1. CHEMICAL BOND Dr. M. Mahboob Hossain Associate Professor, Department of MNS BRAC UNIVERSITY PRESENTED BY

2. VALENCE  Valence:  Valence is the combining capacity of an atom or radical with another atom or radical.  It is determined by the number of electrons that the atom or radical loses, gains or shares when it reacts with other atoms or radicals.  For example, the valence of sodium is +1, because sodium tends to lose one electron easily to become sodium ion (Na + ). Similarly, the valence of chlorine is –1, because chlorine has a tendency to gain one electron to become chloride ion (Cl  ).  Valence Shell: It is the outermost shell (energy level) of an atom.  Valence Electrons:  The electrons present in the outermost shell (energy level) of an atom is known as valence electrons (because it is these electrons that determine the manner in which atoms combine with each other).  For example, the electronic configuration of Na is 2, 8, 1 and that of Cl is 2, 8, 7. Thus, sodium has one valence electron and chlorine 7.

3. VALENCE (Continued)  Bonding and Non-bonding Electrons: The valence electrons actually involved in bond formation are called bonding electrons. The remaining valence electrons still available for bond formation are known as non-bonding electrons.  Lewis Symbols of an Element and a Compound  The valence electrons can be shown by an equal number of dots (  ) or crosses (×) around the symbol of the element. The bonding electrons are shown in appropriate positions and the non-bonding electrons are generally given in pairs. This type of symbol of an element is known as the Lewis symbol of the element. For example, the Lewis symbols for hydrogen, chlorine, oxygen and sulfur may be written as follows:  The structural formulae of compounds built by union of Lewis symbols of the component atoms are known as Electron-dot formulas, Electron-dot structures or Lewis structures. For example, the structural formulae for NaCl, H 2 O and NH 3 can be shown by Lewis structures as follows:

4. VALENCE (Continued)  Electronic Theory of Valence  In 1916, G.N. Lewis and W. Kossel gave a theory to explain ‘why atoms join to form molecules’. This is known as The Electronic Theory of Valence. It states that: In chemical bond formation, atoms interact by losing, gaining, or sharing electrons so as to acquire a stable noble gas configuration in their outermost shell.  Each noble gas, except helium, has a valence shell of eight electrons. Most other elements have either more or less than eight electrons in their valence shell and they tend to assume stable eight-electron configuration of noble gases in their outermost shell by losing, gaining or sharing electrons. This tendency of atoms to have eight electrons in the outermost shell is known as the Octet Rule.  Since hydrogen and lithium have only one and three electrons respectively, they tend to get two electrons in their outermost shell to assume the electronic configuration of the noble gas helium.  Exceptions to the Octet Rule: Certain atoms may have fewer or more than eight electrons in their valence shells when they bond to form molecules.  In BF 3, boron has only six electrons in its valence shell.  Phosphorus and sulfur sometimes can expand their octets to 10 or 12 electrons. For example: in PCl 5 (gaseous state) and SF 6, phosphorus and sulfur have 10 and 12 electrons respectively in their valence shells.

5. CHEMICAL BOND  Chemical Bond It is the force that acts between two or more atoms to hold them together in a molecule or ion.  Causes of Chemical Bond Formation Atoms form chemical bond for the following reasons:  A net attractive force between the bonding atoms  A desire for attaining stable electron configuration of noble gases (8 or 2 electrons in the outermost shell)  A desire for having a lower (overall) energy

6. CHEMICAL BOND (Continued)  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.

7. 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.

8. CHEMICAL BOND (Continued)  The compounds containing such bonds are known as ionic or electrovalent compounds.  In general, metallic atoms tend to lose electrons to form cations and the non- metallic atoms tend to gain electrons to form anions.  Conditions for Formation of Ionic Bonds The conditions favourable for the formation of an ionic bond are:  Number of valence electrons: The atom A (i.e. the donor atom) should possess 1, 2 or 3 valence electrons and the atom B (i.e. the acceptor atom) should have 5, 6 or 7 valence electrons. (The elements of group IA, IIA and IIIA satisfy this condition for the atom A and those of groups VA, VIA and VIIA satisfy this condition for the atom B.)  Net lowering of energy: There must be a net release of energy as a result of electron transfer and formation of ionic compound.  Electronegativity difference between A and B: There should be a large electronegativity difference (2 or more) between A and B.

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

10. CHEMICAL BOND (Continued)  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.

11. 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.

12. CHEMICAL BOND (Continued)  Conditions for Formation of Covalent Bond The conditions favourable 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 electronegativity (in other words, equal attraction for electrons).

13. CHEMICAL BOND (Continued)  Some Examples of Covalent Compounds:  Simple covalent compounds: H 2, Cl 2, HCl, H 2 O, NH 3, CH 4

14. CHEMICAL BOND (Continued)  Multiple covalent compounds: O 2, N 2, CO 2

15. CHEMICAL BOND (Continued)  General Properties of Covalent Compounds  Usually gases, liquids or relatively soft solids at room temperature.  Low melting points or boiling points.  Neither hard nor brittle.  Usually soluble in nonpolar organic solvents (e.g. benzene, ether) and insoluble in water.  Non-conductor of electricity.  Exhibit isomerism.  Molecular reactions are slow.

16. 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.

17. 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

18. 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 State Ionic 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. Solubility Ionic 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. Isomerism Ionic compounds do not show isomerism.Covalent compounds can show isomerism. 8. Reactivity Reactions of ionic compounds are fast.Reactions of covalent compounds are slow.

19. CHEMICAL BOND (Continued) 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.

20. CHEMICAL BOND (Continued)  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.

21. CHEMICAL BOND (Continued) 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.

22. CHEMICAL BOND (Continued)  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:

23. CHEMICAL BOND (Continued)  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

24. CHEMICAL BOND (Continued) 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.

25. CHEMICAL BOND (Continued) 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

26. CHEMICAL BOND (Continued)  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.  Intramolecular 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.

27. CHEMICAL BOND (Continued)  Association of molecules  Abnormally high boiling and melting points (of some molecules e.g. HF, H 2 O and NH 3 )  High solubility of some covalent compounds (e.g. NH 3 and CH 3 OH) in certain hydrogen containing solvents (e.g. H 2 O).  Three dimensional crystal lattice (e.g. of H 2 O in solid state)  High viscosity, high heat of vaporization and high dielectric constant.  Physical Properties Attributed to Hydrogen Bonding

28. CHEMICAL BOND (Continued)  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

29. CHEMICAL BOND (Continued)

31. 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

32. 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

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

34. 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.

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

36. HYBRIDIZATION (Continued)

37.  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

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

39. HYBRIDIZATION (Continued)

40. Sugden s concept of single linkake Sugden postulated that the central atom of the molecule like PCl 5 maintains its octet and in doing so the central atom is linked with some of the combining atoms by single electron bonds called singlet linkages while with the remaining atoms it is linked by normal two electron bond.

42. Variable covalnency ElementsDifferent covanlencies Examples of compouds Phosphorous3535 PCl 3 PCl 5 Sulphur246246 SCl 2 SF 4

44. REFERENCES  Bahl and Tuli (Physical)  Madan (Inorganic)