Bonding between atoms
Bonds
 Forces that hold groups of atoms
together and make them function
as a unit.
 Ionic bonds – transfer of
electrons
 Covalent bonds – sharing of
electrons
Molecule: collection of atoms
that are bound together
• What holds a molecule together? Covalent
bond
• Force of attraction that results from valence
electrons being attracted to two nuclei.
• Valence electrons being shared between two
nuclei
• Core electrons are not involved in bonding
Electronegativity
The ability of an
atom in a molecule to
attract shared
electrons to itself.
Linus Pauling
1901 - 1994
Table of Electronegativities
Covalent Bonds
Polar-Covalent bonds
 Electrons are unequally shared
 Electronegativity difference between .3 and 1.7
Nonpolar-Covalent bonds
 Electrons are equally shared
 Electronegativity difference of 0 to 0.3
Energy is released when forming a
covalent bond and absorbed when
breaking it
Octet Rule and electron sharing
Lewis Dot diagrams
• Show an atom’s valence electrons
• The number of dots is equal to the group
number
• Pairs of dots(electrons) are electron pairs
• Single electrons are unpaired electrons
Examples of Lewis Dot diagrams
• All elements in group IA will have one dot
• All elements in group IIA will have two
dots
• All elements in group IIIA will have three
dots
• Group IV will have four dots, V five, VI
six, VII seven, and VIII eight dots
I
II
III
Transition metals
Metal
Metalloids
Nonmetals
IV
V
VI VII
VIII
Completing a Lewis Structure -CH3Cl
Make carbon the central atom
Add up available valence electrons:
C = 4, H = (3)(1), Cl = 7
Join peripheral atoms
to the central atom
with electron pairs.
H
..
HH
C
..
Complete octets on
atoms other than
hydrogen with remaining
electrons
Total = 14
..
H
Cl
Lewis Structures
1) Count up total number of valence electrons
2) Connect all atoms with single bonds
- “multiple” atoms usually on outside
- “single” atoms usually in center;
C always in center,
H always on outside.
3) Complete octets on exterior atoms (not H, though)
4) Check
- valence electrons math with Step 1
- all atoms (except H) have an octet; if not, try multiple bonds
- any extra electrons? Put on central atom
Comments About the Octet Rule
2nd row elements C, N, O, F observe the octet
rule.
2nd row elements B and Be often have fewer
than 8 electrons around themselves - they are
very reactive.
3rd row and heavier elements CAN exceed the
octet rule using empty valence d orbitals.
When writing Lewis structures, satisfy octets
first, then place electrons around elements having
available d orbitals.
Exceptions to the Octet Rule
• Atoms with less than an octet: BF3
• Atoms with more than an octet: most
notably phosphorus and sulfur (SF4)
additional electrons fill the 3d orbitals
• Molecules with odd number of electrons:
NO (nitrogen monoxide)
Multiple bonds
• Determine the total number of valence
electrons that will be in the final diagram
• Connect atoms with single bonds
• Put in remaining dots two at a time as lone
pairs
• Send lone pairs to atoms that do not have an
octet(lone pairs must come from adjacent
atoms)
Multiple Covalent Bonds:
Double bonds
Two pairs of shared electrons
Multiple Covalent Bonds:
Triple bonds
Three pairs of shared electrons
Ionic Bonding
 Forms between metals and nonmetals
 There is a strong attraction between
oppositely charged ions
 Form ionic lattices (network of positively
and negatively charged ions)
 Examples: sodium oxide, magnesium
chloride, calcium phosphate
Properties of Ionic Compounds
Structure:
Crystalline solids
Melting point: Generally high
Boiling Point:
Generally high
Electrical
Conductivity:
Solubility in
water:
Excellent conductors,
molten and aqueous
Generally soluble
Ionic Bonds
 Electrons are transferred
 Electronegativity differences are
generally greater than 1.7
 The formation of ionic bonds is
always exothermic!
Ionic Bonding:
The Formation of Sodium Chloride
 Sodium has 1 valence electron
 Chlorine has 7 valence electrons
 An electron transferred gives
each an octet
Na 1s22s22p63s1
Cl 1s22s22p63s23p5
Ionic Bonding:
The Formation of Sodium Chloride
This transfer forms ions, each
with an octet:
Na+ 1s22s22p6
Cl- 1s22s22p63s23p6
Ionic Bonding:
The Formation of Sodium Chloride
The resulting ions come together
due to electrostatic attraction
(opposites attract):
+
Na
Cl
The net charge on the compound
must equal zero
Sodium Chloride Crystal Lattice
Ionic compounds form solids
at ordinary temperatures.
Ionic compounds organize in
a characteristic crystal
lattice of alternating positive
and negative ions.
Monatomic
Cations
H+
Li+
Na+
K+
Mg2+
Ca2+
Ba2+
Al3+
Name
Hydrogen
Lithium
Sodium
Potassium
Magnesium
Calcium
Barium
Aluminum
Writing Ionic Compound Formulas
Example: Barium nitrate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges are
balanced.
2+
Ba ( NO3- ) 2
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
Not balanced!
Writing Ionic Compound Formulas
Example: Ammonium sulfate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
( NH4+) SO42-
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
2
Not balanced!
Writing Ionic Compound Formulas
Example: Iron(III) chloride
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
Fe3+ Cl-
3
Not balanced!
Writing Ionic Compound Formulas
Example: Aluminum sulfide
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
3+
Al
2
2S
3
Not balanced!
Writing Ionic Compound Formulas
Example: Magnesium carbonate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
2+
Mg
CO3
2-
They are balanced!
Writing Ionic Compound Formulas
Example: Zinc hydroxide
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges are
balanced.
2+
Zn
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
( OH- )2
Not balanced!
Writing Ionic Compound Formulas
Example: Aluminum phosphate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges are
balanced.
3+
Al
PO4
3-
They ARE balanced!
Naming Ionic Compounds
•1. Cation first, then anion
•2. Monatomic cation = name of the element
•Ca2+ = calcium ion
•3. Monatomic anion = root + -ide
•Cl- = chloride
•CaCl2 = calcium chloride
Naming Ionic Compounds
•1. Cation first, then anion
•2. Monatomic cation = name of the element
•Ca2+ = calcium ion
•3. Monatomic anion = root + -ide
•Cl- = chloride
•CaCl2 = calcium chloride
Naming binary covalent
compounds
• Name less electronegative element first
• The more electronegative element takes a
suffix-ide
• For each element, use a prefix to indicate
number of atoms present





Naming Binary Covalent
Compounds
Compounds between two nonmetals
First element in the formula is named first.
Second element is named as if it were an anion.
Use prefixes
Only use mono on second element P2O5 = diphosphorus pentoxide
CO2 = carbon dioxide
CO = carbon monoxide
N2O = dinitrogen monoxide
Covalent compounds
• Phosphorous trichloride
• Carbon dioxide
• Sulfur hexafluoride
Naming chemical compounds
• Compounds are named according to atoms
or ions that compose them
• Binary molecular are named after
component atoms, with prefixes
• An acid is named after its characteristic
anion
Bond type by electronegativity
• Electronegativity difference less or equal to
0.4  bond type nonpolar covalent
• Between 0.4 and 2.0 polar covalent
• Greater than 2.0 ionic