1. Chemical bond : It is the attractive force that holds two or more atoms together in a molecule or ion. Why do atom Combine? 1. Net attractive force between atoms: 2. Lowering of energy of combining atoms: the process of chemical bonding between atoms decreases the energy of the combining atoms and gives rise to the formation of a system which has lower energy and hence has greater stability

2. 3. Octet rule or rule of eight: Kössel and Lewis in 1916 developed an important theory of chemical combination between atoms known as electronic theory of chemical bonding. ‘According to this, atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an octet in their valence shells. This is known as octet rule’.

3. Limitations of the Octet Rule:  The incomplete octet of the central atom: In some compounds, the number of electrons surrounding the central atom is less than eight. This is especially the case with elements having less than four valence electrons. Examples are LiCl, BeH2 and BCl3.  Li, Be and B have 1,2 and 3 valence electrons only. Some other such compounds are AlCl3 and BF3.  Odd-electron molecules: In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxide, NO 2, the octet rule is not satisfied for all the atoms.

4.  The expanded octet: In a number of compounds of elements in 3 rd period of periodic table, there are more than eight valence electrons around the central atom. This is termed as the expanded octet. Obviously the octet rule does not apply in such cases. Some of the examples of such compounds are: PF5, SF6, H2SO4 and a number of coordination compounds.

5.  Types of Chemical Bond Initially, chemical bonds can be divided into two types: strong bonds and weak bonds.  Strong Bonds: In these bonds, the bonding atoms rearrange (transfer or share) their outermost shell electrons to attain the electron configuration of noble gases. These include following four types of bonds:  Ionic bond  Covalent bond  Coordinate (covalent) bond &  Metallic bond.  Weak Bonds: In these bonds, the bonding atoms do not lose their identity. These include following two types of bonds:  Hydrogen bond &  Van der Waal’s forces.

6. CHEMICAL BOND (Continued) Ionic or Electrovalent Bond  Definition  The ionic or electrovalent bond is the chemical bond formed between two atoms by the transfer of one or more valence electrons from one atom to the other.  In other words, Ionic or electrovalent bond is the electrostatic attraction between the cation (+ve ion) and anion (  ve ion) produced by electron-transfer between two atoms.

7. CHEMICAL BOND (Continued)  Some Examples of Ionic Compounds: NaCl, MgO, CaF 3, MgCl 2, CaO 2, Al 2 O 3

8. 1)No. of valence electrons: A A + (1,2, or 3 valence electrons) IA, IIA and IIIA B B - (5, 6 or 7 valence electrons) VA, VIA & VIIA 2) The ionisation energy of the metal atom should be low: The minimum amount of energy required to remove the most loosely bound electron Ionisation energy A+ Energy required A + +e - ( ionisation energy ) Low I. energy easy formation of ionic bond Factors favoring the formation of ionic compounds

9. 3) Electronic affinity of the non-metal should be high: The amount of energy released when an electron is added to a neutral isolated gaseous atom electron affinity. B + e - B - + energy released (electron affinity) 4) The lattice energy of the ionic compound formed should be high: the energy released when one gram mole of a crystal is formed from its gaseous ions is called the lattice energy of the crystal. A + (1 mole) + B - (1 mole) A + B -

10. Higher the value of lattice energy of a crystal, the greater is the ease of its formation greater will be the stability of the ionic crystal 5) Electronegativity difference of A and B should be high: A difference of 2 or more is necessary for the formation of an ionic bond between atoms. Properties of ionic bond : 1)Crystalline state: ionic crystals are clusters of ions in the crystal lattice. Each Na + is surrounded by equally spaced six Cl - ions placed at the corners of a regular octahedron and similarly each Cl - is surrounded by equally spaced six Na + placed at the corners of an octahedron. 2)Electrical conductivity : Ionic compound in soild state do not conduct electricity Ionic compound in solution/ fused state conduct Electricity 

11. 3) Isomorphism: Ionic solids made up of ions with identical electronic configuration show an identity of crystalline form which is called isomorphism e. g. (a) sodium fluoride & magnesium oxide + Na + (2, 8) F - (2,8) Mg +2 (2,8) O 2- (2,8) (b) Calcium chloride & Potassium sulphide Cl - (2,8,8) Ca +2 (2,8,8) Cl - (2,8,8) K + (2,8,8) S 2- (2,8,8) K + (2,8,8) 4) Dielectric constant: The substances containing ionic bonds have high dielectric constant. 5) High melting and boiling points: very high amount of energy (in the form of heat) is required. e.g. NaCl boils at 1470 0 C CCl4 boils at 77 0 C 6) Solubility in polar and non-polar solvents: ionic compound are generally more soluble in polar (H 2 O, NH 3 ) than in non polar solvents(C 6 H 6, CCl 4 ). The electrostatic force of attraction in ionic solid is reduced by the high value of dielectric constant

12. of the polar solvent ions can move freely and interact with solvent molecules to form the solvated ions. Limitations of ionic bond Electrovalency: Electrovalency of an element is equal to the number of electrons lost by an atom of that element in forming a (+) ion or gained by it in forming a negative ion, both having the noble gas configuration (S 2 P 6 ). Varible electrovalency 1) Unstable configuration Loss of electron from Fe, Co, Ni, Cu remaining part is called the core (unstable) further loss of electron from the core variable electrovalency

13. e.g. Fe (26) 2,8,14,2 or 2,8, 3S 2 3P 6 3d 6 4S 2 Fe +2 (24) 2,8,14 or 2,8, 3S2 3P6 3d6 Fe +3 (23) 2,8,13 or 2,8, 3S2 3P6 3d5 2) Many cations do not always have inert gas structure. There are cations which have outer S 2 P 6 d 10 CU+ Zn+2 3S 2 3P 6 3d 10 Ag+ cd+2 4S 2 4P 6 4d 10 Cu 1S 2 2S 2 2P 6 3S 2 3P 6 3d 10 4S 1 Cu+ 1S 2 2S 2 2P 6 3S 2 3P 6 3d 10 Cu++ 1S 2 2S 2 2P 6 3S 2 3P 6 3d 9

14. 3) Inert Electron Pair Effect: Is simply the reluctance of ns orbital electron pair to take part in bonding. e.g. in case of p-block elements the valency of +3 becomes more predominant than +5 as the ns electrons do not participate in bonding. Example: although the common oxidation sate for elements in group four is +4, most elements in the group can also exist in oxidation state +2.This is because of the inert pair effect. In large atoms, such as those of tin and lead some outer shell electrons are not as well shielded as those in the inner core. They are therefore sucked into the inner core of electrons and thus become inert. Group III A (G=3) Ga (+1, +3), In (+1, +3) TI (+1, +3) Group IV A (G=4) Ge (+2, +4), Sn (+2, +4) Pb (+2, +4) Group V A (G=5) As (+3, +5), Sb (+3, +5) Bi (+3, +5)

15. The higher oxidation state for most of the elements is equal to their group number, G while the lower one is equal to (G- 2). G when all the ns and np electrons from ns 2 p x configuration of p block elements(x= 1, 2,3 and 4 for the elements of gr. IIIA, IVA, VA and VA, respectively) are lost. (G-2) when only the np electrons are lost and the ns electron pair, due to its extra stability, remains inert, i.e. not lost. Such a pair of ns electrons is called inert electron pair and the effect caused by it is known as ‘Inert electron pair effect”.

16. Dielectric constant: When a dipolar or polar molecule is placed in an electrostatic field the positive end of the dipole is directed towards the negative pole of the electric field and vice versa. The oriented dipole oppose the field and the intensity of the electric field is reduced. The reduced tendency is measured by the dielectric constant or the specific inductive capacity of the polar substance. The substances containing ionic bonds have high dielectric constant.

17.  General Properties of Ionic Compounds  Ionic compounds exist as crystals.  They are solid in room temperature.  The crystals are hard and brittle.  They have high melting points.  They are soluble in water.  In molten state and aqueous solution, they are good conductor of electricity.  The ionic compounds do not show isomerism.  The reactions between the ionic compounds take place with great speed.

18. CHEMICAL BOND (Continued) Covalent Bond  Definition  Covalent bond (also called electron-pair bond) may be defined as the chemical bond or attractive force between atoms that results from sharing of an electron-pair. Each of the two bonding atoms contributes one electron to the electron-pair (and has equal claim on the shared electron-pair). The shared electron pair is indicated by a dash (  ) between the two bonded atoms.  To state in a different way, a covalent bond between two atoms results from the overlap of an orbital of one atom with an orbital of another atom. As the two orbitals overlap, they share the same region in space and a new orbital (molecular orbital) is formed.  Two atoms may bind together by one, two or even three covalent bonds.  The compounds containing a covalent bond are known as covalent compounds.

19.  Conditions for Formation of Covalent Bond: The conditions favorable for the formation of covalent bonds are:  Number of valence electrons: Each of the atoms A and B should have 5, 6 or 7 electrons (1 for H) so that both can achieve stable octet by sharing 3, 2 or 1 electron-pair respectively. The non-metals of the groups VA, VIA and VIIA respectively satisfy this condition.  Equal electronegativity: The two atoms should have almost equal electro negativity (in other words, equal attraction for electrons).  High Ionization energy: The atoms with high ionization energy are incapable of loosing electron to form cations but they can form covalent bonds between them  High nuclear charge and small intermolecular distance

20. CHEMICAL BOND (Continued)  General Properties of Covalent Compounds  Physical state: Usually gases, liquids or relatively soft solids at room temperature.  Crystal Structure: a.Giant molecules – where every atom is bonded with other atoms by covalent bonds. E.g: diamond b.Separate layers – the covalent compounds containing separate layers are said to have layer lattice structure.e’g: graphite  Melting point & Boiling Point: Low melting points or boiling points.  Neither hard nor brittle: Covalent compounds are soft and waxy as they contain separate crystal lattice. No repulsion forces are present between the layers so the covalent crystals are easily broken

21.  Solubility : Usually soluble in nonpolar organic solvents (e.g. benzene, ether) and insoluble in water.  Conductivity: Non-conductor of electricity.  Isomerisom: Exhibit isomerism.  Molecular reactions: since there are no strong electrical forces acting in covalent compounds to speed up the reaction between molecules, the molecular reactions are slow.  General Properties of Covalent Compounds

22. CHEMICAL BOND (Continued)  Types of Covalent Bonds: In terms of the molecular orbitals formed, there are two main types of covalent bonds: Sigma (  ) bonds and pi (  ) bonds. Sigma Bonds  A sigma bond is formed by linear (end-to-end) overlap of orbitals.  All single covalent bonds and one bond in multiple covalent bonds are sigma bonds.  It may be obtained by:  the overlap of two s orbitals  the overlap of a p orbital and an s orbital, or  the overlap of two p obitals. Pi Bonds  A pi bond is formed by parallel or (side-by-side) overlap of p orbitals.  A pi bond has two lobes like p orbitals  one half of the bond lies above the plane containing the two nuclei and the other half lies below the plane.  One bond in double bonds and two bonds in triple bonds are pi bonds.

23. CHEMICAL BOND (Continued) Sigma bondPi bond 1It is formed by end-to-end overlapping of half-filled atomic orbitals. It is formed by the sidewise overlapping of half-filled p obitals. 2Overlapping takes place along the inter-nuclear axis. Overlapping takes place perpendicular to the inter-nuclear axis. 3The extent of overlapping is large and the bond formed is stronger. The extent of overlapping is smaller and the bond formed is weaker. 4There is free rotation around the sigma bond and so no geometrical isomerism is possible. There is no free rotation about the pi bond and so geometrical isomerism possible. 5Both s and p orbitals can participate in sigma bond formation. Only p orbitals participate in the formation of pi bonds.  Differences between the Sigma and Pi bonds

24. CHEMICAL BOND (Continued)  Differences/ Distinctions between Ionic Bonds and Covalent Bonds CharacteristicsIonic BondCovalent Bond 1. Formation of Bond Formed by transfer of electrons from a metal to a nonmetal atom. Formed by sharing of electrons between non-metal atoms. 2. Physical StateIonic compounds are solids at room temperature. Covalent compounds are gases, liquids or soft solids. 3. Hardness & Brittleness Ionic compounds are hard and brittle.Covalent compounds are soft and are much readily broken. 4. SolubilityIonic compounds are soluble in water but insoluble in organic solvents. Covalent compounds are insoluble in water but soluble in organic solvents. 5. Melting & Boiling Points Ionic compounds have high melting and boiling points. Covalent compounds have low melting and boiling points. 6. Conduction of Electricity Ionic compounds conduct electricity in molten state or aqueous solution. Covalent compounds do not conduct electricity. 7. IsomerismIonic compounds do not show isomerism. Covalent compounds can show isomerism. 8. ReactivityReactions of ionic compounds are fast. Reactions of covalent compounds are slow.

25. Nonpolar & Polar Covalent Bonds  Nonpolar Covalent Bond:  This is a covalent bond in which the electron pair is shared equally by the linked atoms.  It happens when the two atoms are similar and so have the same electronegativity or zero electronegativity difference.  This is also called homopolar covalent bond or simply covalent bond. As a matter of fact, it is the pure or true covalent bond.

26.  Polar covalent bond:  This is a covalent bond in which the electron pair is shared unequally by the linked atoms so that they acquire a partial positive and negative charge.  This happens when the covalent bond links two dissimilar atoms which have different electronegativity values and so unequal attraction for the electron pair.  The molecules which have polar covalent bonds are known as polar molecules. Examples of polar molecules include HCl, H 2 O, NH 3 etc.  A polar molecule, with its positive and negative charge centers or electric poles at the ends of the covalent bond, becomes dipolar and hence is called a dipole (two poles). The dipole of a bond may be shown by an arrow from positive to negative with a crossed tail.

27. Coordinate Bond  Definition  Coordinate bond is a covalent bond, which is formed by the mutual sharing of two electrons both of which are provided by one of the linked atoms (or ions). Coordinate bond is also sometimes referred to as coordinate covalent bond or dative bond.  If an atom has an unshared pair of electrons (lone pair) and another atom is short of two electrons than the stable number, a coordinate bond is formed. The atom which donates the lone pair is called the donor and the atom which accepts it the acceptor. The coordinate bond is represented by an arrow which points away from the donor to the acceptor.  The compounds containing a coordinate bond are called coordinate compounds and the molecule or ion that contains the donor atom is called the ligand.

28.  Illustration of formation of a coordinate bond: The formation of covalent bond between two atoms, say A and B, can be illustrated as follows: 1) The atom acting as a donor should have a lone pair of electrons. 2) The atom acting as an acceptor should have a vacant orbital to accept the electron pair donated by the donor.  Conditions for the formation of a coordinate bond Conditions necessary for the formation of a coordinate bond are:

29. Characteristics of coordinate covalent compound: (1)Their melting and boiling points are higher than purely covalent compounds and lower than purely ionic compounds. (2) These are sparingly soluble in polar solvent like water but readily soluble in non-polar solvents. (3) Like covalent compounds, these are also bad conductors of electricity. Their solutions or fused masses do not allow the passage to electricity. (4) The bond is rigid and directional.

30.  Some examples of coordinate compounds or ions NH 4 +, H 3 O +, BF 4 , addition compound of NH 3 with BCl 3, CH 3 NO 2, SO 2 & SO 3, Al 2 Cl 6, SO 4 2 , O 3, CO.  Representation of Coordinate Bond Formation by Lewis Symbol and Structure

31. Metallic Bond Metallic bond is the attractive force, acting between the positive metal ions and surrounding freely mobile electrons, that holds the metal atoms together in a metal crystal.

32. Hydrogen Bond  Definition: Hydrogen bond refers to the electrostatic attraction between (i) a H atom covalently bonded to a highly electronegative atom X (O, N or F) and (ii) a lone pair of electrons on X in another molecule  Types of Hydrogen Bond: Hydrogen bond is of two types: i) Intermolecular hydrogen bond (association) & ii) Intramolecular hydrogen bond

33.  Intermolecular Hydrogen Bond: When hydrogen bonding occurs between two molecules of the same or different compounds, it is called intermolecular hydrogen bonding. e.g. hydrogen bonding in between the molecules of H 2 O, NH 3 or HF.  Intermolecular Hydrogen Bond: If hydrogen bonding takes place between an H-atom and an electronegative atom within the same molecule (e.g. in o-nitrophenol), it is referred to as intramolecular hydrogen bonding.

34.  Definition: Van der Waals forces are very short-lived inter-molecular attractive forces which are believed to exist in between all kinds of atoms, molecules and ions when they are sufficiently close to each other.  Types: There are four types of van der Waals forces. These are:  Dipole-dipole interactions (Keesom forces): These are found in polar molecules (e.g. HCl) and are the strongest of all van der Waals forces.  Ion-dipole interactions: e.g. between Na + / Cl  and H 2 O when NaCl dissolves in H 2 O.  Dipole-induced dipole interactions (Debye forces): These are found in a mixture containing polar and non-polar molecules.  Instantaneous dipole-induced interactions (London forces): These are found in nonpolar molecules, e.g. diatomic gases like H 2, O 2, Cl 2, N 2 etc. and monoatomic noble gases like He, Ne, Ar etc. Van der Waals Forces

35. CHEMICAL BOND (Continued)

37. HYBRIDIZATION  Definition:  Hybridization may be defined as the phenomenon of mixing up (or merging) of pure orbitals of nearly equal energy on an atom, giving rise to equal number of entirely new orbitals, which are equal in energy and identical in shape.  It may be noted here that it is the orbitals that undergo hybridization and not the electrons. For example, four orbitals of an oxygen atom (1s 2, 2s 2, 2p x 2, 2p y 1, 2p z 1 ) belonging to second level (i.e. 2s, 2p x, 2p y, 2p z ) can hybridize to give four hybrid orbitals  two of which have two electrons (as before) and the other two have one electron each (as before).  Importance of the concept of hybridization: The concept of hybridization is important for explaining the tendency of atoms like Be, B and C to form bonds and the shape or geometry of their molecules.  Types of Hybridization: Depending upon the number and nature of the orbitals undergoing hybridization, we have various types of hybrid orbitals. For instance, s, p and d orbitals of simple atoms may hybridize in the following manner: (a) sp hybridization, (b) sp 2 hybridization, (c) sp 3 hybridization and (d) Hybridization involving d orbitals, e.g. dsp 3, d 2 sp 3, dsp 2

38. HYBRIDIZATION (Continued) It involves mixing of an s and a p orbital giving rise to two hybrid orbitals known as sp hybrid orbitals.  The resulting orbitals arrange themselves along a line at an angle of 180  (between the axes of the two orbitals) and are, therefore, often referred to as linear hybrid orbitals. sp hybridization

39. HYBRIDIZATION (Continued)  Examples of compounds undergoing sp hybridization: BeF 2, ethynes (e.g. CH  CH)

40. HYBRIDIZATION (Continued) sp 2 hybridization  It involves mixing of an s and two p orbitals giving rise to three new orbitals called sp 2 hybrid orbitals.  The sp 2 hybrid orbitals resemble in shape with that of sp hybrid orbitals but are slightly fatter.  The three orbitals are arranged in the same plane at an angle of 120º to one another and hence sp 2 hybrid orbitals are also called trigonal hybrid orbitals and the process trigonal hybridization.

41. HYBRIDIZATION (Continued)  Example of molecules undergoing sp 2 hybridization: BF 3, ethenes (e.g. CH 2  CH 2 )

42. HYBRIDIZATION (Continued)

43.  It involves mixing of an s and three p orbitals of an atom, giving rise to four new orbitals called sp 3 hybrid orbitals.  They are of the same shape of the previous two types, but bigger in size.  The four orbitlas are arranged in the space in the form of a regular tetrahedron at angle of 109.5° between them, and because of their tetrahedral disposition, this type of hybridization is also called tetrahedral hybridization. sp 3 hybridization

44. HYBRIDIZATION (Continued)  Examples of molecules undergoing sp 3 hybridization: CH 4, NH 3, H 2 O

45. HYBRIDIZATION (Continued)

46. Bond Angle : It is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion. Bond angle is expressed in degree which can be experimentally determined by spectroscopic methods.