Chemical bond: characteristics, how they are formed, types - science - 2023
- Definition of chemical bond
- How are chemical bonds formed?
- Homonuclear compounds A-A
- Heteronuclear compounds A-B
- Types of chemical bonds
- -Covalent bond
- Simple link
- Double link
- Triple bond
- Non-polar bond
- Polar bonds
- Dative or coordination links
- -Ionic bond
- Metallic bond
- Examples of links
- Importance of the chemical bond
The Chemical bond It is the force that manages to hold together the atoms that make up matter. Each type of matter has a characteristic chemical bond, which consists of the participation of one or more electrons. Thus, the forces that bind atoms in gases are different, for example, from metals.
All the elements of the periodic table (with the exception of helium and the light noble gases) can form chemical bonds with each other. However, the nature of these is modified depending on which elements the electrons that form them come from. An essential parameter to explain the type of bonds is electronegativity.
The difference in electronegativity (ΔE) between two atoms defines not only the type of chemical bond, but also the physicochemical properties of the compound. The salts are characterized by having ionic bonds (high ΔE), and many of the organic compounds, such as vitamin B12 (top image), covalent bonds (low ΔE).
In the higher molecular structure, each of the lines represents a covalent bond. The wedges indicate that the link emerges from the plane (towards the reader), and the underlined ones behind the plane (away from the reader). Note that there are double bonds (=) and a cobalt atom coordinated with five nitrogen atoms and an R side chain.
But why do such chemical bonds form? The answer lies in the energy stability of the participating atoms and electrons. This stability must balance the electrostatic repulsions experienced between electron clouds and nuclei, and the attraction exerted by a nucleus on the electrons of the neighboring atom.
Definition of chemical bond
Many authors have given definitions of the chemical bond. Of all of them the most important was that of the physicochemist G. N. Lewis, who defined the chemical bond as the participation of a pair of electrons between two atoms. If atoms A · and · B can contribute a single electron, then the single bond A: B or A – B will form between them.
Before bond formation, both A and B are separated by an indefinite distance, but in bonding there is now a force holding them together in the diatomic compound AB and a bond distance (or length).
What characteristics does this force have that holds the atoms together? These depend more on the type of link between A and B than on their electronic structures. For example, link A – B is directional. What does it mean? That the force exerted by the union of the pair of electrons can be represented on an axis (as if it were a cylinder).
Also, this bond requires energy to break. This amount of energy can be expressed in the units of kJ / mol or cal / mol. Once enough energy has been applied to compound AB (by heat, for example), it will dissociate into the original A · and · B atoms.
The more stable the bond, the more energy it takes to separate the bonded atoms.
On the other hand, if the bond in compound AB were ionic, A+B–, then it would be a non-directional force. Why? Because+ exerts an attractive force on B– (and vice versa) that depends more on the distance that separates both ions in space than on their relative location.
This field of attraction and repulsion brings together other ions to form what is known as the crystal lattice (top image: cation A+ lies surrounded by four anions B–, and these four cations A+ and so on).
How are chemical bonds formed?
Homonuclear compounds A-A
In order for a pair of electrons to form a bond there are many aspects that must be considered first. The nuclei, say those of A, have protons and are therefore positive. When two A atoms are very far apart, that is, at a large internuclear distance (top image), they do not experience any attraction.
As the two A atoms approach their nuclei, they attract the electron cloud of the neighboring atom (the purple circle). This is the attractive force (A on the neighboring purple circle). However, the two nuclei of A repel each other because they are positive, and this force increases the potential energy of the bond (vertical axis).
There is an internuclear distance in which the potential energy reaches a minimum; that is, both the attractive and repulsive forces (the two A atoms in the lower part of the image) are balanced.
If this distance decreases after this point, the bond will cause the two nuclei to repel each other with great force, destabilizing compound A-A.
So for the bond to form there must be an energetically adequate internuclear distance; Furthermore, the atomic orbitals must overlap correctly for the electrons to bond.
Heteronuclear compounds A-B
What if instead of two atoms of A, one of A and the other of B were joined? In this case the upper graph would change because one of the atoms would have more protons than the other, and the electron clouds would have different sizes.
As the A – B bond is formed at the appropriate internuclear distance, the electron pair will be found mainly in the vicinity of the most electronegative atom. This is the case with all heteronuclear chemical compounds, which constitute the vast majority of those that are known (and will be known).
Although not mentioned in depth, there are numerous variables that directly influence how atoms approach and chemical bonds are formed; some are thermodynamic (is the reaction spontaneous?), electronic (how full or empty are the orbitals of the atoms) and others kinetic.
Types of chemical bonds
Links have a series of characteristics that distinguish them from each other. Several of them can be framed in three main classifications: covalent, ionic or metallic.
Although there are compounds whose bonds belong to a single type, many actually consist of a mixture of characters of each. This fact is due to the difference in electronegativity between the atoms that form the bonds. Thus, some compounds may be covalent, but have a certain ionic character in their bonds.
Likewise, the type of bond, the structure and the molecular mass are key factors that define the macroscopic properties of the matter (brightness, hardness, solubility, melting point, etc.).
Covalent bonds are those that have been explained so far. In them, two orbitals (one electron in each) must overlap with the nuclei separated by an appropriate internuclear distance.
According to the molecular orbital theory (TOM), if the overlap of the orbitals is frontal, a sigma σ bond will form (which is also called a simple or simple bond). Whereas if the orbitals are formed by lateral and perpendicular overlaps with respect to the internuclear axis, we will have π bonds (double and triple):
The σ bond, as can be seen in the image, is formed along the internuclear axis. Although not shown, A and B may have other bonds, and therefore their own chemical environments (different parts of the molecular structure). This type of link is characterized by its rotational power (green cylinder) and by being the strongest of all.
For example, the single bond in the hydrogen molecule can rotate on the internuclear axis (H – H). Similarly, a hypothetical CA – AB molecule can.
Links C – A, A – A, and A – B rotate; but if C or B are atoms or a group of bulky atoms, the A – A rotation is sterically impeded (because C and B would collide).
Single bonds are found in practically all molecules. Its atoms can have any chemical hybridization as long as the overlap of their orbitals is frontal. Going back to the structure of vitamin B12, any single line (-) indicates a single link (for example, -CONH links2).
The double bond requires the atoms to be (usually) sp hybridized2. The pure p bond, perpendicular to the three hybrid sp orbitals2, forms the double bond, which appears as a grayish sheet.
Note that both the single bond (green cylinder) and the double bond (gray sheet) coexist at the same time. However, unlike single bonds, double bonds do not have the same freedom of rotation around the internuclear axis. This is because, to rotate, the link (or the foil) must break; process which needs energy.
Also, the bond A = B is more reactive than A – B. Its length is shorter and atoms A and B are at a shorter internuclear distance; therefore, there is greater repulsion between both nuclei. Breaking both the single and double bonds requires more energy than is needed to separate the atoms in the A – B molecule.
In the structure of vitamin B12 Several double bonds can be observed: C = O, P = O, and within aromatic rings.
The triple bond is even shorter than the double bond and its rotation is more energetically impeded. In it, two perpendicular π bonds are formed (the greyish and purple sheets), as well as a single bond.
Ordinarily, the chemical hybridization of the atoms of A and B must be sp: two sp orbitals 180º apart, and two pure p orbitals perpendicular to the first. Note that a triple bond looks like a paddle, but without rotational power. This bond can be represented simply as A≡B (N≡N, nitrogen molecule N2).
Of all the covalent bonds, this is the most reactive; but at the same time, the one that needs more energy for the complete separation of its atoms (· A: +: B ·). If vitamin B12 had a triple bond within its molecular structure, its pharmacological effect would change dramatically.
Six electrons participate in triple bonds; in doubles, four electrons; and in the simple or simple, two.
The formation of one or more of these covalent bonds depends on the electronic availability of the atoms; that is, how many electrons do their orbitals need to acquire one octet of valence.
A covalent bond consists of an equal sharing of a pair of electrons between two atoms. But this is strictly true only in the case where both atoms have equal electronegativities; that is, the same tendency to attract electron density from its surroundings into a compound.
Nonpolar bonds are characterized by a null electronegativity difference (ΔE≈0). This occurs in two situations: in a homonuclear compound (A2), or if the chemical environments on both sides of the bond are equivalent (H3C – CH3, ethane molecule).
Examples of nonpolar bonds are seen in the following compounds:
-Hydrogen (H – H)
-Oxygen (O = O)
-Fluorine (F – F)
-Chloro (Cl – Cl)
When there is a marked difference in electronegativity ΔE between both atoms, a dipole moment is formed along the bond axis: Aδ+–Bδ-. In the case of the heteronuclear compound AB, B is the most electronegative atom, and therefore, it has a higher electron density δ-; while A, the least electronegative, has a δ + charge deficiency.
For polar bonds to occur, two atoms with different electronegativenesses must join; and thus, form heteronuclear compounds. A – B resembles a magnet: it has a positive and a negative pole. This allows it to interact with other molecules through dipole-dipole forces, among which are hydrogen bonds.
Water has two polar covalent bonds, H – O – H, and its molecular geometry is angular, which increases its dipole moment. If its geometry were linear, the oceans would evaporate and the water would have a lower boiling point.
The fact that a compound has polar bonds, does not imply that it is polar. For example, carbon tetrachloride, CCl4, has four polar bonds C – Cl, but due to their tetrahedral arrangement, the dipole moment ends up being vectorially annulled.
Dative or coordination links
When an atom gives up a pair of electrons to form a covalent bond with another atom, then we speak of a dative or coordination bond. For example, having B: the available electron pair, and A (or A+), an electronic vacancy, the B: A link is formed.
In the structure of vitamin B12 the five nitrogen atoms are linked to the metal center of Co by this type of covalent bond. These nitrogens give up their free electron pair to the Co cation.3+, coordinating the metal with them (Co3+: N–)
Another example can be found in the protonation of an ammonia molecule to form ammonia:
H3N: + H+ => NH4+
Note that in both cases it is the nitrogen atom that contributes the electrons; therefore, the dative or coordination covalent bond occurs when an atom alone contributes the pair of electrons.
In the same way, the water molecule can be protonated to become the hydronium (or oxonium) cation:
H2O + H+ => H3OR+
Unlike the ammonium cation, hydronium still has a free electron pair (H3OR:+); however, it is very difficult for it to accept another proton to form the unstable hydronium dication, H4OR2+.
Pictured is a white hill of salt. The salts are characterized by having crystalline structures, that is to say, symmetrical and ordered; high melting and boiling points, high electrical conductivities when melting or dissolving, and also, its ions are strongly bound by electrostatic interactions.
These interactions make up what is known as the ionic bond. In the second image a cation A was shown+ surrounded by four anions B–, but this is a 2D representation. In three dimensions, A+ should have other anions B– forward and behind the plane, forming various structures.
Thus, A+ it can have six, eight, or even twelve neighbors. The number of neighbors surrounding an ion in a crystal is known as the coordination number (N.C). For each N.C a type of crystalline arrangement is associated, which in turn constitutes a solid phase of the salt.
The symmetrical and faceted crystals seen in the salts are due to the equilibrium established by the attraction interactions (A+ B–) and repulsion (A+ TO+, B– B–) electrostatic.
But why A + and B–, or Na+ and Cl–, do not form covalent bonds Na – Cl? Because the chlorine atom is much more electronegative than sodium metal, which is also characterized by very easily giving up its electrons. When these elements meet, they react exothermically to produce table salt:
2Na (s) + Cl2(g) => 2NaCl (s)
Two sodium atoms give up their single valence electron (Na) to the diatomic molecule of Cl2, in order to form the anions Cl–.
The interactions between sodium cations and chloride anions, although they represent a weaker bond than covalents, are capable of keeping them strongly united in the solid; and this fact is reflected in the high melting point of the salt (801ºC).
The last of the types of chemical bond is metallic. This can be found on any metal or alloy part. It is characterized by being special and different from the others, due to the fact that electrons do not pass from one atom to another, but travel, like a sea, the crystal of metals.
Thus, metallic atoms, to say copper, intermingle their valence orbitals with each other to form conduction bands; through which electrons (s, p, d or f) pass around the atoms and hold them tightly together.
Depending on the number of electrons that pass through the metallic crystal, the orbitals provided for the bands, and the packing of its atoms, the metal can be soft (like alkali metals), hard, shiny, or a good conductor of electricity and hot.
The force that holds together the atoms of metals, such as those that make up the little man in the image and his laptop, is greater than that of salts.
This can be verified experimentally because the crystals of the salts can be divided in several halves before a mechanical force; whereas a metallic piece (composed of very small crystals) deforms.
Examples of links
The following four compounds encompass the types of chemical bonds explained:
-Sodium fluoride, NaF (Na+F–): ionic.
-Sodium, Na: metallic.
-Fluorine, F2 (F – F): nonpolar covalent, due to the fact that there is a null ΔE between both atoms because they are identical.
-Hydrogen fluoride, HF (H – F): polar covalent, since in this compound fluorine is more electronegative than hydrogen.
There are compounds, such as vitamin B12, which has both polar and ionic covalent bonds (in the negative charge of its phosphate group -PO4–-). In some complex structures, such as that of metallic clusters, all these types of links can even coexist.
Matter offers in all its manifestations examples of chemical bonds. From the stone at the bottom of a pond and the water that surrounds it, to the toads that croak at its edges.
While the bonds may be simple, the number and spatial arrangement of the atoms in the molecular structure make way for a rich diversity of compounds.
Importance of the chemical bond
What is the importance of the chemical bond? The incalculable number of consequences that the absence of the chemical bond would unleash highlights its enormous importance in nature:
-Without it, colors would not exist, since its electrons would not absorb electromagnetic radiation. The dust and ice particles present in the atmosphere would disappear, and therefore the blue color of the sky would turn dark.
-Carbon could not form its endless chains, from which billions of organic and biological compounds derive.
- Proteins could not even be defined in their constituent amino acids. The sugars and fats would disappear, as well as any carbon compounds in living organisms.
-The Earth would be left without an atmosphere, because in the absence of chemical bonds in its gases, there would be no force to hold them together. Nor would there be the slightest intermolecular interaction between them.
-Mountains might disappear, because their rocks and minerals, although heavy, could not contain their atoms packed inside their crystalline or amorphous structures.
-The world would be made up of solitary atoms incapable of forming solid or liquid substances. This would also result in the disappearance of all transformation of matter; that is, there would be no chemical reaction. Just fleeting gases everywhere.
- Harry B. Gray. (1965). Electrons and Chemical Bonding. W.A. BENJAMIN, INC. P 36-39.
- Whitten, Davis, Peck & Stanley. Chemistry. (8th ed.). CENGAGE Learning, p 233, 251, 278, 279.
- Nave R. (2016). Chemical Bonding. Recovered from: hyperphysics.phy-astr.gsu.edu
- Chemical Bond Types. (October 3, 2006). Taken from: dwb4.unl.edu
- Formation of chemical bonds: The role of electrons. [PDF]. Recovered from: cod.edu
- CK-12 Foundation. (s.f.). Energy and Covalent Bond Formation. Recovered from: chem.libretexts.org
- Quimitube. (2012). Coordinate or dative covalent bond. Recovered from: quimitube.com