In the ground state, carbon atom C (1s 2 2s 2 2p 2) has two unpaired electrons, due to which only two common electron pairs can be formed. However, in most of its compounds, carbon is tetravalent. This is explained by the fact that the carbon atom, absorbing a small amount of energy, goes into an excited state in which it has 4 unpaired electrons, i.e. capable of forming four covalent bonds and take part in the formation of four common electron pairs:
6 С 1 s 2 2s 2 2 p 2 6 С * 1 s 2 2s 1 2 p 3
| ↓ | |||||||||||||
| 1 | ↓ | p | ↓ | p | |||||||||
| s | s | ||||||||||||
The excitation energy is compensated by the formation of chemical bonds, which occurs with the release of energy.
Carbon atoms have the ability to form three types of hybridization of electron orbitals ( sp 3, sp 2, sp) and the formation of multiple (double and triple) bonds among themselves (Table 7).
Table 7
Types of hybridization and molecular geometry
Simple (single) s - communication is carried out at sp 3-hybridization, in which all four hybrid orbitals are equivalent and are directed in space at an angle of 109 o 29 'to each other and oriented to the vertices of a regular tetrahedron.
Rice. 19. Formation of a methane molecule CH 4
If hybrid carbon orbitals overlap with spherical ones s-orbitals of the hydrogen atom, then the simplest organic compound methane CH 4 is formed - a saturated hydrocarbon (Fig. 19).
Rice. 20. Tetrahedral arrangement of bonds in the methane molecule
Of great interest is the study of the bonds of carbon atoms with each other and with atoms of other elements. Let's consider the structure of the molecules of ethane, ethylene and acetylene.
The angles between all bonds in the ethane molecule are almost exactly equal to each other (Fig. 21) and do not differ from the C-H angles in the methane molecule.
Rice. 21. Ethane molecule C 2 H 6
Therefore, the carbon atoms are in a state sp 3-hybridization.
The hybridization of the electronic orbitals of carbon atoms may be incomplete, i.e. two ( sp 2–hybridization) or one ( sp-hybridization) of three R- orbitals. In this case, between the carbon atoms there are formed multiples(double or triple) communications. Hydrocarbons with multiple bonds are called unsaturated or unsaturated. A double bond (C=C) is formed when sp 2– hybridization. In this case, each carbon atom has one of three R- orbitals are not involved in hybridization, resulting in the formation of three sp 2– hybrid orbitals located in the same plane at an angle of 120° to each other, and non-hybrid 2 R The -orbital is located perpendicular to this plane. Two carbon atoms bond together to form one s-bond due to overlapping hybrid orbitals and one p-bond due to overlapping R-orbitals. The interaction of free hybrid orbitals of carbon with 1s-orbitals of hydrogen atoms leads to the formation of the ethylene molecule C 2 H 4 (Fig. 22), the simplest representative of unsaturated hydrocarbons.
Rice. 22. Formation of an ethylene molecule C 2 H 4
The overlap of electron orbitals in the case of a p-bond is less and the zones with increased electron density lie further from the atomic nuclei, therefore this bond is less strong than the s-bond.
A triple bond is formed by one s-bond and two p-bonds. In this case, the electron orbitals are in a state of sp-hybridization, the formation of which occurs due to one s- and one R- orbitals (Fig. 23).
Rice. 23. Formation of an acetylene molecule C 2 H 2
Two hybrid orbitals are located at an angle of 180° relative to each other, and the remaining non-hybrid two R-orbitals are located in two mutually perpendicular planes. The formation of a triple bond takes place in the acetylene molecule C 2 H 2.
A special type of bond occurs during the formation of a benzene molecule (C 6 H 6), the simplest representative of aromatic hydrocarbons.
Benzene contains six carbon atoms linked together in a ring (benzene ring), with each carbon atom in a state of sp 2 hybridization (Fig. 24).
All carbon atoms included in the benzene molecule are located in the same plane. Each carbon atom in the sp 2 hybridization state has one more non-hybrid p-orbital with an unpaired electron, which forms a p-bond (Fig. 25).
The axis of such a p-orbital is located perpendicular to the plane of the benzene molecule.
Rice. 24. sp 2 - orbitals of the benzene molecule C 6 H 6
Rice. 25. - bonds in the benzene molecule C 6 H 6
All six non-hybrid p orbitals form a common bonding molecular p orbital, and all six electrons combine to form a p electron sextet.
The boundary surface of such an orbital is located above and below the plane of the carbon s - skeleton. As a result of circular overlap, a single delocalized p-system arises, covering all carbon atoms of the cycle. Benzene is schematically depicted as a hexagon with a ring inside, which indicates that delocalization of electrons and corresponding bonds takes place.
Variety of inorganic and organic substances
Organic chemistry is chemistry carbon compounds. Inorganic carbon compounds include: carbon oxides, carbonic acid, carbonates and bicarbonates, carbides. Organic substances other than carbon contain hydrogen, oxygen, nitrogen, phosphorus, sulfur and other elements. Carbon atoms can form long unbranched and branched chains, rings, and attach other elements, so the number of organic compounds is close to 20 million, while inorganic substances number just over 100 thousand.
The basis for the development of organic chemistry is the theory of the structure of organic compounds by A. M. Butlerov. An important role in describing the structure of organic compounds belongs to the concept of valency, which characterizes the ability of atoms to form chemical bonds and determines their number. Carbon in organic compounds always tetravalent. The main postulate of the theory of A. M. Butlerov is the position on the chemical structure of matter, i.e. chemical bond. This order is displayed using structural formulas. Butlerov's theory states the idea that every substance has specific chemical structure And properties of substances depend on structure.
Theory of the chemical structure of organic compounds by A. M. Butlerov
Just as for inorganic chemistry the basis of development is the Periodic Law and the Periodic Table of Chemical Elements of D.I. Mendeleev, for organic chemistry it has become fundamental.
Theory of the chemical structure of organic compounds by A. M. Butlerov The main postulate of Butlerov’s theory is the position on the chemical structure of matter, which means the order, the sequence of mutual connection of atoms into molecules, i.e. chemical bond.
Chemical structure- the order of connection of atoms of chemical elements in a molecule according to their valence.
This order can be displayed using structural formulas in which the valencies of atoms are indicated by dashes: one line corresponds to the unit of valence of an atom of a chemical element. For example, for the organic substance methane, which has the molecular formula CH 4, the structural formula looks like this:
The main provisions of the theory of A. M. Butlerov:
Atoms in organic molecules are bonded to each other according to their valency. Carbon in organic compounds is always tetravalent, and its atoms are able to combine with each other, forming various chains.
· The properties of substances are determined not only by their qualitative and quantitative composition, but also by the order of connection of atoms in the molecule, i.e. chemical structure of a substance.
· The properties of organic compounds depend not only on the composition of the substance and the order of connection of atoms in its molecule, but also on mutual influence of atoms and groups of atoms on top of each other.
The theory of the structure of organic compounds is a dynamic and developing doctrine. As knowledge about the nature of chemical bonds and the influence of the electronic structure of molecules of organic substances developed, they began to use, in addition to empirical and structural ones, electronic formulas. Such formulas show the direction displacement of electron pairs in a molecule.
Quantum chemistry and the chemistry of the structure of organic compounds confirmed the doctrine of the spatial direction of chemical bonds (cis- and trans isomerism), studied the energy characteristics of mutual transitions in isomers, made it possible to judge the mutual influence of atoms in the molecules of various substances, created the prerequisites for predicting the types of isomerism and directions and mechanisms of chemical reactions.
Organic substances have a number of features.
· All organic substances contain carbon and hydrogen, so when burned they form carbon dioxide and water.
Organic matter complexly built and can have a huge molecular weight (proteins, fats, carbohydrates).
Organic substances can be arranged in rows similar in composition, structure and properties homologs.
· For organic substances it is characteristic isomerism.
Isomerism and homology of organic substances
The properties of organic substances depend not only on their composition, but also on order of connection of atoms in a molecule.
Isomerism- this is the phenomenon of the existence of different substances - isomers with the same qualitative and quantitative composition, i.e. with the same molecular formula.
There are two types of isomerism: structural and spatial(stereoisomerism). Structural isomers differ from each other in the order of bonding of atoms in the molecule; stereoisomers - the arrangement of atoms in space with the same order of bonds between them.
Main types of isomerism:
· Structural isomerism - substances differ in the order of bonding of atoms in molecules:
1) isomerism of the carbon skeleton;
2) position isomerism:
- multiple bonds;
- deputies;
- functional groups;
3) isomerism of homologous series (interclass).
· Spatial isomerism - molecules of substances differ not in the order of bonding of atoms, but in their position in space: cis-, trans-isomerism (geometric).
Classification of organic substances
It is known that the properties of organic substances are determined by their composition and chemical structure. Therefore, it is not surprising that the classification of organic compounds is based on the theory of structure - the theory of A. M. Butlerov. Organic substances are classified according to the presence and order of connection of atoms in their molecules. The most durable and least changeable part of the molecule of an organic substance is its skeleton - chain of carbon atoms. Depending on the order of connection of carbon atoms in this chain, substances are divided into acyclic, not containing closed chains of carbon atoms in molecules, and carbocyclic containing such chains (cycles) in molecules.
In addition to carbon and hydrogen atoms, molecules of organic substances can contain atoms of other chemical elements. Substances in whose molecules these so-called heteroatoms are included in a closed chain are classified as heterocyclic compounds.
Heteroatoms(oxygen, nitrogen, etc.) can be part of molecules and acyclic compounds, forming functional groups in them, for example,
hydroxyl
carbonyl
,
carboxyl
,
amino group
.
Functional group- a group of atoms that determines the most characteristic chemical properties of a substance and its belonging to a certain class of compounds.
Nomenclature of organic compounds
At the beginning of the development of organic chemistry, compounds to be discovered were assigned trivial names, often associated with the history of their production: acetic acid (which is the basis of wine vinegar), butyric acid (formed in butter), glycol (i.e. “sweet”), etc. As the number of new discovered substances increased, the need arose associate names with their structure. This is how rational names appeared: methylamine, diethylamine, ethyl alcohol, methyl ethyl ketone, which are based on the name of the simplest compound. For more complex compounds, rational nomenclature is not suitable.
The theory of structure of A. M. Butlerov provided the basis for the classification and nomenclature of organic compounds according to structural elements and the arrangement of carbon atoms in the molecule. Currently, the most commonly used nomenclature is developed by International Union of Pure and Applied Chemistry (IUPAC), which is called nomenclature IUPAC. IUPAC rules recommend several principles for the formation of names, one of them is the principle of substitution. Based on this, a replacement nomenclature has been developed, which is the most universal. Let us present several basic rules of substitutive nomenclature and consider their application using the example of a heterofunctional compound containing two functional groups - the amino acid leucine:
1. The names of compounds are based on the parent structure (the main chain of an acyclic molecule, a carbocyclic or heterocyclic system). The name of the parent structure forms the basis of the name, the root of the word.
In this case, the parent structure is a chain of five carbon atoms connected by single bonds. Thus, the root part of the name is pentane.
2. Characteristic groups and substituents (structural elements) are designated by prefixes and suffixes. Characteristic groups are divided by seniority. Order of precedence of the main groups:
The senior characteristic group is identified, which is designated in the suffix. All other substituents are named in the prefix in alphabetical order.
In this case, the senior characteristic group is carboxyl, i.e. this compound belongs to the class of carboxylic acids, so we add -ic acid to the radical part of the name. The second oldest group is the amino group, which is designated by the prefix amino-. In addition, the molecule contains the hydrocarbon substituent methyl-. Thus, the basis of the name is aminomethylpentanoic acid.
3. The name includes the designation of the double and triple bond, which comes immediately after the root.
The compound in question does not contain multiple bonds.
4. The atoms of the parent structure are numbered. Numbering begins from the end of the carbon chain to which the highest characteristic group is located closest:
The numbering of the chain begins with the carbon atom that is part of the carboxyl group, it is assigned the number 1. In this case, the amino group will be at carbon 2, and the methyl group will be at carbon 4.
Thus, the natural amino acid leucine, according to the rules of IUPAC nomenclature, is called 2-amino-4-methylpentanoic acid.
Hydrocarbons. Classification of hydrocarbons
Hydrocarbons- These are compounds consisting only of hydrogen and carbon atoms.
Depending on the structure of the carbon chain, organic compounds are divided into open-chain compounds - acyclic(aliphatic) and cyclic- with a closed chain of atoms.
Cyclic ones are divided into two groups: carbocyclic compounds(cycles are formed only by carbon atoms) and heterocyclic(the cycles also include other atoms, such as oxygen, nitrogen, sulfur).
Carbocyclic compounds, in turn, include two series of compounds: alicyclic And aromatic.
Aromatic compounds based on the molecular structure have flat carbon-containing cycles with a special closed system of p-electrons, forming a common π-system (a single π-electron cloud). Aromaticity is also characteristic of many heterocyclic compounds.
All other carbocyclic compounds belong to the alicyclic series.
Both acyclic (aliphatic) and cyclic hydrocarbons can contain multiple (double or triple) bonds. Such hydrocarbons are called unlimited(unsaturated) in contrast to limiting (saturated), containing only single bonds.
Saturated aliphatic hydrocarbons are called alkanes, they have the general formula C n H 2n+2, where n is the number of carbon atoms. Their old name is often used today - paraffins:
Unsaturated aliphatic hydrocarbons containing one double bond are called alkenes. They have the general formula C n H 2n:
Unsaturated aliphatic hydrocarbons with two double bonds are called alkadienes. Their general formula is C n H 2n-2:
Unsaturated aliphatic hydrocarbons with one triple bond are called alkynes. Their general formula is C n H 2n - 2:
Saturated alicyclic hydrocarbons - cycloalkanes, their general formula is C n H 2n:
A special group of hydrocarbons, aromatic, or arenas(with a closed common n-electronic system), known from the example of hydrocarbons with the general formula C n H 2n - 6:
Thus, if in their molecules one or more hydrogen atoms are replaced by other atoms or groups of atoms (halogens, hydroxyl groups, amino groups, etc.), hydrocarbon derivatives are formed: halogen derivatives, oxygen-containing, nitrogen-containing and other organic compounds.
Homologous series of hydrocarbons
Hydrocarbons and their derivatives with the same functional group form homologous series.
Homologous series name a series of compounds belonging to the same class (homologues), arranged in increasing order of their relative molecular masses, similar in structure and chemical properties, where each member differs from the previous one by the homologous difference CH 2. For example: CH 4 - methane, C 2 H 6 - ethane, C 3 H 8 - propane, C 4 H 10 - butane, etc. The similarity of the chemical properties of homologues greatly simplifies the study of organic compounds.
Hydrocarbon isomers
Those atoms or groups of atoms that determine the most characteristic properties of a given class of substances are called functional groups.
Halogen derivatives of hydrocarbons can be considered as products of the replacement of one or more hydrogen atoms in hydrocarbons with halogen atoms. In accordance with this, there can be finite and unsaturated mono-, di-, tri- (in the general case poly-) halogen derivatives.
General formula of monohalogen derivatives of saturated hydrocarbons:
and the composition is expressed by the formula
where R is the remainder of a saturated hydrocarbon (alkane), a hydrocarbon radical (this designation is used further when considering other classes of organic substances), G is a halogen atom (F, Cl, Br, I).
For example:
Here is one example of a dihalogen derivative:
TO oxygen-containing organic substances include alcohols, phenols, aldehydes, ketones, carboxylic acids, ethers and esters. Alcohols are derivatives of hydrocarbons in which one or more hydrogen atoms are replaced by hydroxyl groups.
Alcohols are called monohydric if they have one hydroxyl group, and saturated if they are derivatives of alkanes.
General formula for limit monohydric alcohols:
and their composition is expressed by the general formula:
For example:
Known examples polyhydric alcohols, i.e. having several hydroxyl groups:
Phenols- derivatives of aromatic hydrocarbons (benzene series), in which one or more hydrogen atoms in the benzene ring are replaced by hydroxyl groups.
The simplest representative with the formula C 6 H 5 OH or
called phenol.
Aldehydes and ketones- derivatives of hydrocarbons containing carbonyl group of atoms
(carbonyl).
In molecules aldehydes one carbonyl bond goes to combine with a hydrogen atom, the other - with a hydrocarbon radical. General formula of aldehydes:
For example:
When ketones the carbonyl group is connected to two (generally different) radicals, the general formula of ketones is:
For example:
The composition of saturated aldehydes and ketones is expressed by the formula C 2n H 2n O.
Carboxylic acids- hydrocarbon derivatives containing carboxyl groups
(or -COOH).
If there is one carboxyl group in an acid molecule, then the carboxylic acid is monobasic. General formula of saturated monobasic acids:
Their composition is expressed by the formula C n H 2n O 2.
For example:
Ethers are organic substances containing two hydrocarbon radicals connected by an oxygen atom: R-O-R or R 1 -O-R 2.
Radicals can be the same or different. The composition of ethers is expressed by the formula C n H 2n+2 O.
For example:
Esters- compounds formed by replacing the hydrogen atom of the carboxyl group in carboxylic acids with a hydrocarbon radical.
General formula of esters:
For example:
Nitro compounds- derivatives of hydrocarbons in which one or more hydrogen atoms are replaced by a nitro group -NO 2.
General formula of saturated mononitro compounds:
and the composition is expressed by the general formula C n H 2n+1 NO 2 .
For example:
Nitro derivatives of arenes:
Amines- compounds that are considered to be derivatives of ammonia (NH 3), in which the hydrogen atoms are replaced by hydrocarbon radicals. Depending on the nature of the radical, amines can be aliphatic, for example:
and aromatic, for example:
Depending on the number of hydrogen atoms replaced by radicals, the following are distinguished:
primary amines with the general formula:
secondary- with the general formula:
tertiary- with the general formula:
In a particular case, secondary and tertiary amines may have the same radicals.
Primary amines can also be considered as derivatives of hydrocarbons (alkanes), in which one hydrogen atom is replaced by an amino group -NH 2. The composition of saturated primary amines is expressed by the formula C n H 2n + 3 N.
For example:
Amino acids contain two functional groups connected to a hydrocarbon radical: amino group -NH 2 and carboxyl -COOH.
The general formula of α-amino acids (they are most important for the construction of proteins that make up living organisms):
The composition of saturated amino acids containing one amino group and one carboxyl is expressed by the formula C n H 2n + 1 NO 2.
For example:
Other important organic compounds are known that have several different or identical functional groups, long linear chains connected to benzene rings. In such cases, a strict determination of whether a substance belongs to a specific class is impossible. These compounds are often classified into specific groups of substances: carbohydrates, proteins, nucleic acids, antibiotics, alkaloids, etc.
Currently, many compounds are also known that can be classified as both organic and inorganic. x are called organoelement compounds. Some of them can be considered as hydrocarbon derivatives.
For example:
There are compounds that have the same molecular formula, expressing the composition of the substances.
The phenomenon of isomerism is that there can be several substances with different properties, having the same molecular composition, but different structures. These substances are called isomers.
In our case, these are interclass isomers: cycloalkanes and alkanes, alkadienes and alkynes, saturated monohydric alcohols and ethers, aldehydes and ketones, saturated monocarboxylic acids and esters.
Structural isomerism
The following varieties are distinguished structural isomerism: isomerism of the carbon skeleton, positional isomerism, isomerism of various classes of organic compounds (interclass isomerism).
Isomerism of the carbon skeleton is due to different bond order between carbon atoms, forming the skeleton of the molecule. As has already been shown, the molecular formula C 4 H 10 corresponds to two hydrocarbons: n-butane and isobutane. For the hydrocarbon C5H12, three isomers are possible: pentane, isopentane and neopentane.
As the number of carbon atoms in a molecule increases, the number of isomers increases rapidly. For hydrocarbon C 10 H 22 there are already 75 of them, and for hydrocarbon C 20 H 44 - 366,319.
Positional isomerism is due to different positions of the multiple bond, substituent, and functional group with the same carbon skeleton of the molecule:
Isomerism of different classes of organic compounds (interclass isomerism) is due to different positions and combinations of atoms in the molecules of substances that have the same molecular formula, but belong to different classes. Thus, the molecular formula C 6 H 12 corresponds to the unsaturated hydrocarbon hexene-1 and the cyclic hydrocarbon cyclohexane.
The isomers are a hydrocarbon related to alkynes - butine-1 and a hydrocarbon with two double bonds in the butadiene-1,3 chain:
Diethyl ether and butyl alcohol have the same molecular formula C 4 H 10 O:
The structural isomers are aminoacetic acid and nitroethane, corresponding to the molecular formula C 2 H 5 NO 2:
Isomers of this type contain different functional groups and belong to different classes of substances. Therefore, they differ in physical and chemical properties much more than carbon skeleton isomers or positional isomers.
Spatial isomerism
Spatial isomerism is divided into two types: geometric and optical.
Geometric isomerism is characteristic of compounds containing double bonds, and cyclic compounds. Since free rotation of atoms around a double bond or in a ring is impossible, the substituents can be located either on the same side of the plane of the double bond or ring (cis position) or on opposite sides (trans position). The designations cis and trans usually refer to a pair of identical substituents.
Geometric isomers differ in physical and chemical properties.
Optical isomerism occurs if the molecule is incompatible with its image in the mirror. This is possible when the carbon atom in the molecule has four different substituents. This atom is called asymmetric. An example of such a molecule is the α-aminopropionic acid (α-alanine) molecule CH 3 CH(NH 2)OH.
The α-alanine molecule cannot coincide with its mirror image during any movement. Such spatial isomers are called mirror, optical antipodes, or enantiomers. All physical and almost all chemical properties of such isomers are identical.
The study of optical isomerism is necessary when considering many reactions occurring in the body. Most of these reactions occur under the action of enzymes - biological catalysts. The molecules of these substances must fit the molecules of the compounds on which they act, like a key to a lock; therefore, the spatial structure, the relative arrangement of molecular sections and other spatial factors are of great importance for the course of these reactions. Such reactions are called stereoselective.
Most natural compounds are individual enantiomers, and their biological effects (from taste and smell to medicinal effects) differ sharply from the properties of their optical antipodes obtained in the laboratory. Such a difference in biological activity is of great importance, since it underlies the most important property of all living organisms - metabolism.
Isomerism Electronic structure of the carbon atom
Carbon, which is part of organic compounds, exhibits a constant valence. The last energy level of the carbon atom contains 4 electrons, two of which occupy a 2s orbital, which has a spherical shape, and two electrons occupy a 2p orbital, which has a dumbbell shape. When excited, one electron from the 2s orbital can move to one of the vacant 2p orbitals. This transition requires some energy expenditure (403 kJ/mol). As a result, the excited carbon atom has 4 unpaired electrons and its electronic configuration is expressed by the formula 2s 1 2p 3 .. Thus, in the case of the methane hydrocarbon (CH 4), the carbon atom forms 4 bonds with the s-electrons of hydrogen atoms. In this case, 1 s-s type bond (between the s-electron of the carbon atom and the s-electron of the hydrogen atom) and 3 p-s bonds (between 3 p-electrons of the carbon atom and 3 s-electrons of 3 hydrogen atoms) should be formed. This leads to the conclusion that the four covalent bonds formed by the carbon atom are unequal. However, practical experience in chemistry indicates that all 4 bonds in a methane molecule are absolutely equivalent, and the methane molecule has a tetrahedral structure with bond angles of 109.5 0, which could not be the case if the bonds were unequal. After all, only the orbitals of p-electrons are oriented in space along mutually perpendicular axes x, y, z, and the orbital of an s-electron has a spherical shape, so the direction of formation of a bond with this electron would be arbitrary. The theory of hybridization was able to explain this contradiction. L. Polling suggested that in any molecules there are no bonds isolated from each other. When bonds are formed, the orbitals of all valence electrons overlap. Several types are known hybridization of electron orbitals. It is assumed that in the molecule of methane and other alkanes, 4 electrons enter into hybridization.
Hybridization of carbon atom orbitals
Orbital hybridization is a change in the shape and energy of some electrons during the formation of a covalent bond, leading to more efficient orbital overlap and increased bond strength. Orbital hybridization occurs whenever electrons belonging to different types of orbitals participate in the formation of bonds.
1. sp 3 -hybridization(first valence state of carbon). During sp 3 hybridization, 3 p orbitals and one s orbital of an excited carbon atom interact in such a way that the resulting orbitals are absolutely identical in energy and symmetrically located in space. This transformation can be written like this:
During hybridization, the total number of orbitals does not change, but only their energy and shape change. It is shown that sp 3 -hybridization orbitals resemble a three-dimensional figure eight, one of the blades of which is much larger than the other. Four hybrid orbitals are extended from the center to the vertices of a regular tetrahedron at angles of 109.5 0. Bonds formed by hybrid electrons (for example, an s-sp 3 bond) are stronger than bonds formed by unhybridized p electrons (for example, an s-p bond). Because the hybrid sp 3 orbital provides a larger area of electron orbital overlap than the non-hybridized p orbital. Molecules in which sp 3 hybridization occurs have a tetrahedral structure. These, in addition to methane, include methane homologues, inorganic molecules such as ammonia. The figures show a hybridized orbital and a tetrahedral methane molecule.
The chemical bonds that arise in methane between carbon and hydrogen atoms are of the type σ-bonds (sp 3 -s-bond). Generally speaking, any sigma bond is characterized by the fact that the electron density of two interconnected atoms overlaps along the line connecting the centers (nuclei) of the atoms. σ-Bonds correspond to the maximum possible degree of overlap of atomic orbitals, so they are quite strong.
2. sp 2 -hybridization(second valence state of carbon). It arises as a result of the overlap of one 2s and two 2p orbitals. The resulting sp 2 -hybrid orbitals are located in the same plane at an angle of 120 0 to each other, and the non-hybridized p-orbital is perpendicular to it. The total number of orbitals does not change - there are four of them.
The sp 2 hybridization state occurs in alkene molecules, in carbonyl and carboxyl groups, i.e. in compounds containing a double bond. Thus, in the ethylene molecule, the hybridized electrons of the carbon atom form 3 σ bonds (two sp 2 -s type bonds between the carbon atom and hydrogen atoms and one sp 2 -sp 2 type bond between the carbon atoms). The remaining unhybridized p-electron of one carbon atom forms a π-bond with the unhybridized p-electron of the second carbon atom. A characteristic feature of the π bond is that the overlap of electron orbitals occurs outside the line connecting the two atoms. The overlap of orbitals occurs above and below the σ bond connecting both carbon atoms. Thus, a double bond is a combination of σ and π bonds. The first two figures show that in the ethylene molecule the bond angles between the atoms forming the ethylene molecule are 120 0 (corresponding to the spatial orientation of the three sp 2 hybrid orbitals). The figures show the formation of a π bond.
Since the overlap area of unhybridized p-orbitals in π bonds is less than the overlap area of orbitals in σ bonds, the π bond is less strong than the σ bond and is more easily broken in chemical reactions.
3. sp hybridization(third valence state of carbon). In the state of sp-hybridization, the carbon atom has two sp-hybrid orbitals located linearly at an angle of 180 0 to each other and two non-hybridized p-orbitals located in two mutually perpendicular planes. sp-hybridization is characteristic of alkynes and nitriles, i.e. for compounds containing a triple bond.
Thus, in an acetylene molecule, the bond angles between atoms are 180 o. The hybridized electrons of a carbon atom form 2 σ bonds (one sp-s bond between a carbon atom and a hydrogen atom and another sp-sp bond between carbon atoms. Two unhybridized p electrons of one carbon atom form two π bonds with unhybridized p electrons of the second carbon atom. The overlap of p-electron orbitals occurs not only above and below the σ bond, but also in front and behind, and the total cloud of p-electrons has a cylindrical shape. Thus, the triple bond is a combination of one σ bond and two π bonds. The presence in the acetylene molecule of less strong two π-bonds ensures the ability of this substance to enter into addition reactions with the cleavage of the triple bond.
Reference material for taking the test:
Mendeleev table
Solubility table
169375 0
Each atom has a certain number of electrons.
When entering into chemical reactions, atoms donate, gain, or share electrons, achieving the most stable electronic configuration. The configuration with the lowest energy (as in noble gas atoms) turns out to be the most stable. This pattern is called the “octet rule” (Fig. 1).
Rice. 1.
This rule applies to everyone types of connections. Electronic connections between atoms allow them to form stable structures, from the simplest crystals to complex biomolecules that ultimately form living systems. They differ from crystals in their continuous metabolism. At the same time, many chemical reactions proceed according to the mechanisms electronic transfer, which play a critical role in energy processes in the body.
A chemical bond is the force that holds together two or more atoms, ions, molecules, or any combination of these.
The nature of a chemical bond is universal: it is an electrostatic force of attraction between negatively charged electrons and positively charged nuclei, determined by the configuration of the electrons of the outer shell of atoms. The ability of an atom to form chemical bonds is called valence, or oxidation state. The concept of valence electrons- electrons that form chemical bonds, that is, located in the highest energy orbitals. Accordingly, the outer shell of the atom containing these orbitals is called valence shell. Currently, it is not enough to indicate the presence of a chemical bond, but it is necessary to clarify its type: ionic, covalent, dipole-dipole, metallic.
The first type of connection isionic connection
According to Lewis and Kossel's electronic valence theory, atoms can achieve a stable electronic configuration in two ways: first, by losing electrons, becoming cations, secondly, acquiring them, turning into anions. As a result of electron transfer, due to the electrostatic force of attraction between ions with charges of opposite signs, a chemical bond is formed, called by Kossel “ electrovalent"(now called ionic).
In this case, anions and cations form a stable electronic configuration with a filled outer electron shell. Typical ionic bonds are formed from cations T and II groups of the periodic table and anions of non-metallic elements of groups VI and VII (16 and 17 subgroups, respectively, chalcogens And halogens). The bonds of ionic compounds are unsaturated and non-directional, so they retain the possibility of electrostatic interaction with other ions. In Fig. Figures 2 and 3 show examples of ionic bonds corresponding to the Kossel model of electron transfer.
Rice. 2.
Rice. 3. Ionic bond in a molecule of table salt (NaCl)
Here it is appropriate to recall some properties that explain the behavior of substances in nature, in particular, consider the idea of acids And reasons.
Aqueous solutions of all these substances are electrolytes. They change color differently indicators. The mechanism of action of indicators was discovered by F.V. Ostwald. He showed that indicators are weak acids or bases, the color of which differs in the undissociated and dissociated states.
Bases can neutralize acids. Not all bases are soluble in water (for example, some organic compounds that do not contain OH groups are insoluble, in particular, triethylamine N(C 2 H 5) 3); soluble bases are called alkalis.
Aqueous solutions of acids undergo characteristic reactions:
a) with metal oxides - with the formation of salt and water;
b) with metals - with the formation of salt and hydrogen;
c) with carbonates - with the formation of salt, CO 2 and N 2 O.
The properties of acids and bases are described by several theories. In accordance with the theory of S.A. Arrhenius, an acid is a substance that dissociates to form ions N+ , while the base forms ions HE- . This theory does not take into account the existence of organic bases that do not have hydroxyl groups.
In accordance with proton According to the theory of Brønsted and Lowry, an acid is a substance containing molecules or ions that donate protons ( donors protons), and a base is a substance consisting of molecules or ions that accept protons ( acceptors protons). Note that in aqueous solutions, hydrogen ions exist in hydrated form, that is, in the form of hydronium ions H3O+ . This theory describes reactions not only with water and hydroxide ions, but also those carried out in the absence of a solvent or with a non-aqueous solvent.
For example, in the reaction between ammonia N.H. 3 (weak base) and hydrogen chloride in the gas phase, solid ammonium chloride is formed, and in an equilibrium mixture of two substances there are always 4 particles, two of which are acids, and the other two are bases:
This equilibrium mixture consists of two conjugate pairs of acids and bases:
1)N.H. 4+ and N.H. 3
2) HCl And Cl ‑
Here, in each conjugate pair, the acid and base differ by one proton. Every acid has a conjugate base. A strong acid has a weak conjugate base, and a weak acid has a strong conjugate base.
The Brønsted-Lowry theory helps explain the unique role of water for the life of the biosphere. Water, depending on the substance interacting with it, can exhibit the properties of either an acid or a base. For example, in reactions with aqueous solutions of acetic acid, water is a base, and in reactions with aqueous solutions of ammonia, it is an acid.
1) CH 3 COOH + H2O ↔ H3O + + CH 3 COO- . Here, an acetic acid molecule donates a proton to a water molecule;
2) NH 3 + H2O ↔ NH 4 + + HE- . Here, an ammonia molecule accepts a proton from a water molecule.
Thus, water can form two conjugate pairs:
1) H2O(acid) and HE- (conjugate base)
2) H 3 O+ (acid) and H2O(conjugate base).
In the first case, water donates a proton, and in the second, it accepts it.
This property is called amphiprotonism. Substances that can react as both acids and bases are called amphoteric. Such substances are often found in living nature. For example, amino acids can form salts with both acids and bases. Therefore, peptides easily form coordination compounds with the metal ions present.
Thus, a characteristic property of an ionic bond is the complete movement of the bonding electrons to one of the nuclei. This means that between the ions there is a region where the electron density is almost zero.
The second type of connection iscovalent connection
Atoms can form stable electronic configurations by sharing electrons.
Such a bond is formed when a pair of electrons is shared one at a time from everyone atom. In this case, the shared bond electrons are distributed equally between the atoms. Examples of covalent bonds include homonuclear diatomic molecules H 2 , N 2 , F 2. The same type of connection is found in allotropes O 2 and ozone O 3 and for a polyatomic molecule S 8 and also heteronuclear molecules hydrogen chloride HCl, carbon dioxide CO 2, methane CH 4, ethanol WITH 2 N 5 HE, sulfur hexafluoride SF 6, acetylene WITH 2 N 2. All these molecules share the same electrons, and their bonds are saturated and directed in the same way (Fig. 4).
It is important for biologists that double and triple bonds have reduced covalent atomic radii compared to a single bond.
Rice. 4. Covalent bond in a Cl 2 molecule.
Ionic and covalent types of bonds are two extreme cases of the many existing types of chemical bonds, and in practice most bonds are intermediate.
Compounds of two elements located at opposite ends of the same or different periods of the periodic system predominantly form ionic bonds. As elements move closer together within a period, the ionic nature of their compounds decreases, and the covalent character increases. For example, the halides and oxides of elements on the left side of the periodic table form predominantly ionic bonds ( NaCl, AgBr, BaSO 4, CaCO 3, KNO 3, CaO, NaOH), and the same compounds of elements on the right side of the table are covalent ( H 2 O, CO 2, NH 3, NO 2, CH 4, phenol C6H5OH, glucose C 6 H 12 O 6, ethanol C 2 H 5 OH).
The covalent bond, in turn, has one more modification.
In polyatomic ions and in complex biological molecules, both electrons can only come from one atom. It is called donor electron pair. An atom that shares this pair of electrons with a donor is called acceptor electron pair. This type of covalent bond is called coordination (donor-acceptor, ordative) communication(Fig. 5). This type of bond is most important for biology and medicine, since the chemistry of the d-elements most important for metabolism is largely described by coordination bonds.
Fig. 5.
As a rule, in a complex compound the metal atom acts as an acceptor of an electron pair; on the contrary, in ionic and covalent bonds the metal atom is an electron donor.
The essence of the covalent bond and its variety - the coordination bond - can be clarified with the help of another theory of acids and bases proposed by GN. Lewis. He somewhat expanded the semantic concept of the terms “acid” and “base” according to the Brønsted-Lowry theory. Lewis's theory explains the nature of the formation of complex ions and the participation of substances in nucleophilic substitution reactions, that is, in the formation of CS.
According to Lewis, an acid is a substance capable of forming a covalent bond by accepting an electron pair from a base. A Lewis base is a substance that has a lone electron pair, which, by donating electrons, forms a covalent bond with Lewis acid.
That is, Lewis's theory expands the range of acid-base reactions also to reactions in which protons do not participate at all. Moreover, the proton itself, according to this theory, is also an acid, since it is capable of accepting an electron pair.
Therefore, according to this theory, the cations are Lewis acids and the anions are Lewis bases. An example would be the following reactions:
It was noted above that the division of substances into ionic and covalent is relative, since complete electron transfer from metal atoms to acceptor atoms does not occur in covalent molecules. In compounds with ionic bonds, each ion is in the electric field of ions of the opposite sign, so they are mutually polarized, and their shells are deformed.
Polarizability determined by the electronic structure, charge and size of the ion; for anions it is higher than for cations. The highest polarizability among cations is for cations of greater charge and smaller size, for example, Hg 2+, Cd 2+, Pb 2+, Al 3+, Tl 3+. Has a strong polarizing effect N+ . Since the influence of ion polarization is two-way, it significantly changes the properties of the compounds they form.
The third type of connection isdipole-dipole connection
In addition to the listed types of communication, there are also dipole-dipole intermolecular interactions, also called van der Waals .
The strength of these interactions depends on the nature of the molecules.
There are three types of interactions: permanent dipole - permanent dipole ( dipole-dipole attraction); permanent dipole - induced dipole ( induction attraction); instantaneous dipole - induced dipole ( dispersive attraction, or London forces; rice. 6).
Rice. 6.
Only molecules with polar covalent bonds have a dipole-dipole moment ( HCl, NH 3, SO 2, H 2 O, C 6 H 5 Cl), and the bond strength is 1-2 Debaya(1D = 3.338 × 10‑30 coulomb meters - C × m).
In biochemistry, there is another type of connection - hydrogen connection, which is a limiting case dipole-dipole attraction. This bond is formed by the attraction between a hydrogen atom and a small electronegative atom, most often oxygen, fluorine and nitrogen. With large atoms that have similar electronegativity (such as chlorine and sulfur), the hydrogen bond is much weaker. The hydrogen atom is distinguished by one significant feature: when the bonding electrons are pulled away, its nucleus - the proton - is exposed and is no longer shielded by electrons.
Therefore, the atom turns into a large dipole.
A hydrogen bond, unlike a van der Waals bond, is formed not only during intermolecular interactions, but also within one molecule - intramolecular hydrogen bond. Hydrogen bonds play an important role in biochemistry, for example, to stabilize the structure of proteins in the form of an a-helix, or for the formation of a double helix of DNA (Fig. 7).
Fig.7.
Hydrogen and van der Waals bonds are much weaker than ionic, covalent and coordination bonds. The energy of intermolecular bonds is indicated in table. 1.
Table 1. Energy of intermolecular forces
Note: The degree of intermolecular interactions is reflected by the enthalpy of melting and evaporation (boiling). Ionic compounds require significantly more energy to separate ions than to separate molecules. The enthalpy of melting of ionic compounds is much higher than that of molecular compounds.
The fourth type of connection ismetal connection
Finally, there is another type of intermolecular bonds - metal: connection of positive ions of a metal lattice with free electrons. This type of connection does not occur in biological objects.
From a brief review of bond types, one detail becomes clear: an important parameter of a metal atom or ion - an electron donor, as well as an atom - an electron acceptor, is its size.
Without going into details, we note that the covalent radii of atoms, the ionic radii of metals and the van der Waals radii of interacting molecules increase as their atomic number increases in groups of the periodic table. In this case, the values of the ion radii are the smallest, and the van der Waals radii are the largest. As a rule, when moving down the group, the radii of all elements increase, both covalent and van der Waals.
Of greatest importance for biologists and physicians are coordination(donor-acceptor) bonds considered by coordination chemistry.
Medical bioinorganics. G.K. Barashkov
Most organic compounds have a molecular structure. Atoms in substances with a molecular type of structure always form only covalent bonds with each other, which is also observed in the case of organic compounds. Let us recall that covalent is a type of bond between atoms that is realized due to the fact that the atoms share part of their external electrons in order to acquire the electronic configuration of a noble gas.
Based on the number of shared electron pairs, covalent bonds in organic substances can be divided into single, double and triple. These types of connections are indicated in the graphic formula by one, two or three lines, respectively:
The multiplicity of a bond leads to a decrease in its length, so a single C-C bond has a length of 0.154 nm, a double C=C bond has a length of 0.134 nm, and a triple C≡C bond has a length of 0.120 nm.
Types of bonds according to the method of overlapping orbitals
As is known, orbitals can have different shapes, for example, s-orbitals are spherical and p-dumbbell-shaped. For this reason, bonds can also differ in the way electron orbitals overlap:
ϭ bonds - are formed when orbitals overlap in such a way that the area of their overlap is intersected by a line connecting the nuclei. Examples of ϭ-connections:
π-bonds - are formed when orbitals overlap, in two regions - above and below the line connecting the nuclei of atoms. Examples of π bonds:
How do you know when a molecule has π and ϭ bonds?
With the covalent type of bond, there is always a ϭ-bond between any two atoms, and a π-bond only in the case of multiple (double, triple) bonds. Wherein:
- Single bond is always a ϭ-bond
- A double bond always consists of one ϭ-bond and one π-bond
- A triple bond is always formed by one ϭ and two π bonds.
Let us indicate these types of bonds in the propinic acid molecule:
Hybridization of carbon atom orbitals
Orbital hybridization is a process in which orbitals that initially have different shapes and energies are mixed, instead forming the same number of hybrid orbitals that are equal in shape and energy.
So, for example, when mixing one s- and three p- four orbitals are formed sp 3-hybrid orbitals:
In the case of carbon atoms, hybridization always takes part s- orbital, and the number p-orbitals that can take part in hybridization vary from one to three p- orbitals.
How to determine the type of hybridization of a carbon atom in an organic molecule?
Depending on how many other atoms a carbon atom is bonded to, it is either in a state sp 3, or able sp 2, or able sp- hybridization:
Let's practice determining the type of hybridization of carbon atoms using the example of the following organic molecule:
The first carbon atom is bonded to two other atoms (1H and 1C), which means it is in a state sp-hybridization.
- The second carbon atom is bonded to two atoms - sp-hybridization
- The third carbon atom is bonded to four other atoms (two C and two H) – sp 3-hybridization
- The fourth carbon atom is bonded to three other atoms (2O and 1C) – sp 2-hybridization.
Radical. Functional group
The term radical most often means a hydrocarbon radical, which is the remainder of a hydrocarbon molecule without one hydrogen atom.
The name of a hydrocarbon radical is formed based on the name of the corresponding hydrocarbon by replacing the suffix –an to suffix –il .
Functional group - a structural fragment of an organic molecule (a certain group of atoms), which is responsible for its specific chemical properties.
Depending on which of the functional groups in the molecule of a substance is the eldest, the compound is classified into one class or another.
R – designation of hydrocarbon substituent (radical).
Radicals can contain multiple bonds, which can also be considered functional groups, since multiple bonds contribute to the chemical properties of a substance.
If a molecule of an organic substance contains two or more functional groups, such compounds are called polyfunctional.
In the ground state, carbon atom C (1s 2 2s 2 2p 2) has two unpaired electrons, due to which only two common electron pairs can be formed. However, in most of its compounds, carbon is tetravalent. This is explained by the fact that the carbon atom, absorbing a small amount of energy, goes into an excited state in which it has 4 unpaired electrons, i.e. capable of forming four covalent bonds and take part in the formation of four common electron pairs:
6 C 1s 2 2s 2 2p 2 6 C * 1s 2 2s 1 2p 3 .
| ↓ | ||||||||||||
| 1 | ↓ | p | ↓ | p | ||||||||
| s | s | |||||||||||
The excitation energy is compensated by the formation of chemical bonds, which occurs with the release of energy.
Carbon atoms have the ability to form three types of hybridization of electron orbitals ( sp 3, sp 2, sp) and the formation of multiple (double and triple) bonds among themselves (Table 2.2).
Table 2.2
Types of hybridization and molecular geometry
A simple (single) s-connection is carried out when sp 3-hybridization, in which all four hybrid orbitals are equivalent and are directed in space at an angle of 109°29 'to each other and oriented to the vertices of a regular tetrahedron (Fig. 2.8).
Rice. 2.8. Formation of methane molecule CH 4
If hybrid carbon orbitals overlap with spherical ones s-orbitals of the hydrogen atom, then the simplest organic compound methane CH 4 is formed - a saturated hydrocarbon.
Of great interest is the study of the bonds of carbon atoms with each other and with atoms of other elements. Let's consider the structure of the molecules of ethane, ethylene and acetylene.
The angles between all bonds in the ethane molecule are almost exactly equal to each other (Fig. 2.9) and do not differ from the C - H angles in the methane molecule.
Therefore, the carbon atoms are in a state sp 3-hybridization.
Rice. 2.9. Ethane molecule C 2 H 6
The hybridization of the electronic orbitals of carbon atoms may be incomplete, i.e. two ( sp 2-hybridization) or one ( sp-hybridization) of three R-orbitals. In this case, between the carbon atoms there are formed multiple connections (double or triple). Hydrocarbons with multiple bonds are called unsaturated or unsaturated. A double bond (C=C) is formed when sp 2-hybridization.
In this case, each carbon atom has one of three R-orbitals are not involved in hybridization, resulting in the formation of three sp 2-hybrid orbitals located in the same plane at an angle of 120° to each other, and non-hybrid 2 R The -orbital is located perpendicular to this plane. Two carbon atoms bond together to form one s-bond due to overlapping hybrid orbitals and one p-bond due to overlapping R-orbitals.
Interaction of free hybrid orbitals of carbon with 1 s-orbitals of hydrogen atoms leads to the formation of an ethylene molecule C 2 H 4 (Fig. 2.10) - the simplest representative of unsaturated hydrocarbons.
Rice. 2.10. Formation of an ethylene molecule C 2 H 4
The overlap of electron orbitals in the case of a p-bond is less and the zones with increased electron density lie further from the atomic nuclei, so this bond is less strong than the s-bond.
A triple bond is formed by one s-bond and two p-bonds. In this case, the electron orbitals are in a state of sp-hybridization, the formation of which occurs due to one s- and one R-orbitals (Fig. 2.11).
The two hybrid orbitals are located at an angle of 180° relative to each other, and the remaining non-hybrid two R-orbitals are located in two mutually perpendicular planes. The formation of a triple bond takes place in the acetylene molecule C 2 H 2 (see Fig. 2.11).
Rice. 2.11. Formation of an acetylene molecule C 2 H 2
A special type of bond occurs during the formation of a benzene molecule (C 6 H 6), the simplest representative of aromatic hydrocarbons.
Benzene contains six carbon atoms linked together in a ring (benzene ring), with each carbon atom in a state of sp 2 hybridization (Fig. 2.12).
Rice. 2.12. sp 2 – orbitals of the benzene molecule C 6 H 6
All carbon atoms included in the benzene molecule are located in the same plane. Each carbon atom in the sp 2 hybridization state has one more non-hybrid p-orbital with an unpaired electron, which forms a p-bond (Fig. 2.13).
The axis is like this R-orbitals are located perpendicular to the plane of the benzene molecule.
All six are non-hybrid R-orbitals form a common bonding molecular p-orbital, and all six electrons combine to form a p-electron sextet.
The boundary surface of such an orbital is located above and below the plane of the carbon s-skeleton. As a result of circular overlap, a single delocalized p-system arises, covering all carbon atoms of the cycle (Fig. 2.13).
Benzene is schematically depicted as a hexagon with a ring inside, which indicates that delocalization of electrons and corresponding bonds takes place.
Rice. 2.13. -bonds in the benzene molecule C 6 H 6
Ionic chemical bond
Ionic bond- a chemical bond formed as a result of mutual electrostatic attraction of oppositely charged ions, in which a stable state is achieved by complete transfer of the total electron density to an atom of a more electronegative element.
A purely ionic bond is an extreme case of a covalent bond.
In practice, the complete transfer of electrons from one atom to another atom through a bond is not realized, since each element has a greater or lesser (but not zero) EO, and any chemical bond will be covalent to some extent.
Such a bond occurs in the case of a large difference in the EO of atoms, for example, between cations s-metals of the first and second groups of the periodic system and anions of non-metals of groups VIА and VIIА (LiF, NaCl, CsF, etc.).
Unlike a covalent bond, ionic bond has no directionality . This is explained by the fact that the electric field of the ion has spherical symmetry, i.e. decreases with distance according to the same law in any direction. Therefore, the interaction between ions is independent of direction.
The interaction of two ions of opposite sign cannot lead to complete mutual compensation of their force fields. Because of this, they retain the ability to attract ions of the opposite sign in other directions. Therefore, unlike a covalent bond, ionic bonding is also characterized by unsaturation .
The lack of directionality and saturation in ionic bonds determines the tendency of ionic molecules to associate. All ionic compounds in the solid state have an ionic crystal lattice, in which each ion is surrounded by several ions of the opposite sign. In this case, all bonds of a given ion with neighboring ions are equivalent.
Metal connection
Metals are characterized by a number of special properties: electrical and thermal conductivity, a characteristic metallic luster, malleability, high ductility, and great strength. These specific properties of metals can be explained by a special type of chemical bond called metal .
A metallic bond is the result of overlapping delocalized orbitals of atoms approaching each other in the crystal lattice of a metal.
Most metals have a significant number of vacant orbitals and a small number of electrons in their outer electronic level.
Therefore, it is energetically more favorable for the electrons not to be localized, but to belong to the entire metal atom. At the lattice nodes of the metal there are positively charged ions, which are immersed in an electron “gas” distributed throughout the metal:
Me ↔ Me n + + n .
There is an electrostatic interaction between positively charged metal ions (Me n +) and non-localized electrons (n), which ensures the stability of the substance. The energy of this interaction is intermediate between the energies of covalent and molecular crystals. Therefore, elements with a purely metallic bond ( s-, And p-elements) are characterized by relatively high melting points and hardness.
The presence of electrons, which can freely move throughout the volume of the crystal, provides the specific properties of the metal
Hydrogen bond
Hydrogen bond – a special type of intermolecular interaction. Hydrogen atoms that are covalently bonded to an atom of an element that has a high electronegativity value (most often F, O, N, but also Cl, S, and C) carry a relatively high effective charge. As a result, such hydrogen atoms can interact electrostatically with atoms of these elements.
Thus, the H d + atom of one water molecule is oriented and interacts accordingly (as shown by three dots) with the O d - atom of another water molecule:
Bonds formed by an H atom located between two atoms of electronegative elements are called hydrogen:
d- d+ d-
A − H ××× B
The energy of a hydrogen bond is significantly less than the energy of a conventional covalent bond (150–400 kJ/mol), but this energy is sufficient to cause aggregation of the molecules of the corresponding compounds in the liquid state, for example, in liquid hydrogen fluoride HF (Fig. 2.14). For fluorine compounds it reaches about 40 kJ/mol.
Rice. 2.14. Aggregation of HF molecules due to hydrogen bonds
The length of a hydrogen bond is also shorter than the length of a covalent bond. Thus, in polymer (HF) n, the bond length is F−H = 0.092 nm, and the bond length is F∙∙∙H = 0.14 nm. For water, the bond length is O−H=0.096 nm, and the bond length O∙∙∙H=0.177 nm.
The formation of intermolecular hydrogen bonds leads to a significant change in the properties of substances: an increase in viscosity, dielectric constant, boiling and melting points.
Related information.