8 Covalent Bonding 8.1 Formation of Covalent Bonds 8.2 Dative Covalent Bonds 8.3 Bond Enthalpies 8.4 Estimation of Average Bond Enthalpies using Data from Energetics
8.5 Use of Average Bond Enthalpies to Estimate the Enthalpy Changes of Reactions 8.6 Bond Enthalpies, Bond Lengths and Covalent Radii 8.7 Shapes of Covalent Molecules and Polyatomic Ions 8.8 Multiple Bonds 8.9 Covalent Crystals
8.1 Formation of Covalent Bonds
Theory of covalent bond formation • Lewis Model (A classical treatment) • Valence Bond Theory (pp.31-35) • (by Linus Pauling) • Molecular Orbital Theory (pp.36-37) • (By Robert Mulliken)
1. Lewis Model A covalent bond is formed by sharing of valence electrons between atoms of non-metals. Reasons : 1. They have small/no difference in E.N. 2. They can achieve stable electronic configurations. (octet /duplet structures) 3. The total energy of the system is lowered when electrons are shared between the two nuclei.
nucleus electron Force between the electron and nucleus. Force along the bond axis. A B If the electron is located between two nuclei, there is an attractive force holding the two nuclei together.
A B A’ B’ BB’ > AA’ If the electron is not located between two nuclei, there is a net force separating the two nuclei.
1. Lewis Model A Covalent bond is the result of shared valence-electron pair(s) positioned between two nuclei. Sharing of electron pair(s) between two nuclei.
H H C H H H H C H H H H C H H Lewis(dot and cross) structure and Bond-line structure • Electrons are represented by dots and/or crosses. • Lewis structures basically obey octet rule. • Hydrogen adopts duplet structrue.
H H C H H Bond pairs H O H Lone pairs Lewis(dot and cross) structure and Bond-line structure Only lone pairs are shown as dots/crosses in bond-line structure
O O N N Multiple bonds Double bond Triple bond
Rules for drawing Lewis structures – A systematic way Using CO2, CO32 and HClO4 as examples • Determine the arrangement of atoms within a molecules/polyatomic ion • Central least electronegative • Terminal H & F (most electronegative)
O C O O C O O O O Cl O O H CO2 CO32 HClO4
2. Determine the total no. of valence electrons(V) in the species. For neutral molecule, V = sum of group numbers of atoms = m For anion, An, V = m + n For cation, Cn+, V = m – n
Species CO2 CO32 HClO4 V 4+26 =16 4+36+2 =24 1+7+46 =32 W W – V
Calculate the total no. of electrons that would be needed if each atom obeys the octet rule(W). • 2e for H • 8e for other atoms • 4. The difference (W – V) gives the no. of electrons that have to be shared such that all atoms in the species can obey octet rule.
Species CO2 CO32 HClO4 V 4+26 =16 4+36+2 =24 1+7+46 =32 W 38=24 48=32 2+58 =42 W – V 24-16 =4 pairs 32-24 =4 pairs 42-32 =5 pairs
O C O O C O O CO2 HClO4 O O CO32 Cl O O H • Assign a single bond to each terminal atom in the species.
HClO4 O O Cl O O H • Assign a single bond to each terminal atom in the species. If the resulting central atom has more than 4 single bonds, rearrange the terminal atoms such that the central atom has 4 single bonds.
Assign a single bond to each terminal atom in the species. • If the resulting central atom has more than 4 single bonds, rearrange the terminal atoms such that the central atom has 4 single bonds. • In SF6, W – V = 78 – (6+67) = 8 • 4 single bonds are required However, S has to form 6 single bonds. Expansion of octet structure may happen for elements in Period 3 and beyond.
CO2 6. If bonding electrons remain, assign them in pair by making some of the bonds double or triple bonds. Two bond pairs left 3 ways
CO32 One bond pair left 3 ways
HClO4 No bond pair left Only one way
CO2 • Assign lone pairs to terminal atoms to give them octet. • If any electrons still remain, assign them to the central atoms as lone pairs. No. of lone pairs left = (16 – 8)/2 = 4
CO32 No. of lone pairs left = (24 – 8)/2 = 8
HClO4 No. of lone pairs left = (32 – 10)/2 = 11
Assume bond pairs are equally shared • Determine the formal charge on each atom. • Formal charge is the charge an atom in a species would have if the bonding electrons are shared equally between the atoms. Formal charge of an atom in a species
• Separation of opposite formal charges • Increase in potential energy of the species • less stable Lewis structures Most stable Lewis structure
There is no separation of opposite formal charges All three structures are stable
2 2 2 CO32
3+ HClO4 Substantial separation of opposite formal charges Very unstable
HClO4 Lone pair electrons must be rearranged to minimize the separation of opposite formal charges.
HClO4 Cl can expand the octet structure to accommodate 14 valence electrons.
Breakdown of the Octet Rule (p.30) • Valence shell expansion (Q.21(g), (h) (k) to (n)) • For elements from Period 3 and beyond can expand the octet by utilizing low-lying d-orbitals in bond formation. Max. no. of bond pairs = Group no. Gp 5 PF5 Gp 6 SF6 Gp 7 IF7
Breakdown of the Octet Rule (p.30) 2. Electron-deficient species (Q.21(o)) Dative bond (coordinate covalent bond) is a covalent bond in which the bond pair electrons are contributed solely by one of the bonding atoms.
The arrow( )is pointing from the electron donor from the electron acceptor. Breakdown of the Octet Rule (p.30) 2. Electron-deficient species (Q.21(o))
unstable Separation of opposite formal charges Not favourable for the most electronegative F to carry a positive formal charge
Q.21(d) Separation of opposite formal charges is unavoidable. Separation of opposite formal charges can be avoided by expansion of octet
Q.21(a) Once formed, dative covalent bond cannot be distinguished from normal covalent bond. The four N – H bonds are identical.
Less stable Breakdown of the Octet Rule (p.30) 3. Odd-electron species Q.22 Separation of opposite formal charges Positive formal charge on the more electronegative atom