BASIC CONCEPTS OF CHEMICAL BONDING
An ionic bond is a chemical bond formed by the electrostatic attraction between cations and anions.
The bonding in sodium chloride, NaCl, is ionic bonding.

2 Na(s) +  Cl2(g) ---> 2  Na1+Cl1-(s)

The sodium cation, Na1+, is formed when a sodium atom loses an electron.

    Na   -------->  Na1+  +  e1-
 [Ne] 3s1      [Ne]

The chloride anion, Cl1-, is formed when a chlorine atom gains an electron.

       Cl     +  e1-    --------->      Cl1-
[Ne] 3s23p5                 [Ne] 3s23p6 = [Ar]

Note that when the ions are formed in the previous equations, both ions adopt a noble gas electron configuration.

A covalent bond is a chemical bond formed by the sharing of a pair of electrons between atoms.   The bonding in hydrogen
chloride, HCl, is covalent bonding.   More detail will be given later in the notes on covalent bonding.



Lewis Symbols
A Lewis symbol is a representation of an element in which the chemical symbol stands for the core electrons
and the valence electrons are represented by dots placed around the letter symbol of the element.  The number of valence
electrons is the same as the group number for groups 1 and 2.   For groups 13 - 18, subtract 10 from the group number to
determine the number of valence electrons.   The Lewis symbol for Li through Ne are as follows.
 
Li .
. Be .
. B .
.
.
. C .
.
.
.  N :
..
.
.  O :
 ..
.
: F :
..
. . 
: Ne :
..



Lattice energy (lattice enthalpy) is the change in energy that occurs when an ionic solid is separated into isolated ions in the gas phase.
In general, strength of ionic bonds and lattice energy are directly proportional.   A large value in the lattice energy indicates a
strong ionic bond and vice versa.  An ionic substance with a large lattice energy will generally have a higher melting point.
LiCl(s) ---> Li1+(g) + Cl1-(g)          DHo = +834 kJ
NaCl(s) ---> Na1+(g) + Cl1-(g)        DHo = +786 kJ            [Melting point of NaCl 801oC]
KCl(s) ---> K1+(g) + Cl1-(g)           DHo = +701 kJ            [Melting point of KCl 770oC]


LEWIS STRUCTURES
A covalent bond is a chemical bond formed by the sharing of a pair of electrons between atoms.

H .  +  H .   ----->  H:H            {":" represents a covalent bond, that is, a shared pair of electrons}

A Lewis structure is a convention used to represent the covalent bonding in a molecule.   Dots or lines are used to represent
valence electrons (electrons available for bonding).  The Lewis structure for the hydrogen molecule is as follows.
     H:H  or  H-H

A bonding pair is an electron pair shared between two atoms.

A nonbonding pair (lone pair, unshared pair) is a pair of electrons that remains on one atom and is not shared.
 
In the Lewis structure for the chlorine molecule, Cl2, there is a single bonding pair located between the Cl atoms and a total of six (6) lone pairs distributed equally around the Cl atoms.   Each Cl atom is surrounded by three (3) lone pairs and shares one lone pair.  This makes a total of eight (8) valence electrons surrounding each Cl atom.  (See octet rule.)              . .     . .
          : Cl - Cl :           
             . .     . .

A coordinate covalent bond is a bond formed when both electrons in a bond are donated by one atom.
A + :B ---> A:B

A coordinate covalent bond is not essentially different from other covalent bonds; it involves
the sharing of a pair of electrons between two atoms.

The octet rule is the tendency of atoms in molecules to have eight electrons in their valence shell.



Multiple Bonds Bonds formed when two atoms share more than one pair of electrons are referred to
as multiple bonds.  Among the atoms which commonly exhibit multiple bonds are  C, N, O, and S.
A single bond consists of one pair of shared electrons between two atoms.
{e.g. H2}
   H-H
A double bond consists of two pairs of shared electrons between two atoms.
{e.g. O2}
   . .  . .
   O=O 
   . .  . .
A triple bond consists of three pairs of shared electrons between two atoms.
{e.g. N2}
 : N=N :


ELECTRONEGATIVITY
Electronegativity is a measure of the ability of an atom in a molecule to draw bonding electrons to itself.
In 1934 Robert S. Mulliken suggest on theoretical grounds that the electronegativity (X) of an atom be
given as half its ionization energy (I.E.) minus electron affinity (E.A.).

X = I.E. - E.A./2

Linus Pauling's scale is based upon bond energies. Fluorine is assigned a value of 4.0 and Li a value of 1.0.
The general trend is that electronegativy increases left to right across a period and bottom to top in a group.

The absolute value of the difference in electronegativity of two bonded atoms gives a rough measure
of the polarity to be expected in a bond.  When this difference is small, the bond is nonpolar.  When it
is large, the bond is polar (or, if the difference is extra large, perhaps ionic).

POLAR COVALENT BONDS
A polar covalent bond is a covalent bond in which the bonding electrons spend more time closer to one atom
than the other.   The majority of covalent bonds are polar.

The bond between hydrogen and chlorine in HCl is a polar bond.  The chlorine is more electronegative
and therefore attracts the bonding electrons to itself (i.e. electrons are not equally shared).   The chlorine atom
in HCl adopts a partial negative charge (d-) and the hydrogen atom adopts a partial positive charge (d+).

Polar covalent bonds can be represented by the following convention.
d+  d-
H--Cl



GUIDELINES FOR DRAWING LEWIS STRUCTURES
1.    Calculate the total number of valence electrons (bonding electrons).
        a.    The number of valence electrons for the elements in groups 1 and 2 is the same as the group number.
               For groups 13 - 18, the number of valence electrons is determined by subtracting 10 from the group number.
        b.   Add 1 for each unit of negative charge.
        c.    Subtract 1 for each unit of positive charge.

2.    Determine the central atom and draw a single bond to each of the atoms surrounding it.
        a.   The central atom will most often be the atom with the lowest electronegativity.
        b.   Molecules and polyatomic ions are symmetrical and compact.
        c.    In oxyacids, hydrogen is bonded to oxygen.  {H2SO4 is an oxyacid}
        d.   When the noble gases Kr and Xe are present in a molecule, they are central atoms.

3.    Complete the octet around each bonded atom.
           Octet Rule Exceptions:
                a.   Hydrogen requires only 2 electrons.  Hydrogen is always a terminal atom and never a central atom.
                b.   Beryllium requires only 4 electrons.
                c.   Boron requires only 6 electrons.
                d.   Elements in Period 3 and beyond can exceed the octet rule. In other words non-metallic atoms
                       with atomic numbers larger than 14 can hold more than 8 electrons.

4.    Distribute any remaining electrons as lone pairs around the central atom (or atoms).

5.    If there are fewer than eight electrons on the central atom move one or more electron pairs from a bonded atom(s)
       to to form a multiple bond to the central atom(s).     Atoms that can form multiple bonds are C, N, O, and S.

NOTE: Lewis structures for ionic species should be enclosed in brackets [ ] and the charge indicated.



FORMAL CHARGE AND LEWIS STRUCTURES
Formal charge is the difference between the valence electrons in an isolated atom and the number of electrons assigned
to that atom in a Lewis structure.

Rules for determining formal charge
1.  Half of the electrons of a bond are assigned to each atom in the bond (counting each dash as two electrons).
2.  Both electrons of a lone pair are assigned to the atom to which the lone pair belongs
3.  Formal charge = Valence electrons - assigned electrons.
4.  Whenever you can write several Lewis formulas for a molecule, choose the one having the lowest magnitudes
     of formal charge.
5.  When two proposed Lewis formulas for a molecule have the same magnitudes of formal charges, choose the one
     having the negative formal charge on the more electronegative atom.



RESONANCE & DELOCALIZED BONDING
For many molecules and ions (e.g. ozone, O3), when the rules for Lewis structures are followed two different valid Lewis structures are possible.  These structures are referred to as resonance structures.   The two valid Lewis structures for ozone are as follows.
 
 . .    . .   . .
: O - O=O 
    . .         . .
<=====>
. .   . .    . .
 O=O - O :
   . .         . .

One misconception which perhaps must be addressed is that electrons in bonds are not sticks or dots which are "stuck" in position but  can be thought of as being "spread out" over several atoms.  (Recall that quantum mechanics describes electrons in terms of three dimensional waves.)   In particular, the data from the measurement of the bonds in ozone, O3, indicates that instead of one single bond and one double bond (as suggested by Lewis structure rules) the bonds are an intermediate between the two.  The electrons in these bonds are said to be delocalized.



EXCEPTIONS TO THE OCTET RULE  See Step 3 in GUIDELINES FOR DRAWING LEWIS STRUCTURES.


BOND LENGTH AND BOND ORDER
Bond length (bond distance) is the distance between the nuclei in a bond.
Covalent radii values assigned to atoms in such a way that the sum of covalent radii of atoms A and B predicts
an approximate A-B bond length.
Bond order is the number of pairs of electrons in a bond.

BOND AND ENERGY
Bond energy (BE) the average enthalpy change for the breaking of an A-B bond in a molecule in the gas phase.
Because it requires energy to break bonds (endothermic process), bond energies are always positive numbers.

When a bond is formed, the energy is equal to the negative of the bond energy (energy is released, exothermic).

Bond formation  is exothermic.            Bond breaking is endothermic.

In general, the enthalpy of reaction (DH) is (approximately) equal to the sum of the bond energies for bonds
broken minus the sum of lthe bond energies formed.

DH = SBE(reactants) - SBE(products)

EXAMPLE
Using bond energies, calculate the enthalpy of reaction for the formation of hydrogen chloride gas from
hydrogen gas and chlorine gas.

    H2(g) + Cl2(g) ----> 2 HCl(g)    DH = ?

Bond Energies:        H-H     436 kJ/mol;            Cl-Cl    243 kJ/mol;            H-Cl    432 kJ/mol

DH = SBE(reactants) - SBE(products)

DH =[ (1 mol H-H bonds)(436 kJ/mol) + (1 mol Cl-Cl bonds)(243 kJ/mol) ] - [(2 mol H-Cl bonds)(432 kJ/mol)]

    = -185 kJ



BONDING THEORY AND MOLECULAR STRUCTURE

Valence-Shell Electron-Pair Repulsion Model:  VSEPR

VSEPR is a simple way of predicting the shapes of species that have main group elements as central atoms.
According to this theory, electron pairs in the valance shell repel each other and therefore will be as far
apart from each other as possible in three dimensional space.

The shape of a molecule is determined by the positions of the atoms that are bonded to the central atom.
The positions of the bonded atoms depend upon the number of bonded atoms in the molecule and upon
the steric number (SN).

The steric number (also referred to as total coordination number) is defined as the total number of atoms and
sets of unshared electron pairs around a central atom.  To determine the steric number draw the
Lewis structure and count the number of bonded atoms and lone pairs on the central atom.   Each double
bond and each triple bond count as 1 when determining steric number.

SN = number of bonded atoms + number of lone pairs on the central atom

Example 1
Calculate the steric number for H2O, SiO2, and ClF3.

For H2O the steric number is 4.  There are 2 bonded atoms and 2 lone pairs on the central atom.

For SiO2 the steric number is 2.  There are 2 bonded atoms and no lone pairs on the central atom.
Notice that a double or triple bond count one each in determining the steric number.
VSEPR theory treats multiple bonds as if they were a single pair of electrons between the bonded atom and the central atom.

For ClF3 the steric number is 5.  There are 3 bonded atoms and 2 lone pairs on the central atom.

VSEPR Formula (Notation)
VSEPR theory predicts shapes of molecules based upon the number of bonded atoms and lone pairs on the central atom.
Molecules with a certain number of bonded atoms and lone pairs are predicted to have a particular geometry.

One convention used to classify molecular shape predicted by VSEPR assigns a specific AXmEn formula to each shape.
A represents the central atom,
X represents the bonded atom(s)
E  represents any lone pairs on the central atom.
The subscripts m and n are the number of bonded atoms and the number of lone pairs respectively.

Different molecules with identical VSEPR formulas will have identical molecular shapes.

Example 2
Use VSEPR notation to describe H2O, SiO2, and ClF3.
H2O        AX2E2

SiO2       AX2

ClF3       AX3E2

Example 3
Predict the shapes of H2O, SiO2, and ClF3.

For H2O the predicted shape is angular.

For SiO2 the predicted shape is linear.

For ClF3 the predicted shape is T-shaped.

Molecular Shapes Predicted by VSEPR Theory
Steric
Number
Number of Bonding Pairs Number of Lone Pairs Molecular Geometry VSEPR Notation Bond Angles Example
2 2 0 Linear AX2  180o BeCl2
3 3 0 Trigonal Planar AX3 120o BH3
2 1 Angular AX2E1 SO2
4 4 0 Tetrahedral AX4 109.5o CH4
3 1 Trigonal Pyramidal AX3E1 NH3
2 2 Angular AX2E2 H2O
1 3 Linear AX1E3 HF
5 5 0 Trigonal Bipyramidal AX5 90o,120o,180o AsCl5
4 1 Seesaw AX4E1 SF4
3 2 T-shaped AX3E2 ICl3
2 3 Linear AX2E3 ICl21-
6 6 0 Octahedral AX6 90o, 180o SF6
5 1 Square Pyramidal AX5E1 BrF5
4 2 Square Planar AX4E2 IF41-




 

Sample molecules and ions for Lewis structures and VSEPR theory.
BINARY MOLECULES
 1 AsCl3
 2 AsF5
 3 BCl3
 4 BeCl2
 5 BH3
 6 Br2
 7 BrF5
 8 CH4
 9 CCl4
10 CO2
11 CO
12 Cl2
13 ClF3
14 H2O
15 H2S
16 ICl3
17 KrF2
18 KrF4
19 N2
20 N2O
21 NCl3
22 NH3
23 O2
24 O3
25 OF2
26 PCl3
27 PF5
28 SCl2
29 SeF4
30 SeF6
31 SF4
32 SF6
33 SiCl4
34 SO2
35 SO3
36 TeF4
37 XeF2
38 XeF4
39 XeF6
40 XeO3
41 XeO4
TERNARY MOLECULES
 1 OSF4
 2 COCl2
 3 POCl3
 4 SOCl2
 5 OCN1-
 6 SCN1-
 7 XeO2F2
 8 XeOF2
 9 XeOF4
10 (CH3)2CO
IONS
 1 AlCl41-
 2 BF41-
 3 BrO41-
 4 ClF41+
 5 ClO1-
 6 ClO21-
 7 ClO31-
 8 ClO41-
 9 CN1-
10 CO32-
11 H3O1+
12 I31-
13 ICl21-
14 ICl3
 
 
 
 

 

15 ICl41-
16 IF41-
17 IF41+
18 N31-
19 NO21-
20 NO21+
21 OH1-
22 PO43-
23 SF51-
24 SiO44-
25 SnCl51-
26 SO32-
27 SO42-
28 XeO64-
ACIDS
Binary
Acids
1 HF
2 HCl
3 HBr
4 HI

Ternary
Acids
(Oxoacids)
 1 H2SO3
 2 H2SO4
 3 H2S2O7
 4 H3PO4
 5 HNO2
 6 HNO3
 7 HClO
 8 HClO2
 9 HClO3
10 HClO4
11 H2C2O4
12 HC2H3O2

  Misc.
  1 C2H6
  2 C2H4
  3 C2H2
  4 C2N2
  5 C3O2
  6 C2O42-
  7 HCN
  8 H2O2
  9 N2F4
 10 N2H4
 11 N2O4
 12 S2O32-
 13 S2O72-
 14 GeF2

Odd Electron
  1 NO
  2 NO2
  3 ClO2