CHM 122 – GENERAL INORGANIC CHEMISTRY 2UNITSMODULE I: CHEMICAL BONDING.

OBJECTIVES

a) To study the formation of the major types of chemical bonds usually encountered in compounds.b) To study the general characteristic properties of the compounds formed by the different chemical bonding.

The types of chemical bonding to be considered arei) Ionic or electrovalent bondingii) Covalent bondingiii) Coordinating covalent (dative) bondingiv) Metallic bondingv) Hydrogen bonding

c) In this module you will also learn about the following concept or terms:(i) Valence electrons (ii) Ionization energy (iii) Electron affinity (iv) Octet rule and exceptions to the Octet rule (v) Lewis formulas or structures(vi) Non – bonding electrons, bonding pairs of electrons and lone pairs of electrons(vii) Polar compounds (viii)Intermolecular forces of attraction etc

1

INTRODUCTION

i) Valence (or Outermost Shell) ElectronsValence electrons are those involved in forming compounds. They are the outermost shell electrons of atoms. For example, nitrogen has the electronic configuration 1s2 2s2 3p3. The outermost shell electrons are those in the 2s and 2p orbital’s.

Atom and Electronic Configuration

Valence Electrons

N[1s2 2s2 2p3] 2s2 2p3 = 5O[1s2 2s2 2p4] 2s2 2p4 = 6CI[1s2 2s2 2p63s23p5] 3s2 3p5 = 7Na[1s2 2s2 2p6 3s1] 3s1 = 1Mg[1s22s22p63s2] 3s2 = 2AL[1s2 2s2 2p6 3s2 3p1] 3s23p1 = 3

ii) Ionization Energy (IE) or Ionization Potential (IP) is the energy required to remove one mole of electrons from one mole of gaseous atoms or ions e.g.

Process Energy ev mol-1

KJ mol-1

a)Na[1s22s22p63s1](g) Na+[1s22s22p6] + e 5.10 495b) Al (g) Al+(g) + e 5.95 577

Al+ (g) Al2+(g) + e 18.82 1816Al2+(g) Al3+(g) + e 28.44 2745

This example illustrates that as the charge on the atom increases IE increases

General Comments on IE In general metals have relatively low IE compared to non-metals In a period IE increases as the nuclear charge, z increases. In a Group IE decreases as z increases e.g. Li Li+ + e, IE = 5200 KJ mol-1

Na Na+ + e, IE = 495 KJ mol-1

K K+ + e, IE = 418 KJ mol-1

Rb Rb+ + e, IE = 403 KJ mol-1

iii) Electron Affinity (EA) is the energy change accompanying the addition of one mole of electrons to one mole of gaseous atoms or ionse.g. F (g) + e F(g)-, EA = -364 KJmol-1

Cl (g) + e Cl-(g), EA = -342 KJmol-1

I(g) + e I(g) -, EA = -295 KJmol-1

O(g) + e O-(g), EA = -141 KJmol-1

O-(g) + e O2-(g), EA = +791 KJmol-1

2

S(g) + e S-(g), EA = -200 KJmol-1

S –(g) + e S2-(g), EA = + 649 KJmol-1

General Comments on EA In most cases, energy is liberated when the first electron is added because it is

attracted to the atom’s nuclear charge, hence EA, is usually negative. EA2 is always positive because energy must be absorbed to overcome

electrostatic repulsive ion. In general non – metals have relatively high EA values than metals As a general trend EA increases as z increases in a Period. In a Group EA decreases as z increases. This is an indication of increasing

metallic properties.

iv) Electronegativity (EN) is the relative ability of a bonded atom to attract the shared electrons. Alternatively, EN is a measure of the relative tendency of an atom to attract electrons to itself when chemically combined with another atom. Elements that tend to attract rather than lose electrons are said to electronegative. Elements with high EN values (nonmetals) often gain electrons to form anions while elements with low EN values (metals) often lose electrons to form cations e.g. O + e O-

F + e F-

Cl + e CI-

Na Na+ + e Ca Ca2+ + 2eEN of the elements are expressed on a somewhat arbitrary scale, called Pauling Scale.

Table 1: Electronegativity Values of the main group elements

3

Group Group Group Group Group Group Group 1 2 13 14 15 16 17(IA) (IIA) (IIIA) (IVA) (VA) (VA) (VIIA) H 2.2

Li1.0

Be1.6

B2.0

C2.6

N3.0

O3.4

F4.0

Na0.9

Mg1.3

AI(III)1.6

Si1.9

P2.2

S2.6

Cl3.2

K0.8

Ca1.0

(d-blockelements)

Ga(III)1.8

Ge(IV)2.0

As(III) 2.2

Se2.6

Br3.0

Rb0.8

Sr0.9

In (III)1.8

Sn(II)1.8Sn(IV)2.0

Sb2.1

Te2.1

I2.7

Cs0.8

Ba0.9

TI(I)1.6TI(III)2.0

Pb(II)1.9Pb(IV)2.3

Bi2.0

Po2.0

At

EN increases as z increases across a Period EN decreases as z increases down a Group EN values can be used to predict the bond type in a particular compound as

indicated in Table 2 below.

Table 2 . Relationship between bond type and electronegativity difference (∆EN) of the two bonded atoms.Bond ∆EN Types of BondCl – Cl 3.2 - 3.2 = 0 CovalentCl – C 3.2 – 2.6 = 0.6 Polar CovalentCl – H 3.2 – 2.2 = 1.0 Polar CovalentF – Na 4.0 -0.9 = 3.1 Ionic

4

Ionic Bonding:

Na (1S 2  2S2 2P6 3S1 ) Na+ (1S 2  2S2 2P6 ) + e

Cation

Cl[Ne]3S2 3P5 + e Cl-[Ne]3S2 3P6

AnionThe ions thus formed are held together by coulombic or electrostatic forces e.g.

Na+ + Cl- NaClThe bond formed is called electrovalent or ionic bond.

Further examples:

a) MgCl2

Mg [Ne] 3S2 Mg2+ [Ne] + 2e Cl[Ne] 3S2 3P5 + e Cl- [Ne] 3S2 3P6

Mg2+ + 2Cl- MgCl2

b) CaO Ca[Ar] 4S2 Ca2+ [Ar] + 2e O [He] 2S2 2P4 + 2e O2- [Ne] Ca2+ + O2- CaO

Factors influencing the formation of ions

i) Ease of formation of cations- i.e. lower IEii)Ease of formation of anions- i.e. the more negative the EA

is, the more stable is the anion formed. e.g. F F- ; Cl Cl- ; Br Br- ; I I-

-333 -364 -342 -295KJ mol-1

O O- ; O- O2- ; S S- ; S- S2-, -141 +791 -200 +600 KJ mol-1

5

iii) Lattice energy : The more exothermic the lattice energy, the more stable is the ionic compound formed. e.g. Na+ (g) + F- (g) NaF(s), ∆HLE = -915 KJ Na+ (g) + Cl- (g) NaCI(s), ∆HLE = -781 KJ Mg2+(g) + O2- (g) MgO(s), ∆HLE = -3933 KJ

By definition, the lattice energy is the energy change when one mole of a crystal is formed from its component ions in the gaseous state.

General Characteristic Properties of Ionic Compoundsi) They exist not in molecular but in ionic forms e.g. NaCl exists in the solid state as Na+CI-

ii) As a result of (i), ionic compounds conduct electricity when in molten state or in solution.

iii) They have relatively high melting and boiling points e.g. melting points (°C) of NaCl is 1413, CaF2 ~ 2500.

iv) They dissolve readily in solvents with high dielectric constants (polar solvents e.g. water) but generally insoluble in non – polar solvents such as benzene, hexane, carbon tetrachloride, chloroform etc.

Covalent Bonding:

Covalent bonding involves the sharing of one or more pairs of electrons between two atoms of either the same element or two different elements e.g.

6

a) Formation of hydrogen molecule, H2

Using a dot/cross representation:

The bond formed is called covalent bond or electron pair bond.

b) Formation of chlorine molecule, CI2

c) Formation of hydrogen chloride, HCI

The electron pairs between the two atoms are called bonding pairs of electrons. The remaining electron pairs on each atom are called lone pairs of electrons or non-bonding electrons.

Octet RuleIn covalent bonding, each atom in the bond often attains the electronic configuration of the nearest noble gas i.e. the atom part from H tends to achieve an octet of electrons in its valence shell.

Further examples of the formation of covalent compoundsi) Water,H 2O

H + xH x H H H H

Covalent bond

or

H2

Cl

+ Cl

x

x

x

xxx

xCl

Clx

x

x

xxx

xCl Cl

Cl2

Cl

xH+ Cl

x H H Cl

H Cl

7

ii) Ammonia, NH3

iii) Methane, CH 4

The structures with the covalent bonds indicated are called Lewis structures.

Remarks:

i) Homonuclear diatomic molecules such as H2, CI2, O2

etc have even distribution of bonding electron density – equal share of the bonding pairs of electrons.

ii) Heteronuclear diatomic molecules such as HCI, HF etc have an uneven distribution of bonding electron density because of the difference in the electronegativities of the elements.

xH2 + O

x H

xH O

H2O

O

H H

xH3 + N

x H

xH N

NH3

N

H Hx

H H

xH4 + C

x H

xH C

CH4

CH Hx

HH

x

H H

H H H Clδ+ δ-

d

8

Partial charges

The existence of partial charges imparts a dipole moment µ to the molecules i.e. thus leading to bond polarity µ = charges x d (distance between charges)

General Characteristic Properties of Covalent Compoundsi) Covalent compounds exist in molecular forms

(i.e discrete molecules or a big giant molecule such as diamond).

ii) The forces between the molecules are weak. As a result, they generally have low melting and boiling points except giant molecules.

iii) At room temperature, they are usually gases (CH4, CO2, SO3, e.t.c), liquids (H2O, C2H5OH, C6H6) or low melting solids (urea, glucose, phenol)

iv) They are usually insoluble in polar solvents but are soluble in non-polar solvents.

v) Covalent compounds do not conduct electricity in the solid or molten state or in solution. They are non-electrolytes.

Failure of Octet RuleThe atoms in some molecules cannot obey the octet rule because there are too few or too many electrons. This generally occurs when an atom forms more than 4 bonds e.g. PCl5 (5bonds), SF6(6bonds).

9

In a few cases, an atom has less than an octet e.g.

Writing of Lewis Structures

The following guidelines will be found useful

Examples1. write the Lewis structure for SO3

Solution Oi) Skeletal structure: O S

PCl

Cl

Cl

Cl

Cl

SF

F

F

F

F

F

8e-s around each Cl, 10e-s around P8e-s around each F, 12e-s around S

Cl

Cl

Be BCl

Cl

Cl

6e-s around B4e-s around Be

10

Decide which atoms are bonded Count all valence electrons

Place any remaining e on central atoms in pairs

If central atom has no octet, form double or triple bonds if necessary.

Oii) Valence electrons: 6 x 3 + 6 = 24 Oiii) Placing 2 electrons in each bond O S Oiv) Remaining electrons: 24 – 6 = 18e

v) Complete octet of the oxygen atom

O

O S OCentral atom has no octet. A double bond will enable it to have octet.

Lewis structure

O O S O S O

2) Write the Lewis structure for C O

Solution.

i) Skeletal structure: C Oii) Valence electrons: 4 + 6 = 10eiii) Form electron pair bond: C: O

11

O

O

iv) Remaining electrons: 10 – 2 = 8e.v) Form octet around oxygen atom

For C and O atom to have octet, triple bond has to be formed.

Study QuestionWrite the Lewis structure for SO4

2-

Formal Charge is defined as the residual charge on an atom in the Lewis structure and is obtained by using the following equation;

Formal charge = Group Number – Number of unshared electrons –Number of bonds (a) OR = Valence electrons – Number of unshared electrons - Number of bonding electrons (b)

Example 1. The Lewis structure of HCN is

Determine the formal charge of each structure and hence the more plausible structure.

SolutionUsing the equation: Formal charge = Group No – No of unshared electrons – No of bonds.

C O

C O C O

C N

H or N CHI II

12

Structure I H: 1 – 0 – 1 = 0 C: 4 – 0 – 4 = 0 Formal charges on the atoms N: 5 - 2 – 3 = 0

Structure II H: 1 – 0 – 1 = 0 C: 4 – 2 – 3 = -1 Formal charges on the atoms N: 5 – 0 – 4 = +1 Based on the formal charges on the atoms, structure I in which the formal charges are zero is more plausible structure.

Example 2.The Lewis structure of thionyl chloride could be written as;

Determine the formal charge on each atom and hence the more probable structure.

SolutionUsing equation (a), the calculated formal charges are:

Structure I CI: 7 – 6 – 1 = 0 S: 6 – 2 – 3 = 1 Formal charges on the atoms O: 6 - 6 - 1 = -1

Structure II CI: 7 – 6 – 1 = 0 S: 6 – 6 – 1 = -1 Formal charges on the atoms O: 6 - 2 - 3 = 1

Cl C O

Cl

or

Cl O S

Cl

I II

13

Oxygen is a more electronegative atom than S. Therefore, structure I is more probable than structure II.

Study Questionsi) Write the Lewis structure of each of the following species:a) HCOF b) CIF3

2) Determine all the possible structures for a) HCOF.

Using the formal charges on the atoms technique, determine the most probable structure. RESONANCE:There are some cases in which the Lewis structure of a specie does not account for the properties (e.g. bond length and bond energy) of the ion or molecule it represents. For example. Consider the formate ion, .

Its Lewis structure is:

The actual structure of the molecule or ion that we cannot draw satisfactorily is considered as an average of a number of Lewis structures that we can draw. For , we can write

I IIThe actual structure of the ion is said to be a resonance hybrid of the two contributing structures I and II.

H C

O

O

CH

O

O CH

O

O

14

The concept of resonance does not suggest that the structure of fluctuates between I and II or the ion exists in two different forms. There is only one form of the ion that can be observed experimentally.

A resonance hybrid may be the characteristics of its “parents” but it never has the exact structure of any one of them.

Further Example:Draw the resonance structures of: i) SO3,

ii)

Solution:i) SO3: The resonance structures are

iv) NO

Lewis Structure:

SOO

O

SOO

O

SOO

O2 2 2

15

Resonance Structures:

Coordinate Covalent (Dative) Bonding:This is a special case of covalent bonding. Here, the electrons shared are all donated by only one of the atoms involved in the bond formation. Usually the donor atom is an atom having an unshared pair of electrons in the valence shell after the formation of the covalent bonds e.g. O, N, CI.

An acceptor atom is an atom which is electron deficient i.e. lacks sufficient valence electrons to attain the octet configuration e.g. B, AI, H+, Mn+ ( where M = a metal and n+, the oxidation number e.g. Mg2+,Cu2+, fe2+)

Formation of the BondExample 1

(H3O+, Hydronium ion)

NO

O

O

NOO

O

NOO

O

NOO

O

H+ + H2O O HH

H

+

16

Denotes the dative bond.

The oxygen atom donates one of its lone pairs of electrons which is shared between it and H+ ion.

Example 2

Denotes the dative bond.

Example 3

Example 4

HF

BF3 + NH3

F B N

F

H

H

H

Cl

AlCl3 + NH3

Cl Al N

Cl

H

H

Cu2+ + 4H2O [Cu(H2O)4]2+

OHH

O

HH

Cu2+O

HH

O

H

H

17

Metallic Bonding is pictured as a series of (+) ions surrounded by a sea of mobile valence electrons (-) i.e

In solid metals, the atoms are packed closely together in regular array: Closed – packed or body – centered cubic systems. The structures are held together by metallic bonding.

Mg2+ + 6H2O

HH

OMg2+

H

O

O

O

O

O

HH

H

H

H

HH

HH

[Mg(H2O)6]2+

Co2+ + 6NH3 Co2+ [Co(NH3)6 ]2+

H3N

NH3

NH3

NH3NH3

H3N

+ + ++ ++ +

+ + ++ ++ +Metal ions

+ + ++ ++ +

+ + ++ ++ +

Valence electrons

18

Properties of metallic compounds1. Metals are solids except Hg2. They are hard (Pt, Fe) or soft (Li, Na, K)3. They have luster appearance4. They are good conductors of heat and electricity.

Hydrogen Bonding:1. IntroductionCovalent bonding between two different electronegative atoms always leads to bond polarity. For instance, a molecule that has a highly electronegative atom connected to a hydrogen atom is strongly polar e.g. HF, NH3, H2O. HF : H F

H2O: O H H NH3 : N

H H H

19

The hydrogen atom, because of its partial positive charge, is attracted to a lone pair of electrons on the atom of another molecule. The new bond formed is called a Hydrogen Bond i.e

i) HF + HF H – F H – F

ii) H2O + H2O H O H O

H H

H H

iii) NH3 + NH3 H N H N

H H

Covalent bond Hydrogen bond

Hydrogen bond is essentially an electrostatic bond, but certainly not as strong as ionic bond.The bond strength (or energy) is weaker than that of covalent bond, but stronger than that of van der waals bonds as indicated in the table below:

Nature of Bond Typical Bond strength (KJ/ mol)

Ionic bondsCovalent bondsHydrogen bondsVan der waal bonds

>1000100 – 90020 – 50<20

20

The higher the electronegativity of the heteroatom X, the stronger the hydrogen bond. Thus N – H N < O-H….O < F-H….F (weakest) (strongest)

This parallels the increasing electronegativity of X (N, O, F) atom bonded to hydrogen; N < O < F

21