covalent bond

[gòng jià jiàn]

Chemical bond

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This entry is made by Compilation and application of scientific encyclopedia entries of "Science Popularization China" to examine.

Covalent bond Chemical bond One of two or more atom Use their outer electrons together, and achieve electrons under ideal conditions saturated The state of the, from which form Relatively stable Chemical structure , such a strong interaction between several adjacent atoms sharing electrons and sharing electrons is called covalent bond. Its essence is Atomic orbital After overlapping, it appears in two Nucleus Between Electronics An electrical interaction with two nuclei.

Chinese name: covalent bond
Foreign name: covalent bond

Category: Chemical bond
Essence: Electrostatic interaction between atoms
Classification: σ bond, π bond, coordination bond, etc

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Early history

Figure 1

stay ancient Greek Chemistry has not yet been separated from natural philosophy, Atomism They had the most original idea of chemical bond, Empedochle （ Empedocles ）He believes that the world consists of "Qi water , soil fire ”These four elements are composed. When these four elements are split and reassembled in a new arrangement under the action of "love" and "hate", the matter has undergone a qualitative change. This force can be regarded as the earliest thought of chemical bond.

Later, the atomists Democritus Assumption, atom There is a kind of“ hook ”It can also be said that the surface is rough, so that they stick together when they collide with each other, forming a stable aggregate. Democritus' idea of chemical bond is more advanced than that of previous natural philosophers, and he has eliminated the idealistic factors in such ideas.

Medieval J R. Grauber put forward the idea that materials of the same kind are mutually exclusive. Then came the theory of affinity for matter binding, which believed that the particles of matter had affinity, so they were attracted to each other and bound together. In short, people's hazy understanding of chemical bonds inspired later chemists.

modern history

In the 18th century, phlogiston（ phlogiston ）The concept of Ernst Starr （ Ernst Stahl ）、 Henry Cavendish （ Henry Cavendish ）And Joseph Priestley （ Joseph Priestley ）And other advanced chemists. At that time, Newtonian mechanics It has been proposed that they hope to combine the force between atoms with Newtonian mechanics to give Classical physics However, limited to the conditions at that time, this is undoubtedly impossible.

In 1916, German chemist Walter Kossel（ Walther Kossel ）After investigating a large number of facts, we came to the conclusion that the atoms of any element must make the outermost layer meet the 8-electron stable structure, but Kossel only explained Ionic compound The formation process of the covalent bond does not explain the formation of the covalent bond.^[16]

In 1919, chemist Owen Langmuir It is the first time to use "covalence" to describe atom Bonding process between^[1] " we shall denote by the term covalence the number of pairs of electrons which a given atom shares with its neighbors^[2] ”(We should use the word "covalence" to express the passage between atoms Shared electron pair Resulting force).

In 1922, Niels Bohr （ N.Bohr ）From the perspective of quantization, Rutherford's nuclear model was reexamined, which provided a new platform for chemists to understand chemical bonds. He believed that electrons should be located in certain orbits, and can transition between different orbits. Stationary transitions can be well explained Atomic hydrogen spectroscopy Each spectral line of.^[3]

Figure 2

1923, American chemist Gilbert Lewis （ G.N.Lewis ）Developed Kossel's theory and proposed the electron pair theory of covalent bond^[1] 。 Lewis hypothesis: an electron from one atom and an electron from another atom in a molecule“ Electron pair ”The form of C forms chemical bonds between atoms. This was a hypothesis that ran counter to orthodox theory at that time, because Coulomb's law It shows that the two electrons are mutually exclusive, but Lewis's idea was soon accepted by the chemical community, leading to electron spin The proposition of opposite hypothesis.

In 1924, Louis de Broglie （ Louis de Broglie ）Propose Wave particle duality A mathematical model of the atom is established to describe the electron as a three-dimensional waveform. Mathematically, it is impossible to obtain the exact values of position and momentum at the same time.

In 1926, Schrodinger The wave equation of quantum mechanics was proposed, which can be directly used to explain the "formation" and "fracture" of chemical bonds, which became the initial beginning of quantum chemistry.

In 1927, Walter Heitler （ W.H.Heitler ）And F. London （ F.London ）Using quantum mechanics to deal with hydrogen molecule, we calculated the wave function , first use Quantum mechanical method Solve the covalent bond problem. Valence bond theory Born in the promotion of this method, their method of studying covalent bonds is called HL method.^[1]

In 1928, Enrico Fermi （ Enrica Fermi ）A new method based on Poisson distribution The single electron density model of atomic structure Question.^[4] After that, Douglas Hartree （ Douglas Rayner Hartree ）Application Iterative method , electronic Hamiltonian operator It is decomposed into the simple addition of several single electron Hamiltonian operators, and then the multi electron wave function of the system is expressed as the product of the single electron wave function. This model is improved, and the Hartley equation is proposed.^[5]

In 1930, Hartley's students Fokker （ Fock ）And John Slater （ John Clarke Slater ）Perfect Hartley equation, called Hartley Fokker equation （ HF ）。 In the early 1950s, Slater obtained the approximate wave function of HF: Hartley Fokker Slater equation（ HFS ）^[6] 。 In 1963, Hermann（ F.Hermann ）And Skillman（ S.Skillman ）Apply HFS to Ground state atom Function.^[7]

In 1950, Clemens Rotern （ C. C. J. Roothaan ）It is further proposed that molecular orbital The famous RHF equation was developed using the linear expansion of atomic orbitals of constituent molecules. In 1964, computer chemist Enrique Clementi（ E.Clementi ）A large number of RHF wave functions have been published,^[8] This equation and its subsequent improved version have become the main method to deal with quantum chemistry problems in modern times.

Figure 3

In 1929, Bert et al proposed Coordination field theory , first used to discuss the energy level splitting of transition metal ions in the crystal field, and later Molecular orbital theory It has developed into a modern coordination field theory. In 1930, American chemist Linus Pauling （ L.C.Pauling ）In the study of carbon Regular tetrahedron Configuration Time raised track Hybridization Theory It is believed that orbits with similar energy levels can hybridize when excited, forming new Degenerate orbit Its theoretical basis is electronic Wave particle duality , and wave can superposition Of. He calculated many kinds Hybrid orbit And obtained due to its outstanding contribution to valence bond theory Nobel Prize in Chemistry 。^[1]

In 1932, Friedrich Hund （ F.Hund ）The covalent bond is divided into σ bond 、 pi bond 、 Delta bond Three, further systematize the valence bond theory and organically combine it with the classical valence theory.^[1]

In the same year, American chemist Robert S.Mulliken （ Robert S.Mulliken ）The molecular orbital theory is proposed. It is believed that the electrons in the compound do not belong to a certain atom, but move in the whole molecule. His method is far from the classical chemistry, and the calculation is very tedious, which is not accepted by the chemical community for the time being. After Robert Milligan （ Robert A.Millikan ）、 Philip Leonard （ Philipp Lenard ）、 Erich Hucker （ Erich Hückel ）The improvement of others is gradually recognized in the chemical industry.^[1]

In 1940, Henry Hijvik （ H.Sidgwick ）And Thomas Powell（ Thomas A.Powell ）On the basis of summarizing the experimental facts, a simple theoretical model To predict the three-dimensional structure of simple molecules or ions. This theoretical model Ronald Gillespie （ R.J.Gillespie ）And Ronald Niholm （ R.S.Nyholm ）It was developed in the 1950s and named Valence shell electron pair repulsion theory , referred to as VSEPR. VSEPR vs Track hybridization Combining the theory, we can semi quantitatively speculate the molecular bonding mode and molecular structure.

In 1951, Kenichi Fukui propose frontier orbital It is believed that the molecular orbital with the highest energy in the molecule（ HOMO ）And the lowest energy molecular orbital not occupied by electrons（ LUMO ）It is the key to determine the chemical reaction of a system. Molecular orbitals of other energies have little influence on the chemical reaction and can be ignored temporarily. HOMO and LUMO are the so-called front track.

1965, American chemist Robert Woodward （ Rober B.Woodward ）With Hoffman's reference to Fukui's frontier track theory Conservation principle of molecular orbital symmetry 。 Molecular orbital theory has been developed.^[1]

Due to the rapid development of computer technology, and Monte Carlo Application of methods, quantum chemistry and Computer Chemistry With each passing day, a large number of excellent chemists were born during the period when the calculation of molecular structure became more accurate. It is estimated that there will be new breakthroughs in quantum chemistry in the middle of the 21st century.

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Saturation

During the formation of covalent bond, because each atom can provide unpaired Number of electrons It is certain that an unpaired electron of an atom can not pair with other electrons after pairing with unpaired electrons of other atoms, that is, the total number of covalent bonds formed by each atom is certain, which is the saturation of covalent bonds.^[9]

The saturation of covalent bond determines the quantitative relationship between various atoms when they form molecules^[9] , Yes Constant ratio law （ law of definite proportion ）Is one of the internal reasons.

directional

Except that the s orbital is spherical, other atomic orbitals have their fixed extension direction, so when the covalent bond is formed, the orbital overlap also has a fixed direction, and the covalent bond also has its own directional The direction of the covalent bond determines the configuration of the molecule.^[9]

The orientation of covalent bonds is affected by the direction of orbital extension.

chemical property

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The essence of chemical change is the breaking of old bonds and the formation of new bonds. In chemical reactions, there are two ways of breaking covalent bonds, which have an important impact on chemical reactions, especially in organic chemistry.

Homolysis and free radical reaction

Covalent bond occurs Homolysis When the bonding electrons are evenly divided into two atoms (clusters), the atoms (clusters) with single electrons produced by homogeneous splitting are called free radicals, which are expressed by "R ·". Free radicals are reactive and can participate in chemical reactions, Free radical reaction It is generally carried out under the action of light or heat.

Heterolysis and Ionic Reaction

Covalent bond occurrence Heteroschisis Generate positive anion , e.g. hydrogen chloride in water ionization Hydrogenation ion and Chloride ion 。 Heterolytic Carbocation And negative ions are Organic reaction The active species often participate in the reaction at the moment of generation, but can prove its existence.^[10]

Reaction Scale Initiated by Heterolysis of Covalent Bond Ionic reaction , which can be divided into two types

· Electrophilic reaction （ electrophilic reaction ）

· Nucleophilic reaction （ nucleophilic reaction ）

Ionic reactions are generally acid-base or Polarity Under the catalysis of substances.

theoretical model

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Lewis theory

Lewis theory, also known as "octet rule" and "electron pairing theory", was first proposed and has epoch-making significance Covalent Bond Theory It has no basis in quantum mechanics, but because it is easy to understand and can also explain the formation of most covalent bonds, it still appears in middle school textbooks.^[11]

The theory of shared electron pair has the following points:

1. When the outermost layer of an atom reaches 8 electrons, it is a stable structure. The number of outermost valence electrons of all atoms in the compound must be 8 (hydrogen is 2);

2. When covalent bonds are formed between atoms, the outermost layer can reach 8 (2) electron stable structure by sharing electrons.

The electron pairing idea of Lewis theory laid the foundation for the development of valence bond theory.^[12] It is worth noting that Lewis theory is not perfect, it can not explain the reason and essence of electron pairing; In addition, non conformance“ Octet Rule ”There are also many compounds, such as: Boron trifluoride (6 Electronic) Phosphorus pentachloride (10 electronic) Sulfur hexafluoride (12 electronic).

Valence bond theory

The valence bond theory is a covalent bond theory developed based on the Lewis theory of electron pairing. The valence bond theory extends the results of solving hydrogen molecular problems by applying quantum mechanics to others Covalent compound Successfully explained many structural problems of molecules.

Heitler -London law

Walter Heitler （ W.H.Heitler ）And F. London （ F.London ）In the process of using quantum mechanics to deal with hydrogen molecules, the relationship curve between molecular energy E and nuclear spacing R is obtained. It is found that if the spin directions of two hydrogen atoms are opposite, with the overlapping of orbits (addition of wave functions), a region with high probability density will appear, and hydrogen atoms will bond at the lowest nuclear spacing of system energy; If the two hydrogen atoms have the same spin direction, the subtractive wave function decreases monotonically, the system energy approaches E=0 infinitely, and there is no lowest point, and no bond can be formed. Therefore, the valence bond theory clarifies the intrinsic reason of electron pairing and the essence of covalent bond through the study of hydrogen molecule, and the valence bond theory was born in the promotion of HL.^[12]

Orbital hybrid theory

When the valence bond theory explains the distribution of atoms in molecules（ L.Pauling ）Proposed the track Hybridization Theory 。 The main theoretical points are

1. Different orbits with similar central atomic energy will hybridize under the influence of the outside world, forming new orbits, called hybrid atomic orbits, or hybrid orbits for short;

2. The angular distribution of hybrid orbitals is more concentrated than that of pure atomic orbitals, so the overlapping degree is greater, which is more conducive to bonding;

3. The number of atomic orbitals participating in hybridization is equal to the number of hybrid orbitals formed. Different types of hybrid orbitals have different spatial orientations.

Table 1
Hybrid type	Hybrid track angle	Spatial orientation
sp^ one	one hundred and eighty	Linear
sp^ two	one hundred and twenty	Plane equilateral triangle
sp^ three	one hundred and nine point two eight	Regular tetrahedron
sp^ three d (dsp^ three )	90 120	Trigonal bipyramid
sp^ three d^ two (d^ two sp^ three )	ninety	Octahedron

Note: This is the spatial orientation of hybrid orbital, not the structure of compound

In compounds, these orbitals may be Lone pair electron Or single electron filling, for example, N atom for sp ² There is a single electron in the hybrid NO2 molecule, and the spatial structure of NO2 is broken line (a vertex of an equilateral triangle is a single electron, and the electron is "invisible").

Mutual exclusion theory

The valence shell electron pair repulsion theory (VSEPR Theory) is a chemical model used to predict the morphology of a single covalent molecule. By calculating the sum of the valence shell electron number of the central atom coordination number To predict the geometry of molecules configuration Its theoretical points include:

1. In covalent molecules, the geometry of the electron pair arrangement around the central atom is mainly determined by the Valence electron layer The number of pairs of electrons in (including bonded pairs and lone pairs of electrons). The position of these electrons tends to separate as far as possible to minimize the repulsion force on each other^[13] ；

2. The mutual repulsion of electron pairs in the electron layer depends on the mutual angle between electron pairs and the bonding of electron pairs. Small distance angle, large repulsion force. The bonding electron pair is attracted by two atoms, Electronic cloud It is relatively tight, and the repulsion force to its adjacent electron pairs is less than that of only one Nucleus The repulsive force of the attracted lone pair of electrons to its adjacent pair of electrons. That is, the order of the repulsive force between electron pairs is: lone pair electron lone pair electron>lone pair electron bonding electron pair>bonding electron pair bonding electron pair;^[13]

3. Intramolecular double bond The triple bond is treated as a single bond;^[13]

Speculative molecular configuration

If the central atom is A, the other n coordination atoms are all represented by B, and m pairs of lone electrons are represented by E, then the substance can be represented as ABnEm. Let z=n+m, and both B and E are represented by Y, then the substance can be represented as AYz, where Y represents Central atom Z represents the number of pairs of electrons in the valence electron layer of the central atom. We can infer the molecular configuration according to the following formula:

N can be seen from the chemical formula

M=1/2 (number of valence electrons of central atom - total number of electrons provided by coordination atom ± number of ion charges)

z=n+m

Table 2
Z=n+m	two	three	four	five	six
structural style	Linear	Plane triangle	tetrahedron	Trigonal bipyramid	Octahedron

Note: For more detailed tables, see Wikipedia, VSEPR theory (extended reading)

Molecular orbital theory

Molecular orbital theory is more accurate than valence bond theory, and its theoretical points are

1. The electrons in the molecule do not belong to an atomic orbit, but belong to the whole molecule;^[13]

2. Molecular orbitals are linear combinations of atomic orbitals. The number of molecular orbitals is equal to the number of atomic orbitals that make up the molecular orbitals. Some of these orbitals have lower energy, become "bonding orbitals", and others have higher energy, become“ Antibonding orbit ”There is still some energy unchanged, called "non bond orbit";^[13]

3. When atomic orbitals are linearly combined, observe“ Symmetry matching principle ”、“ Principle of energy similarity ”、“ Maximum overlap principle ”；^[13]

4. When electrons are arranged in molecular orbital, observe“ Principle of minimum energy ”、“ Pauli exclusion principle ”、“ Hunt rule ”；^[13]

Molecular orbital theory can explain some phenomena that cannot be explained by valence bond theory, such as Oxygen molecule Paramagnetism.

The number of outer electrons of oxygen atom is 6. Four of these six electrons form two pairs, and the other two exist alone.

Figure 4 Each oxygen atom has six outer electrons

These two separate electrons combine with the corresponding separate electrons in another atom to form two new shared electron pairs, thus reaching the state of electron saturation.

Figure 5 Model of oxygen molecule O2

It should be noted that the oxygen molecule model described here is a simplified model, and the actual oxygen molecule is much more complex than that described here, because the six outer atoms are distributed in different orbits, so they cannot form such simple electron pairs. The actual oxygen molecule has three pairs of shared electrons and two separate electrons.

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The covalent bond can perform different classification Each classification includes all covalent bonds (only the classification angle is different).

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Fig. 6 σ bond

σ bond （ sigma bond ）

The covalent bond formed by the overlap of two atomic orbits along the orbital symmetry axis, which increases the probability of electrons appearing between nuclei, is called σ bond, which can be abbreviated as "head to head" (see Figure 6).^[9] σ bond belongs to Local key It can be either a general covalent bond or a coordination covalent bond. Generally, single bonds are σ bonds. Atomic orbital happen Hybridization The covalent bond formed after is also a σ bond. Since the σ bond is along the orbit Axis of symmetry The direction is formed, and the degree of overlap between orbits is large. Therefore, the bond energy of σ bond is usually large, and it is not easy to break. Moreover, since the effective overlap is only once, only one σ bond can be formed between two atoms at most.

pi bond （ pi bond ）

Figure 7 π bond^[14]

Unhybridized p orbitals of bonding atoms form covalent bonds through parallel and side overlapping, called π bonds, which can be abbreviated as "side by side" (see Figure 7).^[9] π bond is different from σ bond Bonding orbital Must be unpaired p orbitals. The π bond has different properties. It has two centers and two electrons. It can also be a localized bond Conjugate π bond and Feedback π key 。 Two atoms can form up to two π bonds, for example, Carbon carbon double bond There is a σ bond, a π bond, and carbon carbon Triple bond There is one σ bond and two π bonds.

Figure 8 Large π bond in benzene molecule

In π bond π electron It can absorb ultraviolet rays and be excited. Therefore, compounds containing π bonds have the function of resisting ultraviolet rays. Sunscreen uses this principle to protect people from ultraviolet rays.^[9] Conjugated π bonds have special stability, such as benzene ring There is a large π bond with 6 centers and 6 electrons in the, which shows aromaticity and is not easy to add and oxidation reaction , but easy to happen Electrophilic substitution Compounds with bond type similar to benzene ring include Heterocyclic compound 、 Polycyclic hydrocarbon And others Hydrocarbons , chemist Erich Hucker By molecular orbital calculation Cycloolefin Aromatic Hull rule (also known as 4n+2 rule ）, other common nonbenzenoid hydrocarbon include Azulene 、 [18] Annulene Etc; Each layer of graphite has an infinite π bond, in which electrons can move freely, similar to Metal bond This is also the reason why graphite can conduct electricity transversely.^[9]

Delta bond （ delta bond ）

Fig. 9 δ bond

The covalent bond formed by the quadruple overlap of two d orbitals is called delta bond, which can be abbreviated as "face-to-face" (see Figure 9).

The delta bond has only two nodal planes (the plane where the density of the electron cloud is zero). Seen from the bond axis, the orbital symmetry of the δ bond is no different from that of the d orbital, and the Greek letter δ is also derived from the d orbital.

The delta bond often occurs in Organometallic compound Medium, especially ruthenium molybdenum and rhenium Compounds formed. What is commonly said“ Quadruple bond ”It refers to one σ bond, two π bonds and one δ bond.

The above three chemical bonds can form various bond types after combination, for example, one σ bond and two π bonds can form a triple bond, but there is evidence that the number of covalent bonds between diatoms cannot exceed six at most.^[15]

Covalent bonds are the overlap of electronic clouds, so the most essential classification of covalent bonds is their overlap. Now there are three overlapping modes known, namely:

σ bond pi bond Delta bond stay Organic compound In general, covalent bonds are divided into Single bond 、 double bond as well as Triple bond 。 The single bond is a σ bond; Double bond and Triple bond They all contain one σ bond, and the other one or two are π bonds.

but inorganic compound Do not use this method. The reason is that conjugate System（ Delocalization π bond ）It is difficult to determine the number of electron pairs shared between two atoms, so it is often taken as Average key level , as Bond energy Rough criteria for.

Press the keying process

1. General covalent bond

The general covalent bond is sometimes called "normal covalent bond", which is a concept used to distinguish it from "coordination covalent bond". It refers to the covalent bond formed when two atoms each provide an unpaired electron when bonding.

2、 Coordination covalent bond （ coordinate covalent bond ）

Coordination covalent bond“ Coordination bond ”It refers to the covalent bond formed when all bonding electrons of two atoms are provided by one atom“ Ligands (referred to as ligand for short). Common ligands are: ammonia (nitrogen atom) carbon monoxide （ carbon atom ）, cyanogen ion (carbon atom), water（ Oxygen atom ）、 Hydroxyl radical (oxygen atom); Receptors are diverse: hydrogen ions, represented by boron trifluoride (boron atom) Electron deficient compound , and a large number of Transition metal elements 。 yes Coordination compound The research of, Coordination chemistry 。

Coordination bond is a special kind of covalent bond, which is characterized by that the shared pair of electrons come from the same atom. The condition for forming coordination bond is that an atom has Lone electron pair And the other atom has an empty orbital.

Central ion : In the complex, the party providing the empty orbit is called the central ion

ligand : In the complex, the party that provides the lone pair electrons is called ligand

Table 3
classification	Chemical bond
covalent bond	σ bond: three center two electron bond (banana bond) · three center four electron bond (hydrogen bond, double hydrogen bond, hydrogen grasping bond) · four center two electron bond
	π bond: feedback π bond· conjugate · Hyperconjugate effect ·Directivity
	δ bond: Quadruple bond · Quintuple bond · Sextuple key
hydrogen bond	Double hydrogen bond · Dihydrogen complex ·Low barrier hydrogen bond· Symmetric hydrogen bond · Hydrophilicity
Noncovalent bond	Van der Waals ·Mechanical combination · embedding· Halogen bond · Metallophilic action ·Overlap· Entropy force · Polarity
other	Intramolecular force and intermolecular force· Coordination bond · Occlusal degree ·Ionic bond · metal bond· Bonding · Antibonding · Disulfide bond · Peptide bond · Phosphodiester bond

2.1 Similarities and differences between coordination covalent bond and general covalent bond

The difference between coordination covalent bond and general covalent bond is only reflected in the bonding process Key Parameters Is the same, for example, nitrogen of ammonium ion hydrogen bond There are three general covalent bonds and one coordination covalent bond, but these four bonds are completely equivalent, and the ammonium ion is also a fully symmetric tetrahedron. In writing, the symbol "-" is generally used for covalent bond; The coordination covalent bond uses the symbol "→" arrow to point from the ligand to the receptor.

Electron bias

1、 Polar covalent bond （ polar bond ）

Figure 10 Polar key, labeled&

In the compound molecule, the covalent bond formed by different atoms, because the two atoms have different ability to attract electrons, the electron cloud is biased towards the atom with stronger ability to attract electrons, so the atom with weaker ability to attract electrons is relatively significantly positive. Such covalent bond is called polar covalent bond Polar bond 。 When forming covalent bonds, polar bonds can be divided into "strong polar bonds" and "weak polar bonds" due to the different deviation degrees of the electronic clouds, but usually the bonding between two different atoms is polar bonds.^[9] Covalent Bond polarity The bond moment can be used for judgment. The polarity of a covalent molecule is equal to all covalent bonds in the molecule dipole moment The vector sum of, therefore, the molecule composed of polar covalent bond can be Polar molecule （ Hydrogen chloride ）It can also be Nonpolar molecule （ carbon dioxide ）。

2、 nonpolar covalent bond （ non-polar bond ）

from Homogeneous element The covalent bond formed between the atoms of N is called non-polar covalent bond. Homoatomic attraction sharing Electron pair The bonding electron pairs are evenly distributed between the two nuclei without bias to any atom, and the bonding atoms are not sensitive to electricity.^[9] Non polar covalent bonds exist in Simple substance It also exists in some compounds and is completely composed of Nonpolar bond The molecules formed must be nonpolar molecules (but some nonpolar molecules contain polar bonds).

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Key Parameters

1、 Key length （ bond length ）

Bond length refers to the distance between the equilibrium nuclei of two bonding atoms. It is the basic configuration parameter to understand the molecular structure and also to understand Chemical bond The parameters of strength, strength and property. For the chemical bond composed of the same A and B atoms, the bond length value is small and the bond strength is strong; The number of keys is large, and the value of key length is small. In actual molecules, due to Conjugate effect 、 Spatial hindrance effect And adjacent Group Electronegativity There are some differences in the bond length of the same chemical bond. The bond length is mainly measured by Molecular spectrum and Thermochemistry Means.^[12] The following table shows the bond lengths of common covalent bonds（ pm ）Data is from Chemistry - Structure and Properties of Substances (Elective) (2007).^[9]

Table 4
covalent bond	Key length	covalent bond	Key length	covalent bond	Key length	covalent bond	Key length
H-H	seventy-four	H-F	ninety-two	H-Cl	one hundred and twenty-seven	H-Br	one hundred and forty-one
H-I	one hundred and sixty-one	C-H	one hundred and nine	C-C	one hundred and fifty-four	C-Si	one hundred and eighty-six
C-N	one hundred and forty-seven	C-O	one hundred and forty-three	C-P	one hundred and eighty-seven	C-S	one hundred and eighty-one
C-F	one hundred and thirty-three	C-Cl	one hundred and seventy-seven	C-Br	one hundred and ninety-four	C-I	two hundred and thirteen
N-H	one hundred and one	N-N	one hundred and forty-six	N-P	one hundred and seventy-seven	N-O	one hundred and forty-four
N-F	one hundred and thirty-nine	N-Cl	one hundred and ninety-one	N-Br	two hundred and fourteen	N-I	two hundred and twenty-two
O-H	one hundred and one	O-P	one hundred and sixty	O-S	one hundred and fifty-one	O-F	one hundred and forty-two
O-Cl	one hundred and sixty-four	O-Br	one hundred and seventy-two	O-I	one hundred and ninety-four	Si-H	one hundred and forty-eight
Si-Si	two hundred and thirty-four	Si-O	one hundred and sixty-one	SI-S	two hundred and ten	Si-F	one hundred and fifty-six
Si-Cl	two hundred and four	Si-Br	two hundred and sixteen	Si-I	two hundred and forty	P-H	one hundred and forty-two
P-Si	two hundred and twenty-seven	P-P	two hundred and twenty-one	P-F	one hundred and fifty-six	P-Br	two hundred and twenty-two
P-I	two hundred and forty-three	S-H	one hundred and thirty-four	S-F	one hundred and fifty-eight	S-Cl	two hundred and one
S-Br	two hundred and twenty-five	S-I	two hundred and thirty-four	F-F	one hundred and forty-three	Cl-Cl	one hundred and ninety-nine
Br-Br	two hundred and twenty-eight	I-I	two hundred and sixty-six	C=C	one hundred and thirty-four	C=N	one hundred and twenty-seven
C=O	one hundred and twenty-three	N=N	one hundred and twenty-two	N=O	one hundred and twenty	O=O	one hundred and twenty-one
C≡C	one hundred and twenty-one	C≡N	one hundred and fifteen	N≡N	one hundred and ten

2、 Bond energy （ bond energy ）

Usually refers to Standard status lower Gaseous The molecules are disassembled into gas atom The average value of the energy required for each key. yes Diatomic molecule For example, the bond energy is the bond energy Dissociation energy 。 The bond energy is approximately equal to the bond enthalpy Atomization energy Equal to the sum of all bond energies.

The following table shows the bond energies of common covalent bonds（ kJ/mol ）Data is from Chemistry - Structure and Properties of Substances (Elective) (2007).^[9]

Table 5
covalent bond	Bond energy	covalent bond	Bond energy	covalent bond	Bond energy	covalent bond	Bond energy
H-H	four hundred and thirty-six	H-F	five hundred and sixty-five	H-Cl	four hundred and thirty-one	H-Br	three hundred and sixty-three
H-I	two hundred and ninety-seven	C-H	four hundred and thirteen	C-C	three hundred and forty-seven	C-Si	three hundred and one
C-N	three hundred and five	C-O	three hundred and fifty-eight	C-P	two hundred and sixty-four	C-S	two hundred and fifty-nine
C-F	four hundred and fifty-three	C-Cl	three hundred and thirty-nine	C-Br	two hundred and seventy-six	C-I	two hundred and sixteen
N-H	three hundred and ninety-one	N-N	one hundred and sixty	N-P	two hundred and nine	N-O	two hundred and one
N-F	two hundred and seventy-two	N-Cl	two hundred	N-Br	two hundred and forty-three	N-I	one hundred and fifty-nine
O-H	four hundred and sixty-seven	O-P	three hundred and fifty-one	O-S	two hundred and sixty-five	O-F	one hundred and ninety
O-Cl	two hundred and three	O-Br	two hundred and thirty-seven point six	O-I	Unknown	Si-H	three hundred and twenty-three
Si-Si	two hundred and twenty-six	Si-O	three hundred and sixty-eight	Si-S	two hundred and twenty-six	Si-F	five hundred and sixty-five
Si-Cl	three hundred and eighty-one	Si-Br	two hundred and thirteen	Si-I	Unknown	P-H	three hundred and twenty
P-Si	three hundred and ten	P-P	two hundred	P-F	four hundred and ninety	P-Br	two hundred and seventy-two
P-I	one hundred and eighty-four	S-H	three hundred and forty-seven	S-F	three hundred and twenty-seven	S-Cl	two hundred and seventy-one
S-Br	two hundred and eighteen	S-I	one hundred and seventy	F-F	one hundred and fifty-nine	Cl-Cl	two hundred and forty-three
Br-Br	one hundred and ninety-three	I-I	one hundred and fifty-one	C=C	six hundred and fourteen	C=N	six hundred and fifteen
C=O	745(799)	N=N	four hundred and eighteen	N=O	six hundred and seven	O=O	four hundred and ninety-eight
C≡C	eight hundred and thirty-nine	C≡N	eight hundred and ninety-one	N≡N	nine hundred and forty-five

3、 Bond angle （ bond angle ）

The bond angle is the included angle of two covalent bonds. Because of the directionality of covalent bonds, the bond angle of covalent compounds is certain, but compounds with similar compositions may not have the same bond angle, and the lone pair of electrons have larger bonding electrons exclude Function, can cause the bond angle to become smaller.

4、 Key level （ bond order ）

Bond order is a concept proposed by molecular orbitals, which is defined as half of the difference between bonding electrons and anti bonding electrons. Bond order can describe the stability of covalent bonds. The larger the bond order, the more stable the covalent bond.

5、 Bond dipole moment （ bond dipole moment ）

key Dipole distance It is called "bond moment" for short. The concept is similar to that of moment, which can describe the polarity of covalent bond. Bond moment Is defined as: μ =q·l

Where μ Is the bond moment (C · m), l is the bond length, and q is Charge quantity

The bond moment is a vector Electronegativity The weak end points to the strong electronegativity end, that is, from positive to negative. Bond moments can also be measured experimentally^[12]

Molecular model

Compared with the description of covalent bonds by bond parameters, the description of various models is more intuitive. The following table shows the colors and corresponding elements commonly used in molecular models.

Table 6
element	oxygen	carbon	nitrogen	sulfur	hydrogen	iodine	fluorine	chlorine	bromine
Color	red	ash	blue	yellow	white	purple	Yellow green	green	orange

Note: The above table only shows the common elements and corresponding colors, which are different from the actual situation.

Baton model （ Ball-and-stick models ）

Figure 11 Ball stick model (left) and scale filled model of methane

The stick model, also known as the "space filling model", is used to represent chemical molecules three-dimensional space Molecular model of distribution. In the stick model, "stick" represents covalent bond, and "ball" represents bonding atom. The ball stick model can represent the bonding angle of molecules and the radius of bonding atoms.

Scale filled model （ Space-filling models ）

The proportional filling model is similar to the stick model, which is used to represent the three-dimensional spatial distribution of molecules. It is a further development of the ball and stick model, which can show a more realistic molecular shape. However, it is difficult to see the bond angle of the compound from the model.

Scope of existence

1. Non metallic simple substance

2. Covalent compounds

3. Some ionic compounds

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