Ch. 9 – Chemical Names and Formulas
I.
Ion Formation
Ionic Formulas
Ionic Nomenclature
I
II
III
IV
Ionic Bonds
When oppositely charged ions attract,
electrostatic force that holds them together
= ionic bond
Compounds containing ionic bonds =
ionic compounds
Electrons are transferred from cations to
anions
Typically bonds between metals and
nonmetals, however polyatomic ions also
involved.
A. Vocabulary
Chemical Bond
attractive force between atoms or ions
that binds them together as a unit
bonds form in order to…
fulfill octet rule
increase stability
Octet Rule
Atoms will gain or lose electrons so that
they have 8 electrons in their highest
energy level
The noble gases already have a full octet,
8 valence electrons, so are chemically
stable.
A. Vocabulary
Cation
Positively charged ion formed when an
atom loses one or more valence eNumber of protons stays the same, but
less electrons gives + charge
Loses an e-
Sodium ion
A. Vocabulary
Anions
Nonmetals easily gain e- to form negative
ions to get to 8 valence eName is changed to root + -ide
Gains an e-
Chloride ion
A. Vocabulary
ION
1 atom
2 or more atoms
Monatomic
Ion
Polyatomic
Ion
+
Na
NO3
-
Monatomic Ions
Ions formed from a single atom of an
element
Most main group metals form one type of
monatomic ion
Group 1 metals: +1 ions (Na+, Li+…)
Group 2 Metals: +2 ions (Mg2+, Ca2+…)
Group 13 Metals: +3 ions (Al3+, Ga3+…)
Monatomic Ions
Main group nonmetals form one type of
anion
Group 15: -3 ions (N3-, P3-)
Group 16: -2 ions (O2-, S2-…)
Group 17: -1 ions (F-, Cl-…)
Monatomic Ions
Some metals, mainly transition metals,
can form more than one type of ion
Iron: Fe2+ or Fe3+
Copper: Cu+ or Cu2+
Lead: Pb2+ or Pb4+
D. Common Ions
B. Formula Unit
Chemical Formulas
Chemical formulas
identify the elements present in a
compound using element symbols
the number of atoms of each element is
indicated with numbers in subscript
Fe2O3
3 Oxygen
2 Iron atoms
atoms
Binary Ionic Compounds
Binary ionic compounds are ionic
compounds formed between two
elements.
Always formed between a metal (positive
ion) and nonmetal (negative ion)
Metal is written first in formula and then
nonmetal
D. Ionic Formulas
Writing Ionic Formulas:
Calcium chloride
Ca2+
Cl11
2
charges do not cancel, must criss-cross charges
Rewrite as complete formula without charges
CaCl2
D. Ionic Formulas
K+ F-

KF
Mg2+ N3-

Mg3N2
Ba2+ Cl-

BaCl2
D. Ionic Formulas
Ca2+ O2- 
CaO

Al2S3
Mg2+ Br- 
MgBr2
Al3+ S2-
C. Lewis Structures
Ionic – show transfer of electrons
G. Ionic Nomenclature
Naming Binary Ionic Compounds
Write names of both ions, cation (metal) first
Change ending of monatomic anions
(nonmetal) to -ide
Use Roman numerals to show the ion’s
charge if more than one is possible
G. Ionic Names with Type II Cations
You must write the charge in parentheses using
Roman numerals. To determine charge know that
overall charge of compound = 0
Cr2O3
Chromium (III) oxide
O: 3 x -2
=-6
Formula:
+3
Cr: 2 x ___
= +6
Element: # atoms x charge = total charge
CrO
Chromium (II) oxide
O: 1 x -2
=-2
+2
Cr: 1 x ___
= +2
Name the following
Compounds
CaCl2
Calcium Chloride
Al2O3
Aluminum Oxide
Na3N
sodium nitride
KI
Potassium Iodide
Name the following cont.
FeN
Iron (III) nitride
Cu2S
Copper (I) sulfide
ZnCl2
Zinc chloride
V3P5
Vanadium (V)
phosphide
Manganese (VI) oxide
MnO3
Write the formula for
Strontium nitride Sr3N2
Lithium sulfide
Li2S
Gallium bromide GaBr3
Barium oxide
BaO
Write the formulas
Tin (IV) oxide
SnO2
Gold (III) Chloride
AuCl3
Chromium (II)
Nitride
Cr3N2
Mercury (I) sulfide
(Hg2)2S
F. Ionic Formulas with Type II Cations
Copper (II) bromide
Cu2+ + Br -

CuBr2
Tin (IV) oxide
Sn4+ + O2-
 Sn2O4
Manganese (II) chloride
Mn2+ + Cl-

MnCl2
 SnO2
E. Polyatomic Ions
Polyatomic Ions
Ions made of more than one atom
Acts as an individual ion and its charge
applies to the entire group of atoms
NEVER change the subscripts – add
parentheses and subscripts outside, if
necessary
Listed on the back of your periodic table
E. Ionic Formulas with PA Ions
potassium chlorate
K+ ClO3-

KClO3
magnesium nitrate
Mg2+ NO3-

Mg(NO3)2
ammonium phosphate
NH4+ PO43-

(NH4)3PO4
E. Ionic Formulas with PA Ions
calcium oxalate
Ca2+ C2O42-

CaC2O4
aluminum perchlorate
Al3+ ClO4-

Al(ClO4)3
strontium phosphate
Sr2+ PO43-

Sr3(PO4)2
Nomenclature with PA
When naming compounds with polyatomic
ions use the polyatomic ions name and
follow all other rules for ionic compounds
Ca3(PO4)2  calcium phosphate
Fe(NO3)3  iron (III) nitrate
G. Ionic Nomenclature
CaBr
calcium bromide
Na2CO3
sodium carbonate
NH4OH
ammonium hydroxide
G. Ionic Names with Transition Metals
Cr2(SO4)3
Chromium (III) sulfate
Cu(NO3)2
Copper (II) nitrate
FeCl3
iron(III) chloride
Ch. 9 – Chemical Names and Formulas
II.
Covalent Bond Formation
Covalent Compound
Names & Formulas
I
II
III
IV
A. What is a covalent bond?
A chemical bond that results from the
sharing of electrons
Molecule = two or more atoms
that are held together by
covalent bonds
H2O
Majority of covalent bonds form between
nonmetals (CLOSE together on periodic
table)
B. Examples:
Which of the following are covalent
compounds?
NaBr
SiO2
CO2
AlCl3
CH4
C. Covalent Bonding Formation
Diatomic molecule
molecule containing only two
atoms
Some elements always exist this
way because they are more
stable than the individual atoms
Cl2
D. Diatomic Elements
The Seven Diatomic Elements
Br2 I2 N2 Cl2 H2 O2 F2
H
N O F
Cl
Br
I
E. Molecular Nomenclature
 Prefix System (binary molecules)
1. Add prefixes to indicate # of atoms.
Omit mono- prefix on first element.
2. Change the ending of the
second element to -ide.
3. Second element ALWAYS gets a
prefix.
E. Molecular Nomenclature
PREFIX
monoditritetrapentahexaheptaoctanonadeca-
NUMBER
1
2
3
4
5
6
7
8
9
10
E. Molecular Nomenclature
CCl4
carbon tetrachloride
N2O
dinitrogen monoxide
SF6
sulfur hexafluoride
E. Molecular Nomenclature
arsenic trichloride
AsCl3
dinitrogen pentoxide
N2O5
tetraphosphorus decoxide
P4O10
Lewis Structures
The Lewis Structures of Covalent
compound represents elements in a
compound, their valence electrons and
how they are shared, and how the
elements orient themselves around each
other.
A. Drawing Lewis Structures
1. Count ALL Valence electrons on all atoms
in the molecule.
 For an anion ion, add one electron for
each negative charge.
 For a cation, subtract one electron for
each positive charge.
A. Drawing Lewis Structures
2. The atom with the least amount is central atom
and place the other atoms around the central
atom.
 Hydrogen is never central atom
 Draw a line connecting the peripheral
atoms to the central atom
 Each line represent 2 electrons
 Check for octet
A. Drawing Lewis Structures
3. Place pairs of valence electrons around
each peripheral atom, except
hydrogen, until octet is reached.
4. If any electrons remain place around the
central atom until octet is reached.
5. If central atom still does not have an
octet, use a lone pair of electrons on a
neighboring atom to form a multiple
bond to the central atom.
Examples
CF4
CO2
 HCN
ClONH4+
B. Drawing Lewis
Diagrams
 CF4
1 C × 4e- = 4e+
4 F × 7e = 28e32e2
16 pairs of e- 4 pairs of e12 pairs of e-
F
F C F
F
A. Drawing Lewis
Diagrams
 CO2
1 C × 4e- = 4e+
2 O × 6e = 12e16e2
8 pairs of e-2 pairs of e6 pairs of e-
O C O
B. Polyatomic Ions
 To find total # of valence e-:
Add 1e- for each negative charge
Subtract 1e- for each positive charge
 Place brackets around the ion and label
the charge
B. Polyatomic Ions
 ClO41 Cl × 7e- = 7e+
4 O × 6e = 24e31e+ 1e32e2 = 16 e- pairs
- 4 e- pairs
12 e- pairs
O
O Cl O
O
B. Polyatomic Ions
 NH4+
1 N × 5e- = 5e+ 4 H × 1e = 4e
9e- 1e8e-
H
H N H
H
2 = 4 pairs of e-4 pairs of e0 pairs of e-
C. Resonance Structures
 Molecules that can’t be correctly
represented by a single Lewis diagram
 Actual structure is an average of all the
possibilities
 Show all possible structures separated by
double-headed arrows
C. Resonance Structures

SO3
O
O S O
O
O S O
O
O S O
A. Octet Rule
 Remember…
Most atoms form bonds in order to
have 8 valence electrons
D. Octet Rule
 Exceptions:
F
F
Hydrogen  2 valence e
F
B
F
Groups 1,2,3
get
2,4,6F
valence e
F
S
H
N
O
O
H
Expanded octet  more than 8
F
valenceVery
e (e.g.
S, P, Xe)
unstable!!
F
F
-
-
-
E. Drawing Lewis
Diagrams
 BeCl2
1 Be × 2e- = 2e+
2 Cl × 7e = 14e16e2
8 pairs of e-2 pairs of e6 pairs of e-
Cl Be Cl
Expanded Octet
Some atoms may have more than 8
electrons around them. Follow the rules for
normal Lewis Diagrams
Nonmetals and metalloids in rows 3 or
higher can have expanded octets
Examples
SF6
PCl5
XeF4
ICl3
E. Drawing Lewis
Diagrams
 SF6
1S× 6e- = 6e+
6F× 7e = 42e48e2
24 pairs of e- 6 pairs of e18 pairs of e-
F
F
F S F
F
F
Octet Deficient Elements
Beryllium and boron and do not need an
octet
Beryllium (Be) covalently bonds and only
needs 4 electrons around it
Boron (B) needs 6 around it
BCl3
BeF2