Unit 4 – Bonding, Nomenclature & Molecular
Structure
I. Introduction to
Bonding
I
II
III
IV
Vocabulary
Chemical Bond
attractive force between atoms or ions
that binds them together as a unit
bonds form in order to…
decrease potential energy (PE)
increase stability
Vocabulary
Molecule
Smallest electrically neutral unit of a
substance that still has the properties of
the substance
Made up of 2 or more atoms
Ion
Atom or group of atoms that have a
positive or negative charge
Vocabulary
ION
1 atom
2 or more atoms
Monatomic
Ion
Polyatomic
Ion
+
Na
NO3
-
Vocabulary
COMPOUND
2 elements
Binary
Compound
NaCl
more than 2
elements
Ternary
Compound
NaNO3
Vocabulary
CHEMICAL FORMULA
IONIC
COVALENT
Formula
Unit
Molecular
Formula
NaCl
CO2
Types of Bonds
IONIC
COVALENT
Bond
Formation
e- are transferred from
metal to nonmetal
e- are shared between
two nonmetals
Type of
Structure
crystal lattice
true molecules
Physical
State
solid
liquid or gas
Melting
Point
high
low
Solubility in
Water
yes
usually not
Electrical
Conductivity
yes
(solution or liquid)
no
Other
Properties
odorous
Types of Bonds/Compounds
Ionic Bonding - Crystal Lattice
Types of Bonds/Compounds
Ionic Bonding
Types of Bonds/Compounds
Covalent Bonding - True Molecules
Diatomic
Molecule
Types of Bonds/Compounds
Covalent Bonding
Types of Bonds/Compounds
METALLIC
Bond Formation
e- are delocalized among metal atoms
Type of Structure
“electron sea”
Physical State
solid
Melting Point
very high
Solubility in Water
no
Electrical Conductivity
yes (any form)
Other Properties
malleable, ductile, lustrous
Types of Bonds/Compounds
Metallic Bonding - “Electron Sea”
Determining Bond Type
Most bonds are
a blend of ionic
and covalent
characteristics.
Difference in
electronegativity
determines bond
type.
Determining Bond Type
Electronegativity
Attraction an atom has for a shared pair
of electrons.
higher e-neg atom  lower e-neg atom +
Determining Bond Type
Electronegativity Trend
Increases up and to the right.
Determining Bond Type
Nonpolar Covalent Bond
e- are shared equally
symmetrical e- density
usually identical atoms
Determining Bond Type
Polar Covalent Bond
e- are shared unequally
asymmetrical e- density
results in partial charges (dipole)
+


Determining Bond Type
Nonpolar
Polar
Ionic
How To Determine the Bond Type Using Electronegativity Differences
nonpolar
0 to 0.4 = ______________
covalent bond
polar
0.5 to 2.0 = _____________
covalent bond
ionic
Above 2.0 = _______________
bond
Practice Problems: Determine the type of bond that forms between
the atoms in the following compounds.
2.5 3.5
0.9 3.0
a) CO2
b) NaCl
1.0 = polar covalent
2.1 = ionic
2.5 2.1
c) CH4
0.4 = nonpolar covalent
Unit 4 – Bonding & Molecular Structure
II. Ionic Compounds
I
II
III
IV
Octet Rule
Octet Rule
Most atoms form bonds in order to
obtain 8 valence eFull energy level stability ~ Noble Gases
Ne
Lewis Structures
Electron Dot Diagrams used to show
bonding
Ionic Compounds – show transfer of e-
Lewis Structures
cation + anion = ionic compound
Cation or Anion?
How do you determine whether an atom
loses or gains electrons?
Look at the group #
What does that tell you?
Ask yourself, “is it easier to LOSE
______ electrons or GAIN ____
electrons to achieve the octet?”
Practice Problem
Aluminum and nitrogen
Al has __ valence electrons
- will it lose __ or gain __?
- Al becomes…
N has __ valence electrons
- will it lose ___ or gain __?
- N becomes…
Barium and Sulfur
Ba
[Xe] ____
6s
Which noble gases
will they look like
after they react?
S
[Ne] ___
3s
___ ___ ___
3p
After reacting…
Ba2+
[Xe] ____
6s
S2[Ne] ___
3s
___ ___ ___
3p
Looks
like
Argon
REVIEW - Common Ion Charges
1+
0
2+
3+ NA 3- 2- 1-
Naming Type I Binary Compounds
(single-charge cations)
Write the names of both ions, cation first.
Change ending of monatomic ions to -ide.
Examples:
SrS
AlBr3
Formulas for Type I Binary Compounds
(single-charge cations)
Write each ion, cation first. Don’t show
charges in the final formula.
Overall charge must equal zero.
If charges cancel, just write symbols.
If not, use subscripts to balance charges.
Some transition metals have only 1 charge:
Ag+
Zn2+
Cd2+
Criss-Cross Rule
Example: Aluminum Chloride
Step 1:
Al3+
Cl1-
Al 1
Cl 3
write symbols & charge of elements
Step 2:
criss-cross charges as subscripts
Step 3:
combine as formula unit
(“1” is never shown)
AlCl 3
Criss-Cross Rule
Example: Aluminum Oxide
Step 1:
Al3+
O2-
Step 2:
Al 2
O3
Step 3:
Al2O3
PRACTICE
Type I Binary Compounds WS (SingleCharge Cations)
Naming Type II Binary Compounds
(multiple-charge cations)
Since the metal ion can have more than one
charge, a Roman numeral is used to specify the
charge.
Determine the charge on the cation using the
charge on the anion
Example: NiBr2
NiBr2 = nickel (II) bromide
Common Cations and Anions
Naming Type II Binary Compounds
(multiple-charge cations)
PRACTICE PROBLEMS
1.
2.
3.
4.
5.
CuCl
HgO
Fe2O3
MnO2
PbCl2
copper (I) chloride
_________________
mercury (II) oxide
_________________
iron (III) oxide
_________________
manganese (IV) oxide
_________________
lead (II) chloride
_________________
Formulas for Type II Binary
Compounds (multiple-charge cations)
Roman numerals indicate the ion’s charge.
Examples:
1. Iron (III) chloride
2. Tin (IV) oxide
3. Lead (II) sulfide
FeCl3
________
SnO2
________
PbS
________
PRACTICE
Multiple-Charge Cations WS
Formulas for Ternary Compounds
(polyatomic ions)
Use parentheses to show more than one
polyatomic ion.
Example: Ca3(PO4)2
Naming Ternary Compounds
PRACTICE PROBLEMS
NaNO2
sodium nitrite
KClO2
potassium chlorite
Ba3(PO4)2
barium phosphate
Fe(OH)3
iron (III) hydroxide
NaHCO3
sodium bicarbonate
‘sodium hydrogen carbonate’
Criss Cross Rule – Ternary Compounds
Example:
Magnesium Phosphate
Step 1:
Mg2+
PO43-
Step 2:
Mg3
(PO4) 2
Step 3:
Mg3(PO4)2
More Ternary Compound Examples
1.
Zn3(PO4)2
________________
2.
(NH4)2CO3
________________
3.
Al2(SO4)3
________________
aluminum sulfate
4.
Na2SO4
sodium sulfate
____________________
5.
LiCN
lithium cyanide
____________________
6.
Ba(ClO3)2
7.
Cu(OH)2
________________
zinc phosphate
ammonium carbonate
barium chlorate
____________________
copper (II) hydroxide
PRACTICE
Ternary Compounds WS
Go Fish For An Ion
 Acceptable Matches
Na+ Na+ SO4 2Al3+ Cl- Cl- Cl Unacceptable Matches
Li+
Na+ SO4 2Al3+ O2-
Go Fish Scoring
Cation
#
Cations
Anion
#
Anions
Chemical
Formula
Cu2+
1
OH-
2
Cu(OH)2
Points
Unit 4 – Bonding, Nomenclature & Molecular
Structure
III. Molecular / Covalent
Compounds (Type III Binary
Compounds)
I
II
III
IV
Energy of Bond Formation
Potential Energy
based on position of an object
low PE =
high stability
Energy of Bond Formation
Potential Energy Diagram
attraction vs. repulsion
no interaction
increased
attraction
Energy of Bond Formation
Potential Energy Diagram
attraction vs. repulsion
increased
repulsion
balanced attraction
& repulsion
Energy of Bond Formation
Bond Energy
Energy required to break a bond
Bond
Energy
Bond
Length
Energy of Bond Formation
Bond Energy
Short bond = high bond energy
Naming Type III Binary Compounds
Prefix System
1. Less e-neg atom comes first.
2. Add prefixes to indicate # of atoms.
Omit mono- prefix on first element.
3. Change the ending of the second
element to -ide.
Naming Type III Binary Compounds
PREFIX
monoditritetrapentahexaheptaoctanonadeca-
NUMBER
1
2
3
4
5
6
7
8
9
10
Naming Type III Binary Compounds
PRACTICE PROBLEMS
CCl4
 carbon tetrachloride
N2O
 dinitrogen monoxide
SF6
 sulfur hexafluoride
Formulas for Type III Binary Compounds
PRACTICE PROBLEMS
arsenic trichloride
 AsCl3
dinitrogen pentoxide
 N2O5
tetraphosphorus decoxide
 P4O10
The 7 Diatomic Elements
Br2 I2 N2 Cl2 H2 O2 F2
H
N O F
Cl
Br
I
Naming Binary Compounds - REVIEW
Unit 4 – Bonding, Nomenclature &
Molecular Structure
IV. Acids
I
II
III
IV
Definition
Acids
Compounds that form H+ in water.
Formulas usually begin with ‘H’.
Examples:
HCl – hydrochloric acid
HNO3 – nitric acid
H2SO4 – sulfuric acid
Acid Nomenclature
Anion
Ending
Acid Name
-ide
hydro-(stem)-ic acid
-ate
(stem)-ic acid
-ite
(stem)-ous acid
Acid Nomenclature
ACIDS
start with 'H'
2 elements
3 elements
hydro- prefix
-ic ending
no hydro- prefix
-ate ending
becomes
-ic ending
-ite ending
becomes
-ous ending
Acid Nomenclature
HBr (aq)
2 elements, -ide

hydrobromic acid

carbonic acid

sulfurous acid
H2CO3 (aq)
3 elements, -ate
H2SO3 (aq)
3 elements, -ite
Acid Nomenclature
hydrofluoric acid
2 elements
 H + F-
HF (aq)
sulfuric acid
3 elements, -ic
 H+ SO42- H2SO4(aq)
nitrous acid
3 elements, -ous  H+ NO2- HNO2 (aq)
PRACTICE
Acids WS
Nomenclature Review Flow
Chart
Formula  Name?
Metal + Nonmetal?
(Except: NH4+)
Two Nonmetals?
Ionic
d,f-block
Pb,Sn
Multiple
Groups 1A, 2A, 3A
Ag+, Zn2+, Cd2+
Single
Covalent
Steps 1 & 4 ONLY
1. Write name of cation (metal)
2. Determine the charge on the metal by balancing the
(-) charge from the anion
3. Write the charge of the metal in Roman Numerals
and put in parentheses
4. Write name of anion
(Monatomic anions need –ide ending!)
Use Prefixes!!!
*Mono*
Hexa
Di
Hepta
Tri
Octa
Tetra
Nona
Penta
Deca
Name  Formula?
No Prefixes?
Ionic
Prefixes?
Covalent
1. Determine the ions present
and the charge on each
(Roman Numeral = cation
charge, otherwise use PT)
1. FORGET CHARGES!!!
2. Balance formula (criss-cross)
3. Do NOT reduce subscripts!
3. Reduce subscripts (if needed)
2. Use prefixes to determine
subscripts