CHEMISTRY 161 Chapter 9 Chemical Bonding

Periodic Table of the Elements ns2npx ns1 ns2 chemical reactivity - valence electrons

THE OCTET RULE ns2np6 atoms combine to form compounds in an attempt to obtain a stable noble gas electron configuration ns2np6 Iso electronic

A + B → AB 1. ELECTRON FULLY TRANSFERED IONIC BONDING 2 Na(s) + Cl2(g)  2 NaCl(s) 2. ELECTRON SHARING COVALENT BONDING 2 H2(g) + O2(g)  2 H2O(l)

represents one valence electron LEWIS MODEL OF BONDING LEWIS DOT SYMBOL DOT represents one valence electron H. Gilbert Lewis (1875-1946)

. . . . . . . . . . with the exception of He, the main group number represents number of ‘dots’ only valence electron are considered

IONIC BONDING Na electron transfer Ne core implied in symbol 1s22s22p63s1 Lewis Symbol

Cl Na Ne core implied in symbol 1s22s22p63s23p5 1s22s22p63s1 Lewis Symbol

IONIC BONDING  Cl Cl Na Na+ the formation of ionic bonds is represented in terms of Lewis symbols Cl Cl Na  Na+ 1s22s22p63s23p6 1s22s22p6 the loss or gain of electrons(dots) until both species have reached an octet of electrons

represents one orbital Cl Cl [Ne] 3s23p6 represents one orbital (Pauli: 2 electrons)

ions stack together in regular crystalline structures electrostatic interaction ionic solids typically 1. high melting and boiling points 2. brittle 3. form electrolyte solutions if they dissolve in water

Li(s) + ½ F2(g) → LiF(s) enthalpy of formation lattice energy (up to few 1000 kJmol-1) Li+(g) + F-(g) → LiF(s) Hess’s Law Born-Haber Cycle

Li(s) + ½ F2(g) → LiF(s) Li+(g) + F-(g) Li(g) + F(g) Li(s) + ½ F2(g) 5 ΔHoR= Σ ΔHoi i=1 ΔHo4 ΔHo3 ΔHo5 Li(g) + F(g) ΔHo1 ΔHo2 ΔHoR Li(s) + ½ F2(g) LiF(s)

Mg(s) + ½ O2(g) → MgO(s) Mg2+(g) + O2-(g) Mg+(g) + O-(g) Mg(g) + O(g) ΔHo6 7 ΔHo5 ΔHoR= Σ ΔHoi i=1 Mg+(g) + O-(g) ΔHo7 ΔHo3 ΔHo4 Mg(g) + O(g) ΔHo1 ΔHo2 ΔHoR MgO(s) Mg(s) + ½ O2(g)

THE OCTET RULE COVALENT BONDING sharing electrons (electron pair) F electronic configuration of F is 1s22s22p5 F F

F F + F F non-bonding, or lone pair of electrons bonding pair of electrons

H2 is the simplest covalent molecule + H H H H + the bond length of H2 is the distance where the total energy of the molecule is minimum

EXAMPLES H2O NH3 CH4 HX single bonds O2 CO2 C2H4 double bonds C2H2 N2 HCN triple bonds

few 100 kJ/mol

IONIC OR COVALENT electronegativity difference between two atoms involved in the bond

energies of the atomic orbital with the unpaired electron ELECTRONEGATIVITY is the tendency of an atom in a bond to attract shared electrons to itself F > O > N, Cl > Br > I, C, S …….. Na, Ba, Ra > K, Rb > Cs, Fr Electronegativity increases F Cl I O N C S Na K Rb Cs Fr Ra Ba Li Br energies of the atomic orbital with the unpaired electron

ELECTRONEGATIVITY F is the most electronegative F > O > N, Cl > Br > I, C, S …….. Na, Ba, Ra > K, Rb > Cs, Fr Electronegativity increases Li Br P I N O Cl F C S Na K Rb Cs Fr Ra Ba Se F is the most electronegative H has an electronegativity about the same a P

IONIC VERSUS COVALENT BONDS bonds are neither completely ionic nor covalent (only in homonuclear molecules)

IONIC VERSUS COVALENT BONDS compounds composed of elements with large difference in ELECTRONEGATIVITY significant ionic character in their bonding B has greater electronegativity A B

IONIC VERSUS COVALENT BONDS B has a greater share A B

Fluorine is more electronegative than hydrogen. HYDROGEN FLUORIDE Fluorine is more electronegative than hydrogen. + F F H H +

Fluorine is more electronegative than hydrogen. HYDROGEN FLUORIDE Fluorine is more electronegative than hydrogen. + F F H H + d+ d– This is a polar covalent bond (dipole moment). The bond has a partly ionic and partly covalent nature.

Microwave Spectroscopy molecules need a dipole moment

Variation of ionic character with electronegativity.

LEWIS SYMBOLS IONIC COMPOUDS COVALENT COMPOUNDS ELECTRONEGATIVITY

single – double – triple Lewis considers only valence electrons H2O O H bonding pair of electrons non-bonding, or lone pair of electrons single – double – triple

LEWIS STRUCTURES 1. concept of resonances 2. exceptions to the octet rule

1. RESONANCES O O NO3- O: 1s22s22p4 N: 1s22s22p3 plus one extra electron for negative charge

- + -

experiment shows all three bonds are the same 128 pm N bond angles 120 0 any one of the structures suggests one is different!

modify the description by blending the structures 128 pm bond angles 120 0 modify the description by blending the structures blending of structures is called resonance

use a double headed arrow between the structures RESONANCE use a double headed arrow between the structures N O O N N O electrons involved are said to be DELOCALIZED over the structure. blended structure is a RESONANCE HYBRID

We use a double headed arrow between the structures.. RESONANCE We use a double headed arrow between the structures.. N O O N N O O N

CO32- NO2-

2. Exceptions to the octet rule 1. more than 8 electrons around central atom 2. less than an octet around central atom 3. molecules with unpaired electrons

1. more than 8 electrons around central atom elements in rows 3 and following can exceed octet rule SF6 S F F S participation of d electrons

Lewis structure for SF6 1s22s22p5 F has seven S has six 1s22s22p63s22p4 SF2 SF4 SF6 PF3 PF5 NF3 NF5 ClO4- SO42- I3-

2. less than an octet around central atom BeH2 AlF3 resonances BF3 NH3 (dative bond) Lewis base Lewis acids

3. molecules with unpaired electrons FREE RADICALS NO but not NO-