Bonding Assignment Help


In a bond ,two atoms are bound by either donating-accepting or sharing of electrons.In a molecule like NaCl, the electronegativity difference between Na and Cl is quiet substantial,which allows the chlorineto 'snatch' the electron from sodium.Thus ionic bond is formed.

In a molecule having bond like C=O,the electronegativity of oxygen is greater than carbon,however not to that extent as would enable the to snatch the electron from .Thus the density of electrons in the bond is 'pulled ' towards oxygen giving rise to polar covalent bond.

In case of C-C bond,since there is no difference between the electronegativities of two atms,it is non-polar covalent bond.

The molecule formed by more than two atoms coming together to form new compound.  The atoms come together and are connected or held together by bonds.

Introduction to molecular bonding

The bonds are composed of electrons from the constituent atoms. The electrons are contributed by the constituent atoms, one each. There are two major types of bondings. They are ionic bonding and covalent bonding

Molecular Bonding : Ionic Bonding

Let us consider an example of sodium chloride molecule.  Sodium has an excess electron than required for the stable configuration.  Similarly the chlorine atom has one less electron in its configuration than required for a stable configuration.
In addition to this, the electronegativity of chlorine is very high.the electronegativity of sodium is very low .  This difference in the in electronegativities of the atoms makes one atom as donor and another as an acceptor. This difference in the electronegativities allows chlorine to snatch an electron from the sodium-atom.  In other words, it is that sodium donates and chlorine accepts the electron.  As a result, a bond is formed between chlorine and sodium atom.  This bond is called as ionic bond.  Another example of and ionic bond is KCl.

Molecular Bonding : Covalent Bonding

Unlike As mentioned above, in a molecule, two atoms may have no difference or very low difference in the electronegativities.  As such, no atom can snatch an electron from the onother atom.  However, they may still share the electrons.  The bonds so formed are called  Covalent bonds.  Covalent bonding can be two types.
Consider the example of C=O bond.  The electronegativity of oxygen is more than that of carbon.  As such the electron density is pulled more towards oxygen and a polarity is built in the molecule.  Such a bond is known as polar covalent bond.
However, when there is no such difference in the electronegativities, the electrons are shared equally and this one is known as non-polar bond.

Types of Chemical Bonding

Let us see about types of chemical bonding. The presence of 8 electrons in the outermost shell of their atoms is responsible for stability of the noble gases, therefore atoms of all elements tend to obtain 8 electrons in their outermost shell by losing/gaining electrons with other atoms and this is the reason of their chemical combinations.

Types of Bonding

There are 6 types of chemical bonding. The types are,

  • Sigma Bonding
  • Pi Bonding
  • Covalent compounds Bonding
  • Co-ordinate compound Bonding
  • Metallic Bonding
  • Hydrogen Bonding



The types of chemical bonding explanation is as follows.

Sigma Bonding

     In the formation of ethane molecule, one of the hybridized bonds of each carbon atom chemical is used up in giving C-c bond and the remaining three bonds of each carbon atom are linked with hydrogen atoms. That means, each bond inethane is formed by the overlapping of orbital along their axes. Such bonds are known as sigma bonds.

Pi Bonding

     In compounds of beryllium, boron and carbon, one pair of electrons is mutual between two atoms in the structure of a bond. In some compounds, particularly of carbon, two or three pairs of electrons are concerned in a bond which is the n known a double or a triple bond.

Covalent Compounds Bonding

     In large number of compounds, atoms also mingle by sharing of the electronics in their outermost shells to inclusive their respective octets. This type of linkage is known called covalent bond. The compounds formed by this method known as covalent compounds bond.

Co-ordinate Compound Bonding

     A different types of covalent linkage can be formed when both electrons for sharing between the two atoms are donated by one atom. This kind of linkage is called as coordinate bond.

Metallic Bonding

     Metals normally have low ionization energies since the electrons in the outer shell can be taken out moderately easily. The force that binds a metal atom to a number of electrons within its sphere of power is called as a metallic bond.

Hydrogen Bonding

     In composites of hydrogen with powerfully electronegative elements a such as fluorine, oxygen and nitrogen the electron pair shared between the two atoms fabrication so far away from the hydrogen nucleus that the latter becomes highly positive.

Two types of covalent bonds are formed depending on the way the two atomic orbitals overlap with each other.

Sigma Bond (s)

When the overlap of orbitals of two atoms takes place along the line joining the two nuclei (orbital axis) then the covalent bond formed is called sigma (s) bond. These bonds can be formed due to 's-s', 's-p' or 'p-p' overlap along the orbital axis. Free rotation around a sigma bond is always possible.


Fig: 6.4 - Formation of sigma bond due to various overlapping

Pi bond (p)

When the two atoms overlap due to the sideways overlap of their 'p' orbitals, the covalent bond is called as pi(p) bond. In a pi bond the electron density is concentrated in the region perpendicular to the bond axis.

Characteristics of the pi (p) bonds

  • The pi bonds are weak because the orbtial overlap is partial.
  • For a complete side ways overlap the 'p' orbitals should be parallel to each other. This is possible when all the atoms of the molecule are in the same plane i.e., there is no rotation of one part of the molecule relative to the other about the pi (p) bond.


Fig: 6.5 - Rotation about a double bond

  • In molecules containing double bond, which has a pi bond, there is no free rotation and such molecules exist in isomeric forms of 'cis' and 'trans'. In the 'cis' form the similar atoms lie on the same side of a plane placed along the internuclear axis while in the 'trans' form the similar atoms lie on the opposite sides.


general representation of cis- trans isomers


  • The electrons in the pi (p) bond are placed above and below the plane of the bonding atoms and so they are more exposed. They are more susceptible to attack by electron seeking or oxidizing agents. Hence, they are the most reactive centers in unsaturated (multiple bonded) compounds.

Comparative properties of sigma and pi bonds

Sigma (σ) bond Pi (π) bond
Formed due to the axial overlap of two orbitals (‘s-s’, ‘s-p’or’p-p’).Formed by the lateral (sideways) overlap of two ‘p’ orbitals.
Only one sigma bond exists between two atoms. There can be more than one pi bonds between the two atoms.
The electron density is maximum and cylindrically symmetrical about the bond axis. The electron density is high along the direction at right angles to the bond axis.
Free rotation about the sigma bond is possible. Free rotation about the pi bond is not possible.
This bond can be independently formed, i.e., without the formation of a pi bond. The pi bond is formed after the sigma bond has been formed,
Sigma bond is relatively strong. Pi bond is a weak bond.




11. Explain how the valence bond theory accounts for the existence of cis- trans isomers.


A double bond between carbon-carbon atoms (C=C) consists of a

sp2-sp2 s bond and a p bond. These bonds are formed between unhybridized '2p' orbitals of each C atom at right angle to the plane of sigma bond. Since the p bond prevents the free rotation about the carbon-carbon axis, a di-substituted alkene can exist in two forms: cis-isomer and trans-isomer.

In cis-isomer, both the like groups are on the same side of the plane whereas in trans-isomer these groups are on the opposite side of the plane.

For example 1, 2-dichloroethene.


In order to transform one isomer into the other, one end of the molecule must be rotated as the other remains fixed. The bond must be broken in order to do this. Breaking the p bond requires considerable energy, so the cis and trans compounds are not easily inter converted.

12. How many and bonds are present in naphthalene?


Naphthalene is


pbonds = 19, p bonds = 5

List of Covalent Bonds:   

          A covalent bond is a compound bond produced as two atoms common distribute a brace of electron.  The atoms achieve constant electronic arrangement. Covalent bonding, go beyond of atomic orbital comprise an electron as of every of two atoms of connection obtain position ensuing in equivalent distribution of brace of electrons.

List of Covalent Bonds:

         The inter-automatic connection hence produced suitable to the overlap of atomic orbital’s of electrons in the valence case of the particle are utilize for electron distribution. The common pair of electrons lies in the middle of the covalent bond.

         Thus in hydrogen molecule a covalent bond results by the overlap of the two s orbital’s every have an electron as of each of the two H atoms of molecule. Every H atom achieves filled K shell. A covalent bond knows how to be produced by distribution of s, p, d, f electrons besides.

Characteristic of covalent bond:

         Covalent composite are produced by the common distribution of electrons. Here is no transmit of electrons as of individual atom to a different with therefore no accuse are fashioned on the atom.

         They group squat melting with boiling position. This is as of the feeble put in the ground molecular services obtainable among covalent molecules. As no brawny columbic services are several of covalent bonds are explosive in environment. Typically covalent composite acquire stumpy melting and boiling position.

         The majority of the covalent bond are non polar with in non-polar solvents like benzene and inexplicable in polar solvents similar to water. Carbon tetrachloride is a covalent non-polar molecule with is soluble in benzene.

List of covalent bonds:

  • Molecule of Hydrogen (H = 1)
  • Molecule of Oxygen (O = 8)
  • Molecule of Nitrogen (N = 7)
  • Molecule of Chlorine (Cl = 17)
  • Molecule of Water (H = 1, O = 8)
  • Molecule of Hydrogen Chloride (H = 1, Cl = 17)
  • Molecule of Ammonia (H = 1, N = 7)
  • Molecule of Methane (H = 1, C = 6)
  • Molecule of Carbon Tetrachloride (C = 6, Cl = 17)
  • Molecule of Carbon Dioxide (C = 6, O = 8)

Types of Co-ordination Compounds

 Name  Formula  Charge  Name of Ligand
 Ammonia  NH3  Zero  ammine
 Water  H2O  Zero  aquo or aqua
 Halide ion  X-(X=Cl,Br,I)  -1  halo
 Hydroxide ion  :OH -  -1  hydroxo
 Cyanide ion  :CN -  -1  Cyano
 Oxide ion  O2 -  - 2  Oxo
 Peroxide ion  O22 - - 2 Peroxo
 Carbonate ion  CO32 - - 2  Carbonato
 Phosphine  PH3  Zero  Phosphine
 Nitrogen oxide  NO Zero Nitrosyl
 Carbon monoxide  CO Zero  Carbonyl
 Sulphate ion  SO42 - -2 Sulphato
 Nitrite ion  NO2-  -1  Nitro or nitrite
 Thiocyanate ion  SCN-  - 1  Thiocyanato or isothiocyanato ion
 Acetate ion  CH3COO-  - 1  Acetate
 Pyridine  C5H5N  Zero  Pyridine(PY)
 Sulphide ion  S2 -  - 2  Sulphido
 Thiosulphate  S2O32 -  - 2  Thiosulphato
 Nitrate ion  NO3-  -1  Nitrato
 Sulphite ion  SO32-  -2  Sulphito
 Triphenyl phosphine  (C6H5)3P  Zero  Triphenyl phosphine
 Thiocarbonyl  CS  Zero  Thiocarbonyl
 Nitrosonium  NO+  +1  Nitrosonium
 Nitronium  NO2+  +1  Nitronium
 Imide ion  NH2 -  - 2  Imido
 Nitrate ion  NO3-  -1  Nitrato
 Amide ion  NH2-  -1  Amido
 Thio urea H2NCSNH2  Zero  Thiourea(tu)


Some multidentate ligands






Examples of ligands



Metallic Bonding


In simple terms metallic bonding is referred to as bonding in metal atoms. It is also defined as interaction between metal nuclei and the delocalized electrons. Delocalized electrons are also called as conduction electrons. Metals nuclei are the positive ions and so Metallic bonding can be imagined as Sea of electrons in which positive metal ions are embedded. Positive metal ions are called as Kernels. Thus metallic bonding can be summarized as: - The force of attraction which binds together the positive metal ions or Kernels with the electrons within its sphere of influence.

Diagrammatic Representation of Metallic Bonding: -

metal ions of Zn

Example: - Zn -------->  Zn2+  +  2e-

These electrons are the delocalized or conduction electrons and Zn2+ are the metal ions.

The geometrical arrangements of metallic bonding can be as follows: - Face centered cubic (fcc), Hexagonal close packed (hcp), or Body centered cubic (bcc). These bonds are weaker than covalent bonds because the electrons are delocalized and mobile, therefore do not experience very strong force of attraction towards the nuclei.

Properties of Metals due to Metallic Bonding: -

1) Metals conduct heat and electricity due to mobile electrons. These act as charge and energy carriers.

2) Metals have high melting and boiling points and the values depend on the strength of the metallic bond which in turn depends on: - packing of metals ions and electrons, and also on the number of delocalized electrons.

3) Metals become malleable and ductile because the delocalized electrons can spread over thus metals can be beaten into sheets and stretched like wires.

4) Metals are lustrous because photons of light are not able to penetrate more and are reflected back quickly.

5) Metallic bonding makes the metal flexible.

Thus metallic bonding effects most of the physical properties of metals and chemically metals becomes electron donors so are highly electropositive forming cations and giving electrons for the reaction to occur. Thus they act as Reducing agents and Lewis bases.

Hydrogen Bonding


The attractive force which binds hydrogen atom of one molecule in the midst of electronegative atom of another molecule of the same substance is known as hydrogen bond. The internal hydrogen bonding also called as internal Hydrogen Bridge, for the reason that hydrogen atoms act a deferment bridge stuck between two electronegative atoms. The internal hydrogen bonding is generally, denoted by a dotted line.

Internal Hydrogen Bonding in Water:

  • Some molecular solid crystals are held together by hydrogen bonding between atoms of different electronegativities.
  • The bond length between internal hydrogen bonding H-----O is more than twice that of O--H bonds in water molecule.
  • Each water molecule is thus tetrahedral surrounded by four other water molecules, thereby giving a strong structure with large number of hydrogen bonding.
  • Strong orientation of hydrogen atoms toward oxygen atoms gives rise to a comparatively less-efficiently packed crystal structure
  • Evidently, ice has an open-cage type structure.
  • When ice melts, the open structure breaks down and molecules pack more closely together, so that water has a higher density.
  • This breaking down proceed is not complete, awaiting a temperature of 4°C is reached.
  • Above 4°C, there is normal effect of expansion and there is increase in volume and consequently, density of water is lowered.

Consequences of Internal Hydrogen Bonding:

  • Association: Two or more molecules linked by internal hydrogen bonding associated with one another to form associated units.
  • Higher melting and boiling point, due to higher effective molecular mass of the associated molecules.
  • The presence of hydrogen bonding in compounds is indicated by higher melting-points, boiling points, heat of vaporisation and viscosities than shown by corresponding normal liquids.
  • The abnormal melting and boiling points of NH3,H2O and HF are due to association caused by the formation of  internal hydrogen bonds.


  • Generally, compounds which can form hydrogen bonds with the solvent molecules are soluble in such solvents.
  • For example, lower alcohols are soluble in water, due to hydrogen bonding, which can exist between alcohol and water molecules.

Hydrogen Bonding

Hydrogen Bonding is a type of interaction or an electrostatic force of attraction between a hydrogen atom and a highly electronegative atom such as fluorine, oxygen and nitrogen.

When a hydrogen atom is attached to a very electronegative atom (such as F, O, or N) the bond is highly polar and the attraction between the H-X (X= F, O, or N) dipole and other polar molecules is greater than would be expected for a typical dipole-dipole interaction.

The hydrogen bond strength generally covers the range of 5 – 50 kJ/mol.  Also, very few hydrogen bonds are observed in the vapor phase.

X-ray diffraction and Spectroscopy are simple methods of comparing hydrogen-bonded systems in the solution or liquid phase.

Introduction to Examples of Hydrogen Bonding

Conditions for hydrogen bond formation, the most important type of molecular forces of attraction are:

i) High electronegativity of the atom bonded to hydrogen atom so that bond is ‘Polar’ in nature.  

ii) The small size of atom bonded to hydrogen so that it is able to attract the bonding electron pair effectively.

Hydrogen bonding can arise between molecules as characterised in intermolecular hydrogen bonding and can arise between side chains within a molecule as in intramolecular hydrogen bonding.

The Intermolecular Hydrogen Bonding:

Such a hydrogen bond exists between the hydrogen atom of one molecule and an electronegative atom of another molecule as in water, ammonia, hydrogen flouride, etc...

The Intramolecular Hydrogen Bonding:

Such a hydrogen bond exists within the same molecule as in salicylic acid, orthonitrophenol, etc...

Examples of Hydrogen Bonding

Case 1:

In the case of salicylaldehyde, the cis isomer of this 2-hydroxybenzaldehyde melts at 1 °C, while the trans isomer has melting and boiling points of 112 °C.  Both compounds exhibit strong hydrogen bonding in the solid state. However, cis-hydroxybenzaldehyde (salicylaldehyde) has a configuration that allows strong intramolecular hydrogen bonding, which precludes any intermolecular hydrogen bonding. The melting point of cis-hydroxybenzaldehyde is going to be controlled by the van der Waal forces between adjacent molecules.

In contrast, since intramolecular hydrogen bonding is precluded in the trans-isomer it can form strong intermolecular hydrogen bonds in the solid state, and thus, it is these that define the melting point. The boiling points are controlled in a similar manner in the case of 2- and 4- hydroxybenzoic acid.  The 2-Hydroxybenzoic acid exhibits strong intramolecular hydrogen bonding while the 4-hydroxybenzoic acid has strong intermolecular hydrogen bonding.

Case 2:

Let us next take the example of o-nitrophenol and p-nitrophenol in water.
o-nitrophenol is capable of forming intramolecular hydrogen bonds, implying that hydrogen bonding is between the oxygen of the phenolic group and the nitrogen of the nitro group of the same molecule.Thus the possibility of hydrogen bonds between the solvent (water) and the solute is decreased whereby the solubility of o- nitrophenol is low.
p-nitrophenol is capable of forming intermolecular hydrogen bonding between water and p-nitrophenol. So p-nitrophenol is highly soluble in water.

Case 3:

Again, let us consider the volatility of these two compounds. In the case of o-nitrophenol, as it forms intramolecular hydrogen bonding, there is no significant change in its molecular weight. There is thus, no change in its boiling point as well. Lower the boiling point, more volatile the substance will be. So, o-nitrophenol is MORE VOLATILE.

In p-nitrophenol, as there are intermolecular hydrogen bonding, it increases the molecular weight, whereby increasing the boiling point. So p-nitrophenol is less volatile.

Case 4:

The acidity of a chemical species that donate H+ can be affected by the presence of hydrogen bonding. For example, consider the di-carboxylic acid derivatives of ethylene (fumaric acid). 

The acidity of the first and second carboxylic group for the trans isomer is similar, the difference being due to the increased charge on the molecule.

In contrast, the second proton in the cis isomer is much less acidic than the first proton because a consideration of the structure of the mono anion of the cis isomer shows a very strong intramolecular hydrogen bond is formed once the first proton is removed. This hydrogen bond makes the second acidic proton much more difficult to remove and thus lowers the acidity of the proton.

Case 5:

The cis isomer of butanedioic acid is called maleic acid and the trans is named fumaric acid. Though they are isomers, maleic acid and fumaric acid have totally different physical properties.

These properties of maleic acid can be explained on the basis of intramolecular hydrogen bonding.

a) Boiling point and melting point:
Polarity is a determining factor in  relative boiling point as it causes increased intermolecular forces, thereby raising the boiling point. In the same manner, symmetry is key in determining relative melting point as it allows for better packing in the solid state.

Thus, trans-alkenes which are less polar and more symmetrical have lower boiling points and higher melting points and cis-alkenes, which are generally more polar and less symmetrical have higher boiling points and lower melting points.

This could also be attributed due to the presence of double bonds among geometric isomers, especially when both substituents are the same. This is because:

Dipoles of substituents in a cis isomer will add up to give a net molecular dipole.

Dipoles of substituents in a trans isomer will cancel out due to their spatial arrangement.

b)    Lower densities:
It has also been observed that trans-isomers have lower densities than their trans structures. This can be explained on the basis of the symmetrical arrangement of the molecule (alkenes) in ‘trans’ structure.

c)  Solubility:
The above is the same reason for the higher solubility of trans alkenes in inert solvents, than their cis counterpart.

d)  Stability:
Trans isomers, usually, are more stable than cis isomers. This is due to its shape-    Straight shape leads to intermolecular hydrogen bonding that helps in increasing the stability of the molecule.

Case 6:

Hydrides of all group XVI elements are gaseous except for water (i.e., hydride of oxygen). This can be explained as follows. Hydrogen bonding is only possible between H and N, O and F. The group XVI elements are oxygen, sulfur, selenium, tellurium and polonium. Among these, hydrogen bonding is possible only between H and O. We know that in water there exists intermolecular hydrogen bonding. Presence of intermolecular hydrogen bonding leads to lower volatility.   Due to this reason, water is a liquid at room temperature.

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