C H E M I S T R Y
Chapter 6
Ionic Bonds and Some Main-Group Chemistry
Copyright © 2008 Pearson Prentice Hall, Inc.
12.1 Types of Chemical Bonds
 Bonds: a force that holds groups of two or more atoms
together and makes them function as a unit
 Required 2 e- to make a bond
 Bond energy: amount of energy required to form or to
break the bond
Ionic Bonding
 Occurs in ionic
compound
 Results from transferring
electron
 Created a strong
attraction among the
closely pack compound
Ions and Their Electron
Configurations
Atoms
Ions
N: 1s2 2s2 2p3
+ 3 e-
N3-: 1s2 2s2 2p6
O: 1s2 2s2 2p4
+ 2 e-
O2-: 1s2 2s2 2p6
F: 1s2 2s2 2p5
+ 1 e-
F1-: 1s2 2s2 2p6
Ne: 1s2 2s2 2p6
Na: 1s2 2s2 2p6 3s1
- 1 e-
Na1+: 1s2 2s2 2p6
Mg: 1s2 2s2 2p6 3s2
- 2 e-
Mg2+: 1s2 2s2 2p6
- 3 e-
Al3+: 1s2 2s2 2p6
Al: 1s2 2s2 2p6 3s2 3p1
Ions and Their Electron
Configurations
Ions and Their Electron
Configurations
Atoms
Ions
Fe: [Ar] 4s2 3d6
- 2 e-
Fe2+: [Ar] 3d6
Fe: [Ar] 4s2 3d6
- 3 e-
Fe3+: [Ar] 3d5
Ionic Radii
Ionic Radii
Ionization Energy
Ionization Energy (Ei): The amount of energy
necessary to remove the highest-energy electron
from an isolated neutral atom in the gaseous state.
Ionization Energy
Ionization Energy
Boron has a lower Ei due to a smaller Zeff
(shielding by the 2s electrons)
Ionization Energy
Oxygen has a lower Ei since the first
electron is removed from a filled orbital
Higher Ionization Energies
M + energy
M1+ + e-
M1+ + energy
M2+ + e-
M2+ + energy
M3+ + e-
Electron Affinity
Electron Affinity (Eea): The energy released when
a neutral atom gains an electron to form an anion.
Ionic Bonds and the Formation
of Ionic Solids
1s2 2s2 2p6 3s1
1s2 2s2 2p6 3s2 3p5
Na + Cl
Na1+
1s2 2s2 2p6
Cl11s2 2s2 2p6 3s2 3p6
The Octet Rule
Octet Rule: Main-group elements tend to undergo
reactions that leave them with eight outer-shell
electrons.
The Octet Rule
Octet Rule: Main-group elements tend to undergo
reactions that leave them with eight outer-shell
electrons.
Metals tend to have low Ei and low Eea.
They tend to lose one or more electrons.
Nonmetals tend to have high Ei and high Eea.
They tend to gain one or more electrons.
The Octet Rule
Chapter 7
Covalent Bond and Molecular Structure
The Covalent Bond
Covalent Bond: A bond that results from the
sharing of electrons between atoms.
Polar Covalent Bonds:
Electronegativity
Electronegativity: The ability of an atom in a molecule
to attract the shared electrons in a covalent bond.
Polar Covalent Bonds:
Electronegativity
NaCl
Cl2
HCl
Polar Covalent Bonds:
Electronegativity
Relationship Between Electronegativity
and Bond Type
Predicting bond polarity




Atoms with similar electronegativity (Δ EN <0.4) –form
nonpolar bond
Atoms whose electronegativity differ by more than two (Δ
EN > 2) – form ionic bonds
Atoms whose electronegativity differ by less than two (Δ
EN < 2) – form polar covalent bonds
Polarity
 Polar covalent bonds – the bonding electrons are attracted
somewhat more strongly by one atom in a bond
 Electrons are not completely transferred
 More electronegative atom: δ- . (δ represents the partial negative charge
formed)
 Less electronegative atom: δ+
Relationship Between Electronegativity
and Bond Type
Predicting bond polarity




Atoms with similar electronegativity (Δ EN <0.4) –form
nonpolar bond
Atoms whose electronegativity differ by more than two (Δ
EN > 2) – form ionic bonds
Atoms whose electronegativity differ by less than two (Δ
EN < 2) – form polar covalent bonds
Examples
 For each of the following pairs of bonds, choose the bond
that will be more polar
a. H-P, H-C
b.
N-O, S-O
Polarity and Dipole Moment
 Dipole moment:
 a vector quantity from the
center of the positive
charge to the center of
negative charge
 Represents with an arrow
E.g Draw the dipole
moment for HF, H2O,
HCl, OF
Lewis Structures
 represents how an atom’s valence electrons are distributed in
a molecule
 Show the bonding involves (the maximum bonds can be
made)
 Try to achieve the noble gas configuration
Examples
 Draw Dot Lewis structure for the following atoms:
 Na
 Mg
 C
 S
 Co
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
6/32
Rules
 Duet Rule: sharing of 2 electrons
 E.g H2
 H:H
 Octet Rule: sharing of 8 electrons
 Carbon, oxygen, nitrogen and fluorine always obey this rule in
a stable molecule
 E.g F2, O2
 Bonding pair: two of which are shared with other
atoms
 Lone pair or nonbonding pair: those that are not
used for bonding
Electron-Dot Structures
Electron-Dot Structures
Rules for Wring Dot Lewis structure
 Draw a dot Lewis structure of ClO4-
Step 1: Calculate the total number of valence
electrons of all atoms in the molecule
 Cl – Valence e- = 7
 O – Valence e - = 6 x 4 = 24e ClO4- => total valence e- = 7 + 24 +1 ( -1 charge) = 32
e-
Rules
Ste 2: Create a skeletal structure using the following
rules:
 Hydrogen atoms (if present) are always on the “outside” of the structure.
They form only one bond
 The central atom is usually least electronegative. It is also often unique
(i.e,. the only one atom of the element in the molecule). Remember, there
might be no “central” atom.
 Connect bonded atoms by line (2-electron, covalent bonds
O
O
Cl
O
O
O
O
Cl
O
O
Rules
Step 3: Place lone pairs around outer atoms (except
hydrogen) so that each atom has an octet
O
O
Cl
O
O
Rules
Step 4: Calculate the number of electrons you haven’t
used. Subtract the number of electrons used so far,
including electrons in lone pair and bonding pairs, from the
total in Step 1. Assign any remaining electrons to the central
atom as lone pair
O
 Cl-O bonds = 4 x 2e- = 8 e O – 4 x 6e- = 24 eCl
O
O
 Total used = 8 + 24 = 32 eO
Rules
Step 5: If the central atom is B (boron) or Be
(beryllium), skip this step

If the central atom has an octet after step 4, skip this step
 If the central atom has only 6 electrons, move a lone pair
from an outer atom to form a double bond between outer
atom and the central atom
 If the central atom has only 4 electrons, do Step 5a to two
different outer atoms (i.e, form two double bonds) or twice to
one outer atom (i.e., form one triple bond)
Electron-Dot Structures of
Polyatomic Molecules
Draw an electron-dot structure for H2O.
Step 1: Total valence electrons
Step 2:
Draw its skeletal
Step 4: Dot Lewis structure
Examples
 Give the Lewis structure for the following
 CF4,
NH4+
 BF3
CO2
 NO3-,
CH2O