bond is called non-polar .
In the ammonia molecule NH 3, a chemical bond is formed different atoms. Therefore, one of the atoms (in this case, the nitrogen atom) attracts a common pair of electrons more strongly. Common pairs of electrons are displaced towards the nitrogen atom, and a small negative charge arises on it, and a positive charge arises on the hydrogen atom, electricity poles arose - a bond polar (Fig. 5).
Most substances built with the help of a covalent bond consist of individual molecules (Fig. 6).
From fig. 6 it can be seen that there are chemical bonds between atoms, but between molecules they are absent or insignificant.
The type of chemical bond affects the properties of a substance, its behavior in solutions. So, the greater, the more significant the attraction between the particles, the more difficult it is to tear them apart and the more difficult it is to transfer a solid substance into a gaseous or liquid state. Try to determine in the diagram below, between which particles the interaction force is greater and what chemical bond is formed in this case (Fig. 7).
If you carefully read the chapter, your answer will be as follows: the maximum interaction between particles occurs in the case of I (ionic bond). Therefore, all such substances are solid. The smallest interaction between uncharged particles (case III - non-polar covalent bond). These substances are usually gases.
Task 4.2. Determine what chemical bond is carried out between atoms in substances: NaCl, Hcl, Cl 2, AlCl 3, H 2 O. Give an explanation.
Task 4.3. Compose electronic and graphical formulas for those substances from task 4.2 in which you have determined the presence of a covalent bond. For an ionic bond, draw up electron transition schemes.
Chapter 5
There is no person on Earth who would not see solutions. And what is it?
A solution is a homogeneous mixture of two or more components (components or substances).
What is a homogeneous mixture? The homogeneity of a mixture implies that between its constituent substances no interface. In this case, it is impossible, at least visually, to determine how many substances formed a given mixture. For example, looking at tap water in a glass, it is difficult to assume that, in addition to water molecules, it contains a dozen more ions and molecules (O 2, CO 2, Ca 2+, etc.). And no microscope will help to see these particles.
But the absence of an interface is not the only sign of homogeneity. in a homogeneous mixture the composition of the mixture at any point is the same. Therefore, to obtain a solution, it is necessary to thoroughly mix the components (substances) that form it.
Solutions can have a different state of aggregation:
Gaseous solutions (for example, air - a mixture of gases O 2, N 2, CO 2, Ar);
Liquid solutions (eg, cologne, syrup, brine);
Solid solutions (for example, alloys).
One of the substances that form a solution is called solvent. The solvent has the same state of aggregation as the solution. So, for liquid solutions, it is a liquid: water, oil, gasoline, etc. Most often, aqueous solutions are used in practice. They will be discussed further (unless an appropriate reservation is made).
What happens when different substances are dissolved in water? Why do some substances dissolve well in water, while others do not? What determines solubility - the ability of a substance to dissolve in water?
Imagine that a piece of sugar is placed in a glass of warm water. He lay down, decreased in size and ... disappeared. Where? Is the law of conservation of matter (its mass, energy) really violated? No. Take a sip of the resulting solution, and you will see that the water is sweet, the sugar has not disappeared. But why is it not visible?
The fact is that in the course of dissolution, crushing (grinding) of the substance occurs. In this case, the sugar cube broke up into molecules, but we cannot see them. Yes, but why doesn't the sugar lying on the table break down into molecules? Why does a piece of margarine, dipped into water, also disappear? But because the crushing of the solute occurs under the action of a solvent, such as water. But the solvent will be able to "pull" the crystal, the solid into molecules, if it can "cling" to these particles. In other words, when a substance is dissolved, there must be interaction between a substance and a solvent.
When is such interaction possible? Only in the case when the structure of substances (both soluble and solvent) is similar, similar. The rule of alchemists has long been known: "like dissolves into like." In our examples, sugar molecules are polar and there are certain interaction forces between them and polar water molecules. Such forces are absent between non-polar fat molecules and polar water molecules. Therefore, fats do not dissolve in water. In this way, solubility depends on the nature of the solute and solvent.
As a result of the interaction between the solute and water, compounds are formed - hydrates. These can be very strong connections:
Such compounds exist as individual substances: bases, oxygen-containing acids. Naturally, during the formation of these compounds, strong chemical bonds arise, and heat is released. So when CaO (quicklime) dissolves in water, so much heat is released that the mixture boils.
But why does not the resulting solution heat up when sugar or salt is dissolved in water? First, not all hydrates are as strong as sulfuric acid or calcium hydroxide. There are salt hydrates (crystalline hydrates), which easily decompose when heated:
Secondly, during dissolution, as already mentioned, the process of crushing takes place. And energy is expended on this, heat is absorbed.
Because both processes occur simultaneously, the solution can either heat up or cool down, depending on which process is dominant.
Task 5.1. Determine which process - crushing or hydration - prevails in each case:
a) when dissolving sulfuric acid in water, if the solution is heated;
b) when dissolving ammonium nitrate in water, if the solution has cooled;
c) when sodium chloride is dissolved in water, if the temperature of the solution has not practically changed.
Since the temperature of the solution changes during dissolution, it is natural to assume that solubility depends on temperature. Indeed, the solubility of most solids increases with heating. The solubility of gases decreases when heated. Therefore, solids are usually dissolved in warm or hot water, and carbonated drinks are stored in the cold.
Solubility(ability to dissolve) substances does not depend on the grinding of the substance or the intensity of mixing. But by raising the temperature, grinding the substance, stirring the finished solution, you can speed up the dissolution process. By changing the conditions for obtaining a solution, it is possible to obtain solutions of different compositions. Naturally, there is a limit, having reached which, it is easy to find that the substance is no longer soluble in water. Such a solution is called rich. For highly soluble substances, a saturated solution will contain a lot of solute. So, a saturated solution of KNO 3 at 100 ° C contains 245 g of salt per 100 g of water (in 345 g of solution), this concentrated solution. Saturated solutions of poorly soluble substances contain negligible masses of dissolved compounds. So, a saturated solution of silver chloride contains 0.15 mg of AgCl in 100 g of water. This is very diluted solution.
Thus, if the solution contains a lot of solute in relation to the solvent, it is called concentrated, if there is little substance - dilute. Very often, its properties depend on the composition of the solution, and hence the application.
Thus, a dilute solution of acetic acid (table vinegar) is used as a flavoring seasoning, and a concentrated solution of this acid (acetic essence when taken orally) can cause a fatal burn.
In order to reflect the quantitative composition of solutions, use a value called mass fraction of solute :
where m(v-va) - the mass of the solute in the solution; m(p-ra) - the total mass of the solution containing the solute and the solvent.
So, if 100 g of vinegar contains 6 g of acetic acid, then we are talking about a 6% solution of acetic acid (this is table vinegar). Ways to solve problems using the concept of the mass fraction of a dissolved substance will be discussed in Chapter 8.
Conclusions on chapter 5. Solutions are homogeneous mixtures consisting of at least two substances, one of which is called a solvent, the other is a solute. When dissolved, this substance interacts with the solvent, due to which the solute is crushed. The composition of a solution is expressed using the mass fraction of the solute in the solution.
* These electron pairs occur at the intersection of electron clouds.
To be continued
3.3.1 Covalent bond - This is a two-center two-electron bond formed due to the overlap of electron clouds carrying unpaired electrons with antiparallel spins. As a rule, it is formed between atoms of one chemical element.
Quantitatively, it is characterized by valency. Element valence - this is its ability to form a certain number of chemical bonds due to free electrons located in the atomic valence zone.
A covalent bond is formed only by a pair of electrons located between atoms. It is called a divided pair. The remaining pairs of electrons are called lone pairs. They fill the shells and do not take part in binding. Communication between atoms can be carried out not only by one, but also by two or even three shared pairs. Such connections are called double and t swarm - multiple bonds.
3.3.1.1 Covalent non-polar bond. A bond carried out by the formation of electron pairs equally belonging to both atoms is called covalent non-polar. It arises between atoms with practically equal electronegativity (0.4 > ΔEO > 0) and, consequently, a uniform distribution of electron density between the nuclei of atoms in homonuclear molecules. For example, H 2 , O 2 , N 2 , Cl 2 , etc. The dipole moment of such bonds is zero. The CH bond in saturated hydrocarbons (for example, in CH 4) is considered practically non-polar, because ΔEO = 2.5 (C) - 2.1 (H) = 0.4.
3.3.1.2 Covalent polar bond. If a molecule is formed by two different atoms, then the overlap zone of electron clouds (orbitals) shifts towards one of the atoms, and such a bond is called polar . With such a connection, the probability of finding electrons near the nucleus of one of the atoms is higher. For example, HCl, H 2 S, PH 3.
Polar (asymmetric) covalent bond - connection between atoms with different electronegativity (2 > ΔEO > 0.4) and asymmetric distribution of a common electron pair. As a rule, it is formed between two non-metals.
The electron density of such a bond is shifted towards a more electronegative atom, which leads to the appearance on it of a partial negative charge (delta minus), and on a less electronegative atom - a partial positive charge (delta plus)
C - Cl
The direction of electron displacement is also indicated by an arrow:
CCl, CO, CN, OH, CMg.
The greater the difference in the electronegativity of the bonded atoms, the higher the polarity of the bond and the greater its dipole moment. Additional forces of attraction act between partial charges of opposite sign. Therefore, the more polar the bond, the stronger it is.
Except polarizability covalent bond has the property satiety - the ability of an atom to form as many covalent bonds as it has energetically available atomic orbitals. The third property of a covalent bond is its orientation.
3.3.2 Ionic bond. The driving force behind its formation is the same aspiration of atoms to the octet shell. But in a number of cases, such an “octet” shell can arise only when electrons are transferred from one atom to another. Therefore, as a rule, an ionic bond is formed between a metal and a non-metal.
Consider as an example the reaction between sodium (3s 1) and fluorine (2s 2 3s 5) atoms. Electronegativity difference in NaF compound
EO = 4.0 - 0.93 = 3.07
Sodium, having donated its 3s 1 electron to fluorine, becomes the Na + ion and remains with a filled 2s 2 2p 6 shell, which corresponds to the electronic configuration of the neon atom. Exactly the same electronic configuration is acquired by fluorine, having accepted one electron donated by sodium. As a result, electrostatic attraction forces arise between oppositely charged ions.
Ionic bond - an extreme case of a polar covalent bond, based on the electrostatic attraction of ions. Such a bond occurs when there is a large difference in the electronegativity of the bonded atoms (EO > 2), when a less electronegative atom almost completely gives up its valence electrons and turns into a cation, and another, more electronegative atom, attaches these electrons and becomes an anion. The interaction of ions of the opposite sign does not depend on the direction, and the Coulomb forces do not have the property of saturation. Because of this ionic bond has no space focus and satiety , since each ion is associated with a certain number of counterions (coordination number of the ion). Therefore, ionically bound compounds do not have a molecular structure and are solid substances that form ionic crystal lattices, with high melting and boiling points, they are highly polar, often salt-like, and electrically conductive in aqueous solutions. For example, MgS, NaCl, A 2 O 3. Compounds with purely ionic bonds practically do not exist, since there is always a certain amount of covalence due to the fact that a complete transition of one electron to another atom is not observed; in the most "ionic" substances, the proportion of bond ionicity does not exceed 90%. For example, in NaF, the bond polarization is about 80%.
In organic compounds, ionic bonds are quite rare, because. a carbon atom tends to neither lose nor gain electrons to form ions.
Valence elements in compounds with ionic bonds very often characterize oxidation state , which, in turn, corresponds to the charge of the ion of the element in the given compound.
Oxidation state is the conditional charge that an atom acquires as a result of the redistribution of electron density. Quantitatively, it is characterized by the number of electrons displaced from a less electronegative element to a more electronegative one. A positively charged ion is formed from the element that gave up its electrons, and a negative ion is formed from the element that received these electrons.
The element in highest oxidation state (maximally positive), has already given up all its valence electrons in the ABD. And since their number is determined by the number of the group in which the element is located, then highest oxidation state for most elements and will be equal to group number . Concerning lowest oxidation state (maximally negative), then it appears during the formation of an eight-electron shell, that is, in the case when the AVZ is completely filled. For non-metals it is calculated according to the formula group number - 8 . For metals is equal to zero because they cannot accept electrons.
For example, the AVZ of sulfur has the form: 3s 2 3p 4 . If an atom gives up all the electrons (six), then it will acquire the highest oxidation state +6 equal to the group number VI , if it takes the two necessary to complete the stable shell, it will acquire the lowest oxidation state –2 equal to Group number - 8 \u003d 6 - 8 \u003d -2.
3.3.3 Metal bond. Most metals have a number of properties that are general in nature and differ from the properties of other substances. Such properties are relatively high melting points, the ability to reflect light, high thermal and electrical conductivity. These features are explained by the existence in metals of a special type of interaction – metallic connection.
In accordance with the position in the periodic system, metal atoms have a small number of valence electrons, which are rather weakly bound to their nuclei and can easily be detached from them. As a result, positively charged ions appear in the crystal lattice of the metal, localized in certain positions of the crystal lattice, and a large number of delocalized (free) electrons that move relatively freely in the field of positive centers and carry out the connection between all metal atoms due to electrostatic attraction.
This is an important difference between metallic bonds and covalent bonds, which have a strict orientation in space. The bonding forces in metals are not localized and not directed, and the free electrons that form the "electron gas" cause high thermal and electrical conductivity. Therefore, in this case it is impossible to talk about the direction of the bonds, since the valence electrons are distributed almost uniformly over the crystal. This is precisely what explains, for example, the plasticity of metals, i.e., the possibility of displacement of ions and atoms in any direction
3.3.4 Donor-acceptor bond. In addition to the mechanism of formation of a covalent bond, according to which a common electron pair arises from the interaction of two electrons, there is also a special donor-acceptor mechanism . It lies in the fact that a covalent bond is formed as a result of the transition of an already existing (lone) electron pair donor (electron supplier) for the general use of the donor and acceptor (supplier of a free atomic orbital).
After formation, it is no different from covalent. The donor-acceptor mechanism is well illustrated by the scheme for the formation of an ammonium ion (Figure 9) (asterisks indicate the electrons of the outer level of the nitrogen atom):
Figure 9 - Scheme of the formation of the ammonium ion
The electronic formula of the AVZ of the nitrogen atom is 2s 2 2p 3, that is, it has three unpaired electrons that enter into a covalent bond with three hydrogen atoms (1s 1), each of which has one valence electron. In this case, an ammonia molecule NH 3 is formed, in which the unshared electron pair of nitrogen is preserved. If a hydrogen proton (1s 0) that does not have electrons approaches this molecule, then nitrogen will transfer its pair of electrons (donor) to this hydrogen atomic orbital (acceptor), resulting in the formation of an ammonium ion. In it, each hydrogen atom is connected to the nitrogen atom by a common electron pair, one of which is realized by the donor-acceptor mechanism. It is important to note that the H-N bonds formed by various mechanisms do not have any differences in properties. This phenomenon is due to the fact that at the moment of bond formation, the orbitals of the 2s– and 2p– electrons of the nitrogen atom change their shape. As a result, four completely identical orbitals arise.
The donors are usually atoms with a large number of electrons, but with a small number of unpaired electrons. For elements of period II, in addition to the nitrogen atom, oxygen (two lone pairs) and fluorine (three lone pairs) have such a possibility. For example, the hydrogen ion H + in aqueous solutions is never in a free state, since the hydronium ion H 3 O + is always formed from water molecules H 2 O and the ion H +. The hydronium ion is present in all aqueous solutions, although for simplicity the spelling is preserved symbol H + .
3.3.5 Hydrogen bond. A hydrogen atom bonded to a strongly electronegative element (nitrogen, oxygen, fluorine, etc.), which “pulls” a common electron pair onto itself, experiences a shortage of electrons and acquires an effective positive charge. Therefore, it is able to interact with the lone pair of electrons of another electronegative atom (which acquires an effective negative charge) of the same (intramolecular bond) or another molecule (intermolecular bond). As a result, there is hydrogen bond , which is graphically indicated by dots:
This bond is much weaker than other chemical bonds (the energy of its formation is 10 – 40 kJ/mol) and mainly has a partly electrostatic, partly donor-acceptor character.
The hydrogen bond plays an extremely important role in biological macromolecules, such inorganic compounds as H 2 O, H 2 F 2, NH 3. For example, O-H bonds in H 2 O have a noticeable polar character with an excess of negative charge – on the oxygen atom. The hydrogen atom, on the contrary, acquires a small positive charge + and can interact with lone pairs of electrons of the oxygen atom of the neighboring water molecule.
The interaction between water molecules turns out to be quite strong, such that even in water vapor there are dimers and trimers of the composition (H 2 O) 2, (H 2 O) 3, etc. In solutions, long chains of associates of this type can occur:
because the oxygen atom has two lone pairs of electrons.
The presence of hydrogen bonds explains the high boiling points of water, alcohols, carboxylic acids. Due to hydrogen bonds, water is characterized by such high melting and boiling points compared to H 2 E (E = S, Se, Te). If there were no hydrogen bonds, then water would melt at –100°C and boil at –80°C. Typical cases of association are observed for alcohols and organic acids.
Hydrogen bonds can occur both between different molecules and within a molecule if this molecule contains groups with donor and acceptor abilities. For example, it is intramolecular hydrogen bonds that play the main role in the formation of peptide chains that determine the structure of proteins. H-bonds affect the physical and chemical properties of a substance.
Hydrogen bonds do not form atoms of other elements , since the forces of electrostatic attraction of the opposite ends of the dipoles of polar bonds (О-Н, N-H, etc.) are rather weak and act only at short distances. Hydrogen, having the smallest atomic radius, allows such dipoles to approach each other so much that attractive forces become noticeable. No other element with a large atomic radius is capable of forming such bonds.
3.3.6 Forces of intermolecular interaction (van der Waals forces). In 1873, the Dutch scientist I. van der Waals suggested that there are forces that cause attraction between molecules. These forces were later called van der Waals forces. – the most versatile form of intermolecular bonding. The energy of the van der Waals bond is less than the hydrogen bond and is 2–20 kJ/∙mol.
Depending on the way the force is generated, they are divided into:
1) orientational (dipole-dipole or ion-dipole) - arise between polar molecules or between ions and polar molecules. When polar molecules approach each other, they are oriented in such a way that the positive side of one dipole is oriented towards the negative side of the other dipole (Figure 10).
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Figure 10 - Orientation interaction |
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2) induction (dipole - induced dipole or ion - induced dipole) - arise between polar molecules or ions and non-polar molecules, but capable of polarization. Dipoles can act on non-polar molecules, turning them into indicated (induced) dipoles. (Figure 11).
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Figure 11 - Inductive interaction |
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3) dispersive (induced dipole - induced dipole) - arise between non-polar molecules capable of polarization. In any molecule or atom of a noble gas, electric density fluctuations arise, as a result of which instantaneous dipoles appear, which in turn induce instantaneous dipoles in neighboring molecules. The movement of instantaneous dipoles becomes coordinated, their appearance and decay occur synchronously. As a result of the interaction of instantaneous dipoles, the energy of the system decreases (Figure 12).
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Figure 12 - Dispersion interaction |
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As a result of studying this topic, you will learn:
As a result of studying this topic, you will learn:
Study questions: |
5.1. covalent bond
A chemical bond is formed when two or more atoms approach each other, if as a result of their interaction a decrease in the total energy of the system occurs. The most stable electronic configurations of the outer electron shells of atoms are configurations of noble gas atoms, consisting of two or eight electrons. The outer electron shells of atoms of other elements contain from one to seven electrons, i.e. are incomplete. When a molecule is formed, atoms tend to acquire a stable two-electron or eight-electron shell. The valence electrons of atoms take part in the formation of a chemical bond.
A covalent bond is a chemical bond between two atoms, which is formed by electron pairs belonging simultaneously to these two atoms.
There are two mechanisms for the formation of a covalent bond: exchange and donor-acceptor.
5.1.1. Exchange mechanism for the formation of a covalent bond
exchange mechanism The formation of a covalent bond is realized due to the overlapping of electron clouds of electrons belonging to different atoms. For example, when two hydrogen atoms approach each other, the 1s electron orbitals overlap. As a result, a common pair of electrons appears, simultaneously belonging to both atoms. In this case, the chemical bond is formed by electrons having antiparallel spins, Fig. 5.1.
Rice. 5.1. Formation of a hydrogen molecule from two H atoms
5.1.2. Donor-acceptor mechanism of covalent bond formation
With the donor-acceptor mechanism for the formation of a covalent bond, the bond is also formed with the help of electron pairs. However, in this case, one atom (donor) provides its electron pair, and the other atom (acceptor) participates in the formation of the bond with its free orbital. An example of the implementation of a donor-acceptor bond is the formation of an ammonium ion NH 4 + during the interaction of ammonia NH 3 with a hydrogen cation H + .
In the NH 3 molecule, three electron pairs form three N - H bonds, the fourth electron pair belonging to the nitrogen atom is unshared. This electron pair can give a bond to the hydrogen ion, which has a free orbital. The result is an ammonium ion NH 4 + , fig. 5.2.
Rice. 5.2. Occurrence of a donor-acceptor bond during the formation of an ammonium ion
It should be noted that the four N – H covalent bonds existing in the NH 4 + ion are equivalent. In the ammonium ion, it is impossible to isolate the bond formed by the donor-acceptor mechanism.
5.1.3. Polar and non-polar covalent bond
If a covalent bond is formed by identical atoms, then the electron pair is located at the same distance between the nuclei of these atoms. Such a covalent bond is called non-polar. An example of molecules with a non-polar covalent bond are H 2, Cl 2, O 2, N 2, etc.
In the case of a polar covalent bond, the shared electron pair is shifted towards the atom with the higher electronegativity. This type of bond is realized in molecules formed by different atoms. The covalent polar bond takes place in the molecules of HCl, HBr, CO, NO, etc. For example, the formation of a polar covalent bond in the HCl molecule can be represented by the scheme, fig. 5.3:
Rice. 5.3. Formation of a covalent polar bond in the HC1 molecule
In the molecule under consideration, the electron pair is shifted to the chlorine atom, since its electronegativity (2.83) is greater than the electronegativity of the hydrogen atom (2.1).
5.1.4. Dipole moment and structure of molecules
The measure of bond polarity is its dipole moment μ:
μ = e l,
where e is the charge of an electron, l is the distance between the centers of positive and negative charges.
The dipole moment is a vector quantity. The concepts of "bond dipole moment" and "dipole moment of a molecule" coincide only for diatomic molecules. The dipole moment of a molecule is equal to the vector sum of the dipole moments of all bonds. Thus, the dipole moment of a polyatomic molecule depends on its structure.
In a linear CO 2 molecule, for example, each of the C–O bonds is polar. However, the CO 2 molecule is generally non-polar, since the dipole moments of the bonds compensate each other (Fig. 5.4). The dipole moment of a carbon dioxide molecule is m = 0.
In the corner H 2 O molecule, the polar H–O bonds are located at an angle of 104.5 o. The vector sum of the dipole moments of two H–O bonds is expressed by the diagonal of the parallelogram (Fig. 5.4). As a result, the dipole moment of the water molecule m is not equal to zero.
Rice. 5.4. Dipole moments of CO 2 and H 2 O molecules
5.1.5. Valency of elements in compounds with a covalent bond
The valency of atoms is determined by the number of unpaired electrons participating in the formation of common electron pairs with electrons of other atoms. Having one unpaired electron on the outer electron layer, the halogen atoms in the F 2, HCl, PBr 3 and CCl 4 molecules are monovalent. Elements of the oxygen subgroup contain two unpaired electrons on the outer layer, so in compounds such as O 2, H 2 O, H 2 S and SCl 2 they are divalent.
Since, in addition to the usual covalent bonds, a bond can be formed in molecules by a donor-acceptor mechanism, the valence of atoms also depends on the presence of lone electron pairs and free electron orbitals in them. A quantitative measure of valence is the number of chemical bonds by which a given atom is connected to other atoms.
The maximum valence of elements, as a rule, cannot exceed the number of the group in which they are located. The exception is the elements of the side subgroup of the first group Cu, Ag, Au, whose valency in compounds is greater than one. The valence electrons primarily include the electrons of the outer layers, however, for the elements of the secondary subgroups, the electrons of the penultimate (anterior) layers also take part in the formation of a chemical bond.
5.1.6. Valency of elements in normal and excited states
The valency of most chemical elements depends on whether these elements are in a normal or excited state. Electronic configuration of the Li atom: 1s 2 2s 1. The lithium atom at the outer level has one unpaired electron, i.e. lithium is monovalent. A very large expenditure of energy is required, associated with the transition of a 1s electron to a 2p orbital, in order to obtain trivalent lithium. This energy expenditure is so great that it is not compensated by the energy released during the formation of chemical bonds. In this regard, there are no compounds of trivalent lithium.
The configuration of the outer electron layer of the elements of the subgroup of beryllium ns 2 . This means that on the outer electron layer of these elements, there are two electrons with opposite spins in the ns cell orbital. The elements of the beryllium subgroup do not contain unpaired electrons, so their valency in the normal state is zero. In the excited state, the electronic configuration of the elements of the beryllium subgroup is ns 1 nр 1, i.e. elements form compounds in which they are divalent.
Valence possibilities of the boron atom
Consider the electronic configuration of the boron atom in the ground state: 1s 2 2s 2 2р 1 . The boron atom in the ground state contains one unpaired electron (Fig. 5.5), i.e. he is univalent. However, boron is not characterized by the formation of compounds in which it is monovalent. When a boron atom is excited, a transition of one 2s-electron to a 2p-orbital occurs (Fig. 5.5). The boron atom in an excited state has 3 unpaired electrons and can form compounds in which its valency is three.
Rice. 5.5. Valence states of the boron atom in the normal and excited states
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 one free 2p orbital in the boron atom, boron in compounds can form a fourth covalent bond, acting as an electron pair acceptor. Figure 5.6 shows how the BF molecule interacts with the F ion - , as a result of which an ion - is formed, in which boron forms four covalent bonds.
Rice. 5.6. Donor-acceptor mechanism for the formation of the fourth covalent bond at the boron atom
Valence possibilities of the nitrogen atom
Consider the electronic structure of the nitrogen atom (Fig. 5.7).
Rice. 5.7. The distribution of electrons in the orbitals of the nitrogen atom
From the presented diagram it can be seen that nitrogen has three unpaired electrons, it can form three chemical bonds and its valency is three. The transition of the nitrogen atom to an excited state is impossible, since the second energy level does not contain d-orbitals. At the same time, the nitrogen atom can provide an unshared electron pair of outer electrons 2s 2 to an atom that has a free orbital (acceptor). As a result, a fourth chemical bond of the nitrogen atom arises, as is the case, for example, in the ammonium ion (Fig. 5.2). Thus, the maximum covalence (the number of formed covalent bonds) of the nitrogen atom is four. In its compounds, nitrogen, unlike other elements of the fifth group, cannot be pentavalent.
Valence possibilities of phosphorus, sulfur and halogen atoms
Unlike nitrogen, oxygen, and fluorine atoms, phosphorus, sulfur, and chlorine atoms in the third period have free 3d cells, to which electrons can transfer. When a phosphorus atom is excited (Fig. 5.8), it has 5 unpaired electrons on its outer electron layer. As a result, in compounds, the phosphorus atom can be not only tri-, but also pentavalent.
Rice. 5.8. Distribution of valence electrons in orbits for a phosphorus atom in an excited state
In an excited state, sulfur, in addition to a valency of two, also exhibits a valence of four and six. In this case, the depairing of 3p and 3s electrons occurs sequentially (Fig. 5.9).
Rice. 5.9. Valence possibilities of the sulfur atom in an excited state
In the excited state, for all elements of the main subgroup of group V, except for fluorine, sequential depairing of first p- and then s-electron pairs is possible. As a result, these elements become tri-, penta-, and heptavalent (Fig. 5.10).
Rice. 5.10. Valence possibilities of chlorine, bromine and iodine atoms in an excited state
5.1.7. Length, energy and direction of a covalent bond
A covalent bond, as a rule, is formed between the atoms of non-metals. The main characteristics of a covalent bond are length, energy, and directionality.
Covalent bond length
The bond length is the distance between the nuclei of the atoms that form this bond. It is determined by experimental physical methods. The bond length can be estimated using the additivity rule, according to which the bond length in the AB molecule is approximately equal to half the sum of the bond lengths in the A 2 and B 2 molecules:
.
From top to bottom in the subgroups of the periodic system of elements, the length of the chemical bond increases, since the radii of atoms increase in this direction (Table 5.1). As the bond multiplicity increases, its length decreases.
Table 5.1.
The length of some chemical bonds
chemical bond |
Communication length, pm |
chemical bond |
Communication length, pm |
C - C |
|||
Bond energy
The measure of bond strength is the bond energy. Bond energy is determined by the energy required to break the bond and remove the atoms that form this bond to an infinite distance from each other. The covalent bond is very strong. Its energy ranges from several tens to several hundreds of kJ/mol. For an IСl 3 molecule, for example, Ebonds ≈40, and for N 2 and CO molecules, Ebonds ≈1000 kJ/mol.
From top to bottom in the subgroups of the periodic system of elements, the energy of a chemical bond decreases, since the bond length increases in this direction (Table 5.1). With an increase in the multiplicity of the connection, its energy increases (Table 5.2).
Table 5.2.
The energies of some chemical bonds
chemical bond |
bond energy, |
chemical bond |
bond energy, |
C - C |
|||
Saturation and directionality of a covalent bond
The most important properties of a covalent bond are its saturation and directionality. Saturation can be defined as the ability of atoms to form a limited number of covalent bonds. So a carbon atom can form only four covalent bonds, and an oxygen atom can form two. The maximum number of ordinary covalent bonds that an atom can form (excluding bonds formed by the donor-acceptor mechanism) is equal to the number of unpaired electrons.
Covalent bonds have a spatial orientation, since the overlap of orbitals during the formation of a single bond occurs along the line connecting the nuclei of atoms. The spatial arrangement of the electron orbitals of a molecule determines its geometry. The angles between chemical bonds are called bond angles.
The saturation and directionality of a covalent bond distinguishes this bond from an ionic bond, which, unlike a covalent bond, is unsaturated and non-directional.
Spatial structure of H 2 O and NH 3 molecules
Let us consider the orientation of a covalent bond using the example of H 2 O and NH 3 molecules.
The H 2 O molecule is formed from an oxygen atom and two hydrogen atoms. The oxygen atom has two unpaired p-electrons that occupy two orbitals located at right angles to each other. Hydrogen atoms have unpaired 1s electrons. The angle between the bonds formed by the p-electrons should be close to the angle between the orbitals of the p-electrons. Experimentally, however, it was found that the angle between the O–H bonds in a water molecule is 104.50. The increase in the angle compared to the angle of 90 o can be explained by the repulsive forces that act between the hydrogen atoms, fig. 5.11. Thus, the H 2 O molecule has an angular shape.
Three unpaired p-electrons of the nitrogen atom participate in the formation of the NH 3 molecule, the orbitals of which are located in three mutually perpendicular directions. Therefore, the three N–H bonds must be at angles to each other close to 90° (Fig. 5.11). The experimental value of the angle between bonds in the NH 3 molecule is 107.3°. The difference in the values of the angles between the bonds from the theoretical values is due, as in the case of the water molecule, to the mutual repulsion of the hydrogen atoms. In addition, the presented schemes do not take into account the possibility of participation of two electrons in 2s orbitals in the formation of chemical bonds.
Rice. 5.11. Overlapping of electronic orbitals during the formation of chemical bonds in H 2 O (a) and NH 3 (b) molecules
Consider the formation of the BeCl 2 molecule. An atom of beryllium in an excited state has two unpaired electrons: 2s and 2p. It can be assumed that the beryllium atom should form two bonds: one bond formed by the s-electron and one bond formed by the p-electron. These bonds must have different energies and different lengths. The BeCl 2 molecule in this case should not be linear, but angular. Experience, however, shows that the BeCl 2 molecule has a linear structure and both chemical bonds in it are equivalent. A similar situation is observed when considering the structure of BCl 3 and CCl 4 molecules – all bonds in these molecules are equivalent. The BC1 3 molecule has a planar structure, CC1 4 is tetrahedral.
To explain the structure of molecules such as BeCl 2, BCl 3 and CCl 4, Pauling and Slater(USA) introduced the concept of hybridization of atomic orbitals. They proposed to replace several atomic orbitals, not very different in their energy, with the same number of equivalent orbitals, called hybrid ones. These hybrid orbitals are made up of atomic orbitals as a result of their linear combination.
According to L. Pauling, when chemical bonds are formed by an atom that has electrons of various types in one layer and, therefore, not very different in energy (for example, s and p), it is possible to change the configuration of orbitals of various types, in which they are aligned in shape and energy . As a result, hybrid orbitals are formed, which have an asymmetric shape and are strongly elongated on one side of the nucleus. It is important to emphasize that the hybridization model is used in the case when electrons of different types participate in the formation of bonds, for example, s and p.
5.1.8.2. Various types of hybridization of atomic orbitals
sp hybridization
Hybridization of one s- and one R- orbitals ( sp- hybridization) realized, for example, in the formation of beryllium chloride. As shown above, in the excited state, the Be atom has two unpaired electrons, one of which occupies the 2s orbital and the other, the 2p orbital. When a chemical bond is formed, these two different orbitals are transformed into two identical hybrid orbitals directed at an angle of 180 ° to each other (Fig. 5.12). The linear arrangement of two hybrid orbitals corresponds to their minimum repulsion from each other. As a result, the BeCl 2 molecule has a linear structure - all three atoms are located on the same line.
Rice. 5.12. Scheme of overlapping electron orbitals during the formation of the BeCl 2 molecule
The structure of the acetylene molecule; sigma and pi bonds
Consider the scheme of overlapping electron orbitals in the formation of an acetylene molecule. In the acetylene molecule, each carbon atom is in the sp hybrid state. Two sp-hybrid orbitals are located at an angle of 1800 to each other; they form one σ-bond between carbon atoms and two σ-bonds with hydrogen atoms (Fig. 5.13).
Rice. 5.13. Scheme of the formation of s-bonds in the acetylene molecule
A σ-bond is a bond formed as a result of the overlap of electron orbitals along the line connecting the nuclei of atoms.
Each carbon atom in the acetylene molecule contains two more p-electrons, which do not take part in the formation of σ-bonds. The electron clouds of these electrons are located in mutually perpendicular planes and, overlapping with each other, form two more π-bonds between carbon atoms due to the lateral overlap of non-hybrid R-clouds (Fig. 5.14).
A π bond is a covalent chemical bond formed as a result of an increase in electron density on either side of a line connecting the nuclei of atoms.
Rice. 5.14. Scheme of the formation of σ - and π -bonds in the acetylene molecule.
Thus, in an acetylene molecule, a triple bond is formed between carbon atoms, which consists of one σ bond and two π bonds; σ -bonds are stronger than π-bonds.
sp2 hybridization
The structure of the BCl 3 molecule can be explained in terms of sp 2- hybridization. The boron atom in the excited state contains one s-electron and two p-electrons on the outer electron layer, i.e. three unpaired electrons. These three electron clouds can be converted into three equivalent hybrid orbitals. The minimum repulsion of three hybrid orbitals from each other corresponds to their location in the same plane at an angle of 120 o to each other (Fig. 5.15). Thus, the BCl 3 molecule has a planar shape.
Rice. 5.15. The planar structure of the BCl 3 molecule
sp 3 - hybridization
The valence orbitals of the carbon atom (s, p x, p y, p z) can be converted into four equivalent hybrid orbitals, which are located in space at an angle of 109.5 o to each other and directed to the vertices of the tetrahedron, in the center of which is the nucleus of the carbon atom (Fig. 5.16).
Rice. 5.16. Tetrahedral structure of the methane molecule
5.1.8.3. Hybridization involving lone electron pairs
The hybridization model can be used to explain the structure of molecules in which, in addition to binding, there are also unshared electron pairs. In water and ammonia molecules, the total number of electron pairs of the central atom (O and N) is four. In this case, the water molecule has two, and the ammonia molecule has one unshared electron pair. The formation of chemical bonds in these molecules can be explained by assuming that lone electron pairs can also fill hybrid orbitals. Unshared electron pairs occupy much more space in space than bonding pairs. As a result of the repulsion that occurs between lone and bonding electron pairs, the bond angles in water and ammonia molecules decrease, which are less than 109.5 o.
Rice. 5.17. sp 3 - hybridization involving lone electron pairs in H 2 O (A) and NH 3 (B) molecules
5.1.8.4. Establishment of the type of hybridization and determination of the structure of molecules
To establish the type of hybridization, and, consequently, the structure of molecules, the following rules must be used.
1. The type of hybridization of the central atom, which does not contain unshared electron pairs, is determined by the number of sigma bonds. If there are two such bonds, sp-hybridization takes place, three - sp 2 -hybridization, four - sp 3 -hybridization. Unshared electron pairs (in the absence of bonds formed by the donor-acceptor mechanism) are absent in molecules formed by atoms of beryllium, boron, carbon, silicon, i.e. the elements of the main subgroups II - IV groups.
2. If the central atom contains unshared electron pairs, then the number of hybrid orbitals and the type of hybridization are determined by the sum of the number of sigma bonds and the number of unshared electron pairs. Hybridization involving unshared electron pairs takes place in molecules formed by nitrogen, phosphorus, oxygen, and sulfur atoms, i.e. elements of the main subgroups of groups V and VI.
3. The geometric shape of the molecules is determined by the type of hybridization of the central atom (Table 5.3).
Table 5.3.
Valence angles, geometric shape of molecules depending on the number of hybrid orbitals and the type of hybridization of the central atom
5.2. Ionic bond
Ionic bonding is carried out by electrostatic attraction between oppositely charged ions. These ions are formed as a result of the transfer of electrons from one atom to another. An ionic bond is formed between atoms that have large differences in electronegativity (usually greater than 1.7 on the Pauling scale), for example, between alkali metals and halogens.
Let us consider the appearance of an ionic bond using the example of the formation of NaCl. From the electronic formulas of the atoms Na 1s 2 2s 2 2p 6 3s 1 and Cl 1s 2 2s 2 2p 6 3s 2 3p 5, it is clear that to complete the external level, it is easier for the sodium atom to give one electron than to attach seven, and it is easier for the chlorine atom to attach one, than give seven. In chemical reactions, the sodium atom donates one electron, and the chlorine atom accepts it. As a result, the electron shells of sodium and chlorine atoms turn into stable electron shells of noble gases (the electronic configuration of the sodium cation is Na + 1s 2 2s 2 2p 6, and the electronic configuration of the chlorine anion Cl is 1s 2 2s 2 2p 6 3s 2 3p 6). The electrostatic interaction of ions leads to the formation of the NaCl molecule.
The main characteristics of the ionic bond and the properties of ionic compounds
1. An ionic bond is a strong chemical bond. The energy of this bond is about 300 – 700 kJ/mol.
2. Unlike a covalent bond, an ionic bond is non-directional, since an ion can attract ions of the opposite sign to itself in any direction.
3. Unlike a covalent bond, an ionic bond is unsaturated, since the interaction of ions of the opposite sign does not lead to complete mutual compensation of their force fields.
4. In the process of formation of molecules with an ionic bond, there is no complete transfer of electrons, therefore, a 100% ionic bond does not exist in nature. In the NaCl molecule, the chemical bond is only 80% ionic.
5. Ionic compounds are crystalline solids with high melting and boiling points.
6. Most ionic compounds dissolve in water. Solutions and melts of ionic compounds conduct electric current.
5.3. metal connection
Metal atoms at the outer energy level contain a small number of valence electrons. Since the ionization energy of metal atoms is low, valence electrons are weakly retained in these atoms. As a result, positively charged ions and free electrons appear in the crystal lattice of metals. In this case, the metal cations are located at the nodes of their crystal lattice, and the electrons move freely in the field of positive centers, forming the so-called "electron gas". The presence of a negatively charged electron between two cations leads to the fact that each cation interacts with this electron. Thus, a metallic bond is a bond between positive ions in metal crystals, which is carried out by the attraction of electrons that move freely throughout the crystal.
Since the valence electrons in the metal are evenly distributed throughout the crystal, the metallic bond, like the ionic one, is an undirected bond. Unlike a covalent bond, a metallic bond is an unsaturated bond. From a covalent bond metallic bond differs also in durability. The energy of a metallic bond is about three to four times less than the energy of a covalent bond.
Due to the high mobility of the electron gas, metals are characterized by high electrical and thermal conductivity.
5.4. hydrogen bond
In the molecules of compounds HF, H 2 O, NH 3, there are hydrogen bonds with a strongly electronegative element (H–F, H–O, H–N). Between the molecules of such compounds can be formed intermolecular hydrogen bonds. In some organic molecules containing H–O, H–N bonds, intramolecular hydrogen bonds.
The mechanism of hydrogen bond formation is partly electrostatic, partly donor-acceptor. In this case, the atom of a strongly electronegative element (F, O, N) acts as an electron pair donor, and the hydrogen atoms connected to these atoms act as an acceptor. As with covalent bonds, hydrogen bonds are characterized by orientation in space and saturability.
The hydrogen bond is usually denoted by dots: H ··· F. The hydrogen bond is more pronounced, the greater the electronegativity of the partner atom and the smaller its size. It is characteristic primarily for fluorine compounds, as well as oxygen, to a lesser extent nitrogen, to an even lesser extent for chlorine and sulfur. Accordingly, the energy of the hydrogen bond also changes (Table 5.4).
Table 5.4.
Average values of hydrogen bond energies
Intermolecular and intramolecular hydrogen bonding
Thanks to hydrogen bonds, molecules are combined into dimers and more complex associates. For example, the formation of a formic acid dimer can be represented by the following scheme (Fig. 5.18).
Rice. 5.18. Formation of intermolecular hydrogen bonds in formic acid
Long chains of associates (H 2 O) n can appear in water (Fig. 5.19).
Rice. 5.19. Formation of a chain of associates in liquid water due to intermolecular hydrogen bonds
Each H 2 O molecule can form four hydrogen bonds, while an HF molecule can form only two.
Hydrogen bonds can occur both between different molecules (intermolecular hydrogen bond) and within a molecule (intramolecular hydrogen bond). Examples of the formation of an intramolecular bond for some organic substances are shown in fig. 5.20.
Rice. 5.20. Formation of an intramolecular hydrogen bond in the molecules of various organic compounds
The effect of hydrogen bonding on the properties of substances
The most convenient indicator of the existence of an intermolecular hydrogen bond is the boiling point of a substance. The higher boiling point of water (100 o C compared to hydrogen compounds of the elements of the oxygen subgroup (H 2 S, H 2 Se, H 2 Te) is due to the presence of hydrogen bonds: additional energy is required to destroy intermolecular hydrogen bonds in water.
The hydrogen bond can significantly affect the structure and properties of substances. The existence of intermolecular hydrogen bonds increases the melting and boiling points of substances. The presence of an intramolecular hydrogen bond leads to the fact that the molecule of deoxyribonucleic acid (DNA) is folded into a double helix in water.
Hydrogen bonding also plays an important role in dissolution processes, since solubility also depends on the ability of the compound to form hydrogen bonds with the solvent. As a result, substances containing OH groups such as sugar, glucose, alcohols, carboxylic acids, as a rule, are highly soluble in water.
5.5. Types of crystal lattices
Solids, as a rule, have a crystalline structure. The particles that make up crystals (atoms, ions or molecules) are located at strictly defined points in space, forming a crystal lattice. The crystal lattice consists of elementary cells that retain the structural features characteristic of this lattice. The points where the particles are located are called lattice nodes. Depending on the type of particles located at the lattice sites and on the nature of the connection between them, 4 types of crystal lattices are distinguished.
5.5.1. Atomic crystal lattice
At the nodes of atomic crystal lattices there are atoms interconnected by covalent bonds. Substances having an atomic lattice include diamond, silicon, carbides, silicides, etc. In the structure of an atomic crystal, it is impossible to single out individual molecules; the entire crystal is considered as one giant molecule. The structure of diamond is shown in fig. 5.21. A diamond is made up of carbon atoms, each bonded to four neighboring atoms. Due to the fact that covalent bonds are strong, all substances having atomic lattices are refractory, solid and low volatile. They are slightly soluble in water.
Rice. 5.21. Diamond crystal lattice
5.5.2. Molecular crystal lattice
Molecules are located at the nodes of molecular crystal lattices, interconnected by weak intermolecular forces. Therefore, substances with a molecular lattice have low hardness, they are fusible, are characterized by significant volatility, are slightly soluble in water, and their solutions, as a rule, do not conduct electric current. A lot of substances with a molecular crystal lattice are known. These are solid hydrogen, chlorine, carbon monoxide (IV) and other substances that are in a gaseous state at ordinary temperatures. Most crystalline organic compounds have a molecular lattice.
5.5.3. Ionic crystal lattice
Crystal lattices, at the nodes of which ions are located, are called ionic. They are formed by substances with an ionic bond, for example, alkali metal halides. In ionic crystals, individual molecules cannot be distinguished; the entire crystal can be considered as one macromolecule. The bonds between ions are strong, so substances with an ionic lattice have low volatility, high melting and boiling points. The crystal lattice of sodium chloride is shown in fig. 5.22.
Rice. 5.22. Crystal lattice of sodium chloride
In this figure, light balls are Na + ions, dark balls are Cl - ions. On the left in fig. 5.22 shows the unit cell of NaCI.
5.5.4. metal crystal lattice
Metals in the solid state form metallic crystal lattices. At the nodes of such lattices there are positive metal ions, and valence electrons move freely between them. The electrons electrostatically attract the cations, thereby giving stability to the metal lattice. Such a structure of the lattice determines the high thermal conductivity, electrical conductivity and plasticity of metals - during mechanical deformation, bonds are not broken and the crystal is not destroyed, since the ions that make it up seem to float in a cloud of electron gas. On fig. 5.23 shows the crystal lattice of sodium.
Rice. 5.23. The crystal lattice of sodium