The atoms and molecules in a solid state are more closely packed as compared to gaseous and liquid states and are held together by strong mutual forces of attraction. These interatomic forces are electrostatic in nature and depend upon the electronic structure of matter.
When atoms come closer and finally unite to form molecules their electrons rearrange themselves in such a way as to achieve a stable configuration. This arrangement of electrons gives rise to different types ul bonds due to which atoms are held together.
Bonds can be broadly classified as:
1. Primary bonds
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2. Secondary bonds
1. Primary Bonds (or Chemical Bonds):
These are the strongest bonds between atoms which can be further divided as follows:
(i) Ionic (or electrostatic) bond.
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(ii) Covalent (or atomic or homopolar) bond.
(iii) Metallic bond.
2. Secondary (or Molecular) Bonds:
Attraction forces (also called ven der Waals forces) exist between atoms or molecules. These bonds are weaker than primary bonds.
1. Primary Bonds:
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i. Ionic Bond:
Ionic bonds are mainly formed in inorganic compound like NaCl and KOH etc., and never in pure elements. An ionic bond is really the attractive force between a positive ion and a negative ion when they are brought into close proximity. These ions, of course, are formed when the atoms involved loss or gain electrons in order to stabilise their outer shell electron configuration. Elements are classified as either electropositive or electronegative, depending upon when they tend to lose or gain electrons in order to achieve this stable outer shell electron configuration.
Let us consider the combination which lakes place between the Sodium (metal) and Chlorine (non-metal) to form sodium chloride. The sodium has a single electron in its outer shell and this transfer to join the seven electrons in the outer shell of chlorine atom.
This type of atomic interaction, involving the outright transfer of one electron form one atom to another, leads to formation of ions which are held together by electrostatic attraction. Because of the electrostatic nature of the binding force, the bond between atoms is said to be ionic or electrovalent.
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Characteristics of Ionic Compounds:
(i) Ionic bonds are unidirectional.
(ii) They are generally crystalline in nature.
(iii) They are rigid.
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(iv) They have high melting and also high boiling point because strong electrostatic forces bind them.
(v) They are generally non-conductors of electricity but their metals and solutions conduct electricity.
(vi) They are usually highly soluble in water and other polar solvents but insoluble in organic solvents.
ii. Covalent Bond:
The covalent bond is formed by sharing of electrons between atoms rather than by transfer of electrons. Only a few solids are held together by covalent bonds. Covalent bonding alone is not sufficient to build three dimensional solids. The majority of solids incorporating covalent bonds are bound also either by ionic or van der Waals bonds.
An excellent example of covalent bonding is found in the Chlorine molecule. Here the outer shell of each atom possesses seven electrons. Each chlorine atom would like to gain an electron and thus form a stable octet. This can be done by sharing of two electrons between pairs of chlorine atoms thereby producing stable diatomic molecules.
Characteristics of Covalent Compound:
(i) Covalent compounds are directional in nature.
(ii) They can be solids, liquids and gasses.
(iii) They are soft and have low melting and boiling points.
(iv) They are generally insulators electrically.
(v) They are soluble in non-polar solvents such as benzene and carbon tetrachloride and insoluble in water.
iii. Metallic Bond:
A metallic bond results from the sharing of variable number of electrons by a variable number of atoms. The atoms in metal and alloys are held together by such bonds. This type of bond is characteristic of the elements having small number of valence electrons, which are loosely held, so that they can be easily released to the common pool.
The bonding takes place when each of the atoms of the metal contributes its valence electrons to the formation of an electron cloud that pervades the solid metal. The valence electrons are not bonded directly to an individual atom but they move freely in the sphere of influence of other atoms and are bound to different atoms at different times and that too for a short time.
Fig. 42 shows a schematic picture of the metal ions (positive) and that of electron cloud (negative). The cohesion of a metallic crystal is due to the attraction of the positive nuclei and the valence electrons passing between them. A metallic bond thus conceived can exist only between a large aggregate of metallic atoms and must therefore, be non-directional. The high electrical conductivity of metals is given by the free electrons moving freely in an electric field.
Characteristics of metallic compound:
(i) Metallic compounds are crystalline in nature.
(ii) They are very good conductors of electricity.
(iii) They have usually moderate to high, melting temperature (since the metallic bonds are not very strong).
(iv) They are opaque to light (since free electrons in a metal absorb light energy).
(v) They have high thermal conductivity
(vi) They have high reflectivity and lustre.
Comparison of Primary Bonds:
1. Bond Energy:
Ionic Bond:
Higher than metallic bond.
Covalent Bond:
Higher than metallic bond.
Metallic Bond:
Generally lower than other primary bonds.
2. Nature of Bond:
Ionic Bond:
Non-Directional.
Covalent Bond:
Directional.
Metallic Bond:
Non-Directional.
3. Bond Formation:
Ionic Bond:
This type of bond is most easily formed when one of atoms has small number of valance electrons such as alkali metals and alkali earths.
Covalent Bond:
This type of bond is formed when atomic orbitals of two atoms overlap. An electron in each atom can exchange with an electron in its partner atom. Elements that form molecules with this type of bonding include those with four or more valence electrons, but hydrogen is an exception.
Metallic Bond:
A metallic bond results from the sharing of variable number of electrons by a variable number of atoms. This type of bond is characteristic of elements having small number of valence electrons which are loosely held, so that they can be easy released to the common pool.
4. Bonding Force:
Ionic Bond:
This bond exists due to electrostatic force.
Covalent Bond:
Electrostatic force of attraction between two atoms.
Metallic Bond:
Electrostatic force of attraction between the free electron cloud of valence electrons and positive ions of the same or different metallic elements.
5. Conductivity:
Ionic Bond:
Ionic compounds have low conductivity.
Covalent Bond:
Low conductivity.
Metallic Bond:
Most of the metallic compounds have high thermal and electrical conductivity.
6. Melting and Boiling Point:
Ionic Bond:
Higher melting and boiling points in general.
Covalent Bond:
They generally possess lower melting and boiling points but exceptions are there with very high melting and boiling points.
Metallic Bond:
Melting and boiling points generally high.
7. Mechanical Properties:
Ionic Bond:
High hardness. These crystals tend to cleave along certain planes of atoms rather than to deform in a ductile fashion when subjected to stresses.
Covalent Bond:
Solids formed are generally hard.
Metallic Bond:
Good strength and ductility in general.
8. Solubility:
Ionic Bond:
Soluble in water liquid, ammonia, etc.
Covalent Bond:
Not soluble in water, however, easily soluble in organic solvents like benzene.
Metallic Bond:
Not soluble in water and organic solvents.
2. Secondary or Molecular Bonds:
Molecular bonds are formed in case of those elements or compounds whose electron configuration is such that little transfer takes place between atoms. These bonds are formed as a result of weak van der Waals forces of attraction which exist between various atoms.
These forces are due to the electrostatic attraction between the nucleus of one atom and the electrons of the other. This is largely but not completely neutralized by the electrostatic repulsion of the nucleus of one atom by the nucleus of the other. The resultant weak attraction between the two atoms is called van der Walls force.
Three types of intermolecular bonds known are:
i. Dispersion bond
ii. Dipole bond
iii. Hydrogen bond.
i. Dispersion Bond:
These bonds are made possible largely because the electrons of adjacent atoms in a molecule tend to repel each other. As the electrons rotate around their nuclei, they end to keep in phase for a hydrogen molecule. The result is that the molecule has a small fluctuating net charge on each end and acts as an oscillating dipole.
The hydrogen molecule is instantaneously charged negatively on the right end and positively on the left. This fluctuating charge on one molecule tends to interact with the fluctuating charge on a neighbouring molecule, resulting in a net attraction. The strength of the bond depends on the ease with which one atom can influence the other.
Molecules of inert gases which consists of single atoms, are held together by dispersion forces when the gases are solidified. In many organic solids the most important bonding forces between molecules are of this type.
ii. Dipole Bond:
I. These bonds are caused by permanent electric charges on some molecules.
Example:
In the covalent bonded hydrogen chloride (HCl) molecule the net effect of the electron-sharing process is to give the chlorine atom a slightly negative charge while the hydrogen atom has a corresponding positive charge. The charges are actually very small being 0.816 x 10-10 esu (electrostatic units) with 4.8 x 10-10 esu for a single electron.
The spacing of the atoms is 1.28 x 10-8 cm, resulting in net dipole moment of the 1.04 x 10-18 cm x esu. Adjacent HCl molecules therefore, attract each other by means of the electrostatic attraction between their oppsitively charge ends. The attraction is small compared with that between ions because the charge on an ion is at least equal to that of one electron, 4.8 x 10-10 esu.
Dipole bonds are much weaker than bonds, but at the same time they are considerably stronger than disperson bonds.
Some other materials subject to dipolar bonding, and their dipole moments in cm x esu are –
(i) Sulphur dioxide (SO2)… 1.60 x 10-18
(ii) Hydrogen bromide (HBr)…0.78 x 10-18
(iii) Hydrogen cyanide (HCN)…2.93 x 10-18
iii. Hydrogen Bond:
The hydrogen bond might be considered as a special type of dipole bond, but one that is considerably stronger. It occurs between molecules in which one end is a hydrogen atom. The one electron belonging to the hydrogen atom is fairly loosely held, and if the adjacent atom in the molecule is strongly electro-negative, it may keep all the electrons around itself, leaving the hydrogen atom in effect a positive ion. A strong permanent dipole is created that can bond to other similar dipoles with a force near that involved in the ionic bond.
Example:
A good example of hydrogen bonding is water or ice. In water the hydrogen and oxygen atoms are held by covalent bonds in a configuration as shown in Fig. 44. It is seen that this special form of dipole bond is connected as the dipole moment of the H-0 bond rather than that of the molecule as a whole.
Characteristics of Molecular Bonds:
(i) Molecular bonds can be both crystalline and non-crystalline.
(ii) They have low melting point (because of weak molecular bond).
(iii) They are usually transparent to light.
(iv) They are good insulators (since they do not have valence electrons).
Mixed Bonding:
It is possible to have mixed bonding. In fact bonding between atoms in many materials cannot be classified as one of the four ideal types, i.e. – ionic, covalent, metallic and molecular (van der Waals) bonds but rather as a mixture of those types.
The apices represent van der Waal’s ionic, metallic and covalent bonding.
The various types of mixed bonding are represented on the edges.
The metals gradually change from pure metallic bonding as in sodium to the less perfect metals such as tellurium and arsenic, finally reaching the pure covalent bonding of carbon in the form of diamond.
Starting at diamond, we find graphite, benzene rings and high polymers. Ultimately, we reach the rare gases and pure van der Waals bonding.
In general, there are no mixtures between metallic and van der Waals bonding.