CHEMICAL BONDING AND SHAPES OF MOLECULES

CHEMICAL BONDING AND SHAPES OF

MOLECULESSANJAYA DWARE

CHEMISTRY DEPARTMENTLITTLE ANGELS’ COLLEGE

Chemical Bonding by SD 12021/09/01

Syllabus– Show structure atoms and ions by Lewis dot method.

– Explain the ionic bond and the properties of ionic compounds.

– Explain the covalent bond, co-ordinate bond and the properties of covalent compound.

– Describe the feature of sigma and Pi-bond

– Describe the co-ordinate covalent compounds with some examples.

– Write the lewis dot diagrams of some ionic and covalent compounds (NaCl, MgCl2, NH4Cl, Oxides of Hydrogen, Nitrogen and Phosphorous, common mineral acids).

– Write the resonance structure of some covalent species.

– Explain the properties of molecular and metallic solids on the basis of vanderwaal’s and metallic bonding.

– Use VSEPR theory to describe the shapes of simple covalent molecules.

– Describe the concept of hybridization in simple covalent molecules.

– Explain the characteristics of bond in terms of dipole moment, Ionic character and bond length.

– Describe the hydrogen bondng and outline the importance of hydrogen bonding to the physical properties of substances, including ice and water (for example, boiling and melting points, viscosity, surface tension and solubility)

Terminology

■ Valence shells

■ Valence electrons

■ Lewis symbol

Types of bond

■ Electrovalent bond

■ Covalent bond

■ Coordinate covalent bond

■ Hydrogen bond

■ Metallic bond

■ Van der waal’s bond

Electrovalent bond or ionic bond

■ The bond formed between two or more than two atoms by the transfer of one or

more electrons among atoms

■ Loss of electrons by an atom of an element forms cation and gain of electrons by an

atom forms anion

■ Cations and anions are held together by an electrostatic force of attraction.

■ The number of electrons lost or gained by an atom of an element is called

electrovalency of the element.

Conditions for the formation of ionic bond

■ Ionisation energy: Energy required to remove the most loosely held electron from an

isolated gaseous atom in its ground state is called ionization energy of that element.

Low ionization energy favors ionic bond.

■ Electron affinity: Energy released when a neutral gaseous atom gains electron

giving an anion is called electron affinity.

High electron affinity favors ionic bond.

■ Lattice energy: Energy released when one mole of the crystal of the compound is

formed from its ions.

Higher lattice energy favors ionic bond.

Formation of ionic bonds

■ Sodium chloride

■ Magnesium chloride

■ Sodium oxide

Characteristics of ionic compounds

■ Physical state

■ Solubility

■ Electrical conductivity

■ Brittleness

■ Melting and boiling point

■ Ionic bond is non-directional

COVALENT BOND

■ Definition: The bond formed by mutual sharing of electrons between combining

atoms of same or different elements

■ Covalency: The number of electrons which its atom contributes for sharing during

the covalent bond formation.

■ Conditions for the formation of covalent bond:

– Electron affinity: elements having nearly equal electron affinity

– Electronegativity: the combining elements having low electronegativity

difference.

– Ionization potential: elements having high ionization potential form covalent

Electronegativity is defined as a power of an element to attract electrons towards itself in a molecule.

Formation of covalent bond

■ H2

■ O2

■ Cl2

■ N2

■ NH3

■ HCl

■ CO2 , H2O, CH4 , C2H4 , CH3OH , H2S

HNO3, H2SO4

CO3-- , NH4

+ , SO4--, NO3

Properties of covalent compounds

■ Physical state: exist in all three states i.e. solid, liquid and gases.

■ Melting and boiling point: low melting and boiling point.

■ Electrical conductivity: donot conduct electricity as they donot give charged particles

in solution.

■ Solubility: soluble in nonpolar solvents. Eg. Benzene, carbon tetrachloride

■ Directional nature: covalent bonds have directional nature thus have particular

geometrical shape.

Polarity of a covalent bond:Ionic character of a covalent bond

■ In homonuclear diatomic molecules (H2 , O2 , Cl2 etc. ) the shared electrons are equally distributed between the participating atoms. Such.bonds are called non-polar bonds. Eg. H-H, Cl-Cl, F-F etc.

■ When covalent bond is formed between the atoms of different elements, the atoms of the element having higher electronegativity attracts the shared electrons more towards itself and thus the polarity develops in the bond. Such bonds are called polar covalent bonds. Eg. HCl, HBr, HI etc.

■ Electronegativities in Pauling Scale:

H = 2.1, F = 4.0 , Cl = 3.0, Br = 2.8, I = 2.5

H F H Cl, H Br, H I

Electronegativities diff: 1.9 0.9 0.7 0.4

Polarity of H----X bond decreases

Dipole moments

■ Definition:

The product of charge (q) on either of the atoms and distant(d) between them is called dipole moment.

It is denoted by µ.

Mathematically,

µ = q x d

where, q = charge on any one atom,

d = distance between the atoms

Dipole moment is a vector quantity.

Unit: Debye (D)

Applications of dipole moment

■ In the calculation of percentage ionic character

■ In the determination of the polarity of bonds: Greater is the magnitude of dipole

moment of the molecule, higher will be the polarity of the bond.

– If µ = 0, the molecule is non – polar, and

– If µ = 0, the molecule is polar.

■ In the determination of shape of the molecules.

Knowing the dipole moment, the shape of the

molecule can be predicted.

Co-ordinate covalent bond or dative bond

■ Definition:

The bond formed by the sharing of electrons between combining atoms, in which both

the shared electrons are contributed by one of the bonded atoms only.

■ The atom which donates the shared pair of electrons is called a donor atom, and the

atom which accepts the shared pair of electrons is called an acceptor atom.

■ Conditions for the formation of a coordinate bond

– The donor atom after the formation of other bonds must be in octet state and

must possess a lone pair of electrons.

– The acceptor atom must require two electrons to attain its nearest noble gas

configuration.

Formation of coordinate bonds

■ O3 , SO2 , SO3

■ NH4+

■ H3O +

HNO3, H2SO4, H3PO4 , HClO3

CO3-- , NH4

+ , SO4--, NO3

Na2 SO4 , K NO3 , CaCO3

Formation of coordinate bonds

■ H2SO3,

■ H3PO3 ,

■ SO2Cl2 ,

■ HIO3 ,

■ CH3COONa

FAILURE OF OCTET RULE

■ Molecules having incomplete octet: BeCl2 , BF3 , AlCl3 , BCl3

■ Molecules having super octet structures : PCl5 , IF7

HYDROGEN BONDING

■ Definition:

The force of attraction between hydrogen bonded with highly electronegative element (F

or O or N) and highly electronegative element (F or O or N) of same or different molecule

is called hydrogen bond.

Example: water, ammonia, HF, methanol

It is represented by a dotted line.

H----------F- - - - - - - H-------------F - - - - - -H----------Fδ+ δ--

Hydrogen bond

δ--δ--δ+ δ+

Note: Bond energy of hydrogen bond is 2-10kcal/mol whereas bond energy of covalent bond is 50-100kcal/mol.

TYPES OF HYDROGEN BOND

■ INTERMOLECULAR HYDROGEN BOND:

■ Definition:

■ The force of attraction between hydrogen bonded with highly electronegative

element (F or O or N) and highly electronegative element (F or O or N) of different

molecule is called intermolecular hydrogen bond.

■ Examples: Water, water with methanol, methyl alcohol

TYPES OF HYDROGEN BOND

■ INTRAMOLECULAR HYDROGEN BOND: The force of attraction between hydrogen

bonded with highly electronegative element (F or O or N) and highly electronegative

element (F or O or N) of same molecule is called intramolecular hydrogen bond.

o-hydroxybenzoic

O-nitrophenol has lower boiling point than p-nitrophenol. Why?

Intermolecular

Consequences of hydrogen bonding

■ Water is liquid but H2S is gas at room temperature.

■ Glycerol is more viscous than water.

■ Boiling point of o-nitrophenol is less than that of p-nitrophenol.

■ Decrease in volume during melting of ice.

■ Unusual melting and boiling points of hydrides of group VA, VIA and VIIA

Unusual melting and boiling points of hydrides of group VA, VIA and VIIA

Consequences of hydrogen bonding

Van der Waals’ Force

■ Definition:

The short lived attractive force existing between all kind of atoms, molecules and ions,

when they come closer, is called van der Waals’ force.

In 1873, J.D. van der Waals pointed out the force of interaction among the non polar

molecules and atoms.

In 1930, it was explained by Fritz London that the van der waals interaction arises due

to unsymmetrical distribution of electron cloud motion of the electrons around the

nucleus.

Also called as London Forces.

Metallic Bond

■ Definition:

The simultaneous force of attraction between the mobile electrons and the positively charged core is responsible for holding metal atoms together, which is known as metallic bond.

Electron pool theory■ Metallic lustre

■ Electrical conductivity

■ Malleability and ductility

■ Hardness

■ Thermal conductivity.

Resonance

■ Definition:

The phenomenon in which a molecule or ion cannot be represented be a

single structure rather it should be represented by more than one

structures to explain its properties, is called resonance.

Examples:

Benzene

Resonating structures of Benzene

Bonding in solid state

■ Ionic solids

■ Covalent solids

■ Molecular solids

■ Metallic solids

Thank you

34Chemical Bonding by SD2021/09/01

CHEMICAL BONDING AND SHAPES OF MOLECULES

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