Discovery of the electron

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Before the discovery of the electron, an understanding of the nature of a chemical bond was not possible. It is true that the idea of the valence existed as early as 1852, and some appreciation of molecular geometry existed shortly thereafter. In this respect, the tetrahedral structure of the carbon atom was recognized by vannt Hoff and by le Bel, and the stereochemistry of complex ions was deduced by Werner. For such structures to exist, it was apparent that some type of bonding force must be prevent. For the lack of something better, the chemical bond was represent by a straight line b/w the symbols of the bonded atoms . This indicates the existence of a bond, but it, of cource, fails to give any description of the nature of the bond.

Even before the discovery of the electron, the independent existence of ions was

According to Paulling, a chemical bond exist b/w two atoms when the bonding force b/w them is of such strength as to lead to an aggregate of sufficient stability to warrant their consideration as an independent molecular species. Althrough it would appear that this definition permits some freedom of choices, we ordinarily find it convenint to consider five types of chemical bonds. These are ionic bonds, covalent bonds, metallic bonds,hydrogen bonds, and vanderwaals forces. All of these are important, and the first these are quite strong. However , in spite of its importance, we will not consider the metallic bond.

Since opposite charges attract via a basic electromagnetic force, the negatively-charged electrons orbiting the nucleus and the positively-charged protons in the nucleus attract each other. Also, an electron positioned between two nuclei will be attracted to both of them. Thus, the most stable configuration of nuclei and electrons is one in which the electrons spend more time between nuclei, than anywhere else in space. These electrons cause the nuclei to be attracted to each other, and this attraction results in the bond.

Bonds in chemical formula

The 3-dimensionality of atoms and molecules makes it difficult to use a single technique for indicating orbitals and bonds. In molecular formulae the chemical bonds (binding orbitals) between atoms are indicated by various different methods according to the type of discussion. Sometimes, they are completely neglected. For example, in organic chemistry chemists are sometimes concerned only with the functional groups of the molecule. Thus, the molecular formula of ethanol (a compound in alcoholic beverages) may be written in a paper in conformational, 3-dimensional, full 2-dimensional (indicating every bond with no 3-dimensional directions), compressed 2-dimensional (CH3-CH2-OH), separating the functional group from another part of the molecule (C2H5OH), or by its atomic constituents (C2H6O), according to what is discussed. Sometimes, even the non-bonding valence shell electrons (with the 2-dimensional approximate directions) are marked, i.e. for elemental carbon .'C'. Some chemists may also mark the respective orbitals, i.e. the hypothetical ethene−4 anion (\/C=C/\ −4) indicating the possibility of bond formation.

Strong chemical bonds

Typical bond lengths in pm
and bond energies in kJ/mol.
Bond lengths can be converted to Å
by division by 100 (1 Å = 100 pm).
Data taken from [1].




H — Hydrogen













C — Carbon

























N — Nitrogen










O — Oxygen







F, Cl, Br, I — Halogens



















Strong chemical bonds are the intramolecular forces which hold atoms together in molecules. A strong chemical bond is formed from the transfer or sharing of electrons between atomic centers and relies on the electrostatic attraction between the protons in nuclei and the electrons in the orbitals. Although these bonds typically involve the transfer of integer numbers of electrons (this is the bond order), some systems can have intermediate numbers. An example of this is the organic molecule benzene, where the bond order is 1.5 for each carbon atom.

The types of strong bond differ due to the difference in electronegativity of the constituent elements. A large difference in electronegativity leads to more polar (ionic) character in the bond.

Types of Chemical Bonding

There are several ways in which atoms can combine or chemically bond together to form a molecule. The most common type is covalent bonding, where the atoms share pairs of outer shell or valence electrons. Covalent bonding may be single or multiple, depending on the number of pairs the atoms share. Ionic bonding is another common way atoms combine, where one atom passes its electron to the other element, creating positive (+) and negative (-) ions




From a mathmetical standaredpoint, the simplest type of chemical bond is one that can be considered strictly electrostatic in a character . Such a treatment has proved successful for the alkali halides where the bonding occurs b/w the cation of a highly electropositive atom and the anion of a highly electronegative atom. Althrough it is possible to consider a bond to be partially covalent and partially ionic, the lexrent of ionic character in the bond is dependent on the difference in electronegativity b/w the combining atoms. In the instance of alkali halides, it would be safe to cinsider the bond to be almost exclusively ionic. However, the test of this postulate will actually rest on the success we have in qualitatively evaluating varios properties of the resultant compounds. In general, we can define a bond as purely ionic in terms of the success of the electrostatic model.

It is electrostatic force of attraction b/w two opposite charged ions.

Na+(g) + cl-(g) -> NaCl(s)

Size of cation is always smaller than its neutral parent atom because effective nuclear charge increases. Electrovalent bonds are more common in inorganic compoundes. The hylides, halides, oxides, sulphides, nitrides and carbides of alkai metals(IA) and alkaline earth metal(IIA) are ionic compounds.


  1. Electronegativity difference mest be large b/w two combining elements.
  2. I.E of electro-positive elements should be low.
  3. E.A of electro-negative should be high.
  4. Lattice energy of ionic crystal must be high.

Sodium chloride

Salt or Sodium Chloride (NaCl) is a good example of a ionic bonding. Sodium (Na) has 1 valance electron and Chlorine (Cl) has 7 electrons in its outer orbit. If Sodium lost its valance electron, its next shell will be full. But that would also make Sodium a positive ion. If Chlorine gained 1 valance electron, its shell would be full with a maximum of 8 electrons, and it would then be a negative ion.

Thus Sodium Chloride (NaCl) is a bonding of the Na+ ion and the Cl- ion.

Sodium lets Chlorine use its valance electron

In its solid form as table salt, the Na+ and the Cl- ions are held in place in a crystalline lattice. When dissolved in water, the ions freely roam about the solution.

Note that the combination of these two elements can result in a violent reaction, giving off heat and perhaps even an explosion. Seldom is Na directly combined with Cl to form NaCl. Usually the combination is done indirectly with other compounds or in a water solution. But the fact that the bonding process gives off energy means that the molecule is fairly stable and not easy to separate.

One- and three-electron bonds

Bonds with one or three electrons can be found in radical species, which have an odd number of electrons. The simplest example of a 1-electron bond is found in the hydrogen molecular cation, H2+. One-electron bonds often have about half the bond energy of a 2-electron bond, and are therefore called "half bonds". However, there are exceptions: in the case of dilithium, the bond is actually stronger for the 1-electron Li2+ than for the 2-electron Li2. This exception can be explained in terms of hybridization and inner-shell effects.[5]

The simplest example of three-electron bonding can be found in the helium dimer cation, He2+, and can also be considered a "half bond" because, in molecular orbital terms, the third electron is in an anti-bonding orbital which cancels out half of the bond formed by the other two electrons. Another example of a molecule containing a 3-electron bond, in addition to two 2-electron bonds, is nitric oxide, NO. The oxygen molecule, O2 can also be regarded as having two 3-electron bonds and one 2-electron bond, which accounts for its paramagnetism and its formal bond order of 2.[6]

Molecules with odd-electron bonds are usually highly reactive. These types of bond are only stable between atoms with similar electronegativities.[6]

Bent bonds

Bent bonds, also known as banana bonds, are bonds in strained or otherwise sterically hindered molecules whose binding orbitals are forced into a banana-like form. Bent bonds are often more susceptible to reactions than ordinary bonds.

3c-2e and 3c-4e bonds

In three-center two-electron bonds ("3c-2e") three atoms share two electrons in bonding. This type of bonding occurs in electron deficient compounds like diborane. Each such bond (2 per molecule in diborane) contains a pair of electrons which connect the boron atoms to each other in a banana shape, with a proton (nucleus of a hydrogen atom) in the middle of the bond, sharing electrons with both boron atoms.

Three-center four-electron bonds ("3c-4e") also exist which explain the bonding in hypervalent molecules. In certain cluster compounds, so-called four-center two-electron bonds also have been postulated.

In certain conjugated π (pi) systems, such as benzene and other aromatic compounds (see below), and in conjugated network solids such as graphite, the electrons in the conjugated system of π-bonds are spread over as many nuclear centers as exist in the molecule or the network.

Aromatic bond

In organic chemistry, certain configurations of electrons and orbitals infer extra stability to a molecule. This occurs when π orbitals overlap and combine with others on different atomic centres, forming a long range bond. For a molecule to be aromatic, it must obey Hückel's rule, where the number of π electrons fit the formula 4n + 2, where n is an integer. The bonds involved in the aromaticity are all planar.

In benzene, the prototypical aromatic compound, 18 (n = 4) bonding electrons bind 6 carbon atoms together to form a planar ring structure. The bond "order" (average number of bonds) between the different carbon atoms may be said to be (18/6)/2=1.5, but in this case the bonds are all identical from the chemical point of view. They may sometimes be written as single bonds alternating with double bonds, but the view of all ring bonds as being equivalently about 1.5 bonds in strength, is much closer to truth.

In the case of heterocyclic aromatics and substituted benzenes, the electronegativity differences between different parts of the ring may dominate the chemical behaviour of aromatic ring bonds, which otherwise are equivalent.

General properties of ionic compounds :

a)physical state -

At room temp ionic compounds exist either in solid state or in solution phase but not in gaseous state.

b) Isomorphism

Simple ionic compounds do not show isomerism but isomorphism is their imp characteristic. Isomorphism means similar crystal structure

c) Electrical conductivity

Ionic solids are almost non-conductors.

However they conduct a very little amount of current due to crystal defects.

All Ionic solids are good conductor in molten state as well as in their aqueos solution because their ions are free to move.


Covalent bond is formed by sharing of same no of electrons b/w two atoms to complete their octate.

Atoms taking part in covalent bond formation may share one, two or three electron pairs and thus forms single, double or triple bonds resp. Covalent bonding is a common type of bonding, in which the electronegativity difference between the bonded atoms is small or nonexistent. Bonds within most organic compounds are described as covalent. See sigma bonds and pi bonds for LCAO-description of such bonding.

A polar covalent bond is a covalent bond with a significant ionic character. This means that the electrons are closer to one of the atoms than the other, creating an imbalance of charge. They occur as a bond between two atoms with moderately different electronegativities, and give rise to dipole-dipole interactions.

A coordinate covalent bond is one where both bonding electrons are from one of the atoms involved in the bond. These bonds give rise to Lewis acids and bases. The electrons are shared roughly equally between the atoms in contrast to ionic bonding. Such bonding occurs in molecules such as the ammonium ion (NH4+) and are shown by an arrow pointing to the Lewis acid.

Molecules which are formed primarily from non-polar covalent bonds are often immiscible in water or other polar solvents, but much more soluble in non-polar solvents such as hexane.

Single covalent bonding

Most common type of chemical bonding is single covalent bonding, where one pair of valence electrons is shared by the two atoms. Valence electrons are those that are in the outer orbit or shell of an atom.

Hydrogen molecule

A good example of single covalent bonding is the Hydrogen molecule (H2).

Each Hydrogen atom shares the other's valance electron

Since the atoms are sharing the other's electron, both appear to have the first orbit or shell filled with the maximum of two electrons.

Water molecule

Another example of single covalent bonding is the water (H2O) molecule. Adding hydrogen gas molecules (H2) to oxygen gas molecules (O2) can result in an explosion if lit by a flame or spark. The end product is the very stable water molecule (H2O). The chemical equation is:

2H2 + O2 → 2H2O

The resulting molecule has single covalent bonding:

Oxygen has single covalent bonding with each of the two Hydrogen atoms

You can see that with the sharing of electrons, each Hydrogen atom has two valence electrons, thus filling their outer orbits. Likewise, Oxygen now has 8 outer orbit +electrons. This makes for a good chemical bond and a stable molecule. Usually, electrolysis is required to separate the hydrogen and oxygen from water.

Multiple covalent bonding

Some molecules are held together with double, triple and even quadruple covalent bonding. That means that two, three or four pairs of electrons are shared. A vast majority of multiple covalent boding is double covalent bonding.

Double covalent bonding

Since its outer orbit is missing 2 electrons, you never see an Oxygen atom by itself, because and it can readily combine with another Oxygen atom to form a more stable Oxygen molecule (O2).

The Oxygen molecule is held together by a double covalent bonding.

Oxygen molecule employs double covalent bonding

The problem with this "solar system" diagram of the atoms is that sharing two paris of electrons just doesn't look right. We know it happens, but the illustration seems confusing. It gets even more confusing in molecules having triple or quadruple covalent bonding.

Electron dot notation

Thus a method to simplify a diagram of the molecule was devised. It only shows the valence electrons as dots. It is called the electron dot notation. (It is also often called the Lewis dot notation, after the person who invented it.)

Oxygen molecule in dot notation

Water molecule in dot notation

There are a few other ways of diagramming molecules to better illustrate the covalent and even ionic bonding.

There are also a few other types of chemical bonding that are not common enough to go into in this lesson.

Valence bond theory

In the year 1927, valence bond theory was formulated which argued essentially that a chemical bond forms when two valence electrons, in their respective atomic orbitals, work or function to hold two nuclei together, by virtue of system energy lowering effects. In 1931, building on this theory, chemist Linus Pauling published what some consider one of the most important papers in the history of chemistry: “On the Nature of the Chemical Bond”. In this paper, building on the works of Lewis, and the valence bond theory (VB) of Heitler and London, and his own earlier work, he presented six rules for the shared electron bond, the first three of which were already generally known:

  1. The electron-pair bond forms through the interaction of an unpaired electron on each of two atoms.
  2. The spins of the electrons have to be opposed.
  3. Once paired, the two electrons cannot take part in additional bonds.

His last three rules were new:

  1. The electron-exchange terms for the bond involves only one wave function from each atom.
  2. The available electrons in the lowest energy level form the strongest bonds.
  3. Of two orbitals in an atom, the one that can overlap the most with an orbital from another atom will form the strongest bond, and this bond will tend to lie in the direction of the concentrated orbital.

Building on this article, Pauling's 1939 textbook: On the Nature of the Chemical Bond would become what some have called the “bible” of modern chemistry. This book helped experimental chemists to understand the impact of quantum theory on chemistry. However, the later edition in 1959 failed to address adequately the problems that appeared to be better understood by molecular orbital theory. The impact of valence theory declined during the 1960s and 1970s as molecular orbital theory grew in popularity and was implemented in many large computer programs. Since the 1980s, the more difficult problems of implementing valence bond theory into computer programs have been largely solved and valence bond theory has seen a resurgence.

Comparison of valence bond and molecular orbital theory

In some respects valence bond theory is superior to molecular orbital theory. When applied to the simplest two-electron molecule, H2, valence bond theory, even at the simplest Heitler-London approach, gives a much closer approximation to the bond energy, and it provides a much more accurate representation of the behavior of the electrons as chemical bonds are formed and broken. In contrast simple molecular orbital theory predicts that the hydrogen molecule dissociates into a linear superposition of hydrogen atoms and positive and negative hydrogen ions, a completely unphysical result. This explains in part why the curve of total energy against interatomic distance for the valence bond method lies above the curve for the molecular orbital method at all distances and most particularly so for large distances. This situation arises for all homonuclear diatomic molecules and is particularly a problem for F2, where the minimum energy of the curve with molecular orbital theory is still higher in energy than the energy of two F atoms.

The concepts of hybridization are so versatile, and the variability in bonding in most organic compounds is so modest, that valence bond theory remains an integral part of the vocabulary of organic chemistry. However, the work of Friedrich Hund, Robert Mulliken, and Gerhard Herzberg showed that molecular orbital theory provided a more appropriate description of the spectroscopic, ionization and magnetic properties of molecules. The deficiencies of valence bond theory became apparent when hypervalent molecules (e.g. PF5) were explained without the use of d orbitals that were crucial to the bonding hybridisation scheme proposed for such molecules by Pauling. Metal complexes and electron deficient compounds (e.g. diborane) also appeared to be well described by molecular orbital theory, although valence bond descriptions have been made.

In the 1930s the two methods strongly competed until it was realised that they are both approximations to a better theory. If we take the simple valence bond structure and mix in all possible covalent and ionic structures arising from a particular set of atomic orbitals, we reach what is called the full configuration interaction wave function. If we take the simple molecular orbital description of the ground state and combine that function with the functions describing all possible excited states using unoccupied orbitals arising from the same set of atomic orbitals, we also reach the full configuration interaction wavefunction. It can be then seen that the simple molecular orbital approach gives too much weight to the ionic structures, while the simple valence bond approach gives too little. This can also be described as saying that the molecular orbital approach is too

, while the valence bond approach is too localised.

The two approaches are now regarded as complementary, each providing its own insights into the problem of chemical bonding. Modern calculations in quantum chemistry usually start from (but ultimately go far beyond) a molecular orbital rather than a valence bond approach, not because of any intrinsic superiority in the former but rather because the MO approach is more readily adapted to numerical computations. However better valence bond programs are now available.

Metallic bond

In a metallic bond, bonding electrons are delocalized over a lattice of atoms. By contrast, in ionic compounds, the locations of the binding electrons and their charges are static. Because of delocalization or the free moving of electrons, it leads to the metallic properties such as conductivity, ductility and hardness.

Intermolecular bonding

There are four basic types of bonds that can be formed between two or more (otherwise non-associated) molecules, ions or atoms. Intermolecular forces cause molecules to be attracted or repulsed by each other. Often, these define some of the physical characteristics (such as the melting point) of a substance.

  • A large difference in electronegativity between two bonded atoms will cause dipole-dipole interactions. The bonding electrons will, on the whole, be closer to the more electronegative atom more frequently than the less electronegative one, giving rise to partial charges on each atomic center, and causing electrostatic forces between molecules.
  • A hydrogen bond is effectively a strong example of a permanent dipole. The large difference in electronegativities between hydrogen and any of fluorine, nitrogen and oxygen, coupled with their lone pairs of electrons cause strong electrostatic forces between molecules. Hydrogen bonds are responsible for the high boiling points of water and ammonia with respect to their heavier analogues.
  • The London dispersion force arises due to instantaneous dipoles in neighbouring atoms. As the negative charge of the electron is not uniform around the whole atom, there is always a charge imbalance. This small charge will induce a corresponding dipole in a nearby molecule; causing an attraction between the two. The electron then moves to another part of the electron cloud and the attraction is broken.
  • A cation-pi interaction occurs between the negative charges of pi bonds above and below an aromatic ring and a cation.

Electrons in chemical bonds

In the (unrealistic) limit of "pure" ionic bonding, electrons are perfectly localized on one of the two atoms in the bond. Such bonds can be understood by classical physics. The forces between the atoms are characterized by isotropic continuum electrostatic potentials. Their magnitude is in simple proportion to the charge difference.

Covalent bonds are better understood by valence bond theory or molecular orbital theory. The properties of the atoms involved can be understood using concepts such as oxidation number. The electron density within a bond is not assigned to individual atoms, but is instead delocalized between atoms. In valence bond theory, the two electrons on the two atoms are coupled together with the bond strength depending on the overlap between them. In molecular orbital theory, the linear combination of atomic orbitals (LCAO) helps describe the delocalized molecular orbital structures and energies based on the atomic orbitals of the atoms they came from. Unlike pure ionic bonds, covalent bonds may have directed anisotropic properties. These may have their own names, such as Sigma and Pi bond.

In the general case, atoms form bonds that are intermediates between ionic and covalent, depending on the relative electronegativity of the atoms involved. This type of bond is sometimes called polar covalent.