Chapter 5
Names and Formulas
of Compounds
Homework
 Assigned Problems (odd numbers only)
 “Questions and Problems” 5.1 to 5.61 (begins
on page 131)
 “Additional Questions and Problems” 5.71 to
5.95 (page 157-158)
 “Challenge Questions” 5.97 to 5.103, (page
158-159)
Octet Rule and Ions
 Compounds are pure substances
 The result of a combination of two or
more elements held together by
chemical bonds
 Chemical bonds are the attractive forces
that hold atoms or ions together
 They can be ultimately broken down into
two or more simpler substances:
Elements
Octet Rule and Ions
 Two types of attractive forces
 Ionic: Involves the transfer of one (or more)
electrons from one atom (or group) to another
 For example, NaCl
 Covalent: When two or more atoms share one
or more electrons between them
 For example, HF
 In compounds with covalent bonds it is the
outermost electrons involved in the chemical
bonding
Octet Rule and Ions
 When a compound forms, the atoms
must lose, gain, or share electrons to
produce a noble gas electron
configuration The octet rule
 When the sodium atom loses its only
valence electron, it obtains the
electron configuration of its nearest
noble gas: Neon
Positive Ions
 Form when an electron or electrons are lost
from a metal
 Named with element name, then add “ion”
 Atom becomes charged
 The charge on an ion is equal to the number of
the electrons lost


Sodium Ion
Na  Na  e
2

Magnesium Mg  Mg
 2e Magnesium Ion
Sodium
Aluminum
Calcium
 Aluminum Ion
 3e
2
 Calcium Ion
3
Al  Al
Ca  Ca
 2e
Positive Ions
 Ionic bonding involves
transferring one or
more electrons
between two or more
atoms
 Produces a (+) charged
atom: cation
 Metals in groups IA, IIA,
IIIA easily lose
electrons to acquire the
noble gas electron
configuration
Negative Ions
 Form when an electron or electrons are gained
 Named with root of parent atom and adding -ide to
the end
 Atom becomes charged
 The charge on an ion is equal to the number of
the electrons gained

 Fluoride
Fluorine
Bromine
Oxygen
Sulfur
Fe F


Br  e  Br

2
O  2e  O

2
S  2e  S
Bromide
Oxide
Sulfide
Negative Ions
 Ionic bonding and the
transfer of electrons
also produces a (-)
charged atom: anion
 Nonmetals in groups
VA, VIA, VIIA will gain
the necessary number
of electrons to acquire
the noble gas
electronic configuration
Ionic Compounds
 Compounds which are held together by
the attraction of positive and negative ions:
ionic compounds
 Solid crystals formed by a very ordered
packing of oppositely charged ions
 Most ionic compounds are composed of a
metal and a nonmetal
 High melting temperatures
Ionic Compounds
 Solids
 They do not exist as
single molecules
 The formula represents
the simplest ratio that
these atoms combine
together
 Ions are packed
together into a “lattice”
 Held strongly together,
high melting
temperature
Cl–
Na+
Cl–
Cl–
Na+
Cl–
Na+
Cl–
Na+
Cl–
Na+
Na+
Fig4_21a
Charge Balance in Ionic
Compounds
 Binary ionic compounds are composed of
only two elements (metal and nonmetal)
 The symbol of the cation always
precedes the symbol of the anion
 The sum of the positive charges (cation)
must equal the sum of the negative
charges (anion)
 Net charge is zero
 Subscripts written as whole numbers
indicate the number of each ion in the
formula unit
Subscripts in Formulas
 Sodium Chloride
 Formed from
sodium and
chlorine atoms
 An ionic bond
forms consisting
of a sodium ion
(+ charge) and a
chloride ion
(- charge)
 Each sodium loses
one electron to
achieve an octet
 Each chlorine atom
gains one electron
to achieve an octet
 Formula is NaCl
Subscripts in Formulas
 Magnesium
Chloride
 Formed from
magnesium and
two chlorines
 An ionic bond
forms consisting
of a magnesium
ion (2+ charge)
and
two chloride ions
(- charge each)
 Each magnesium
loses two electrons
to achieve an octet
 Each chlorine atom
gains one electron
to achieve an octet
 Formula is MgCl2
Writing Ionic Formulas from
Ionic Charges
 Subscripts in a formula represent the number of
positive and negative ions
 Write the formula for the ionic compound containing
Na+ and N3-
Na+
N
••
• • 3-
N
••
••
••
•
Na•
••
••
Na+
••
Na•
••
•
Na•
Gains
Each loses 1e-
3e-
•
•
•
Na+
Net charge: 3(+1) + 1(3-)=0
Formula: Na3N
Naming and Writing Ionic Formulas
 Ionic Compounds Containing Two Elements
 Compounds containing a metal and a
nonmetal are binary ionic compounds
 Single Cation Metals: Form one positive ion
 Multiple Cation Metals: Form more than one
positive ion
 The systematic naming uses the name of the
cation first, followed by the name of the anion
 Subscripts in the formula are not included in
the name
Types of Metal Ions
 Single Cation Metals
 Form only one type of ion (one possible
charge)
 Main group metals in groups IA, IIA, and
some IIIA
 i.e. Sodium only forms one ion (Na+) in
chemical reactions
 Determine charge by position on the
periodic table (also see table 5.3 on
page 136)
Naming Ionic Compounds Containing
Two Elements
 Name metal cation first, name nonmetal
anion second
 Single metal cation name is the metal
name only, drop the word “ion”
 Nonmetal anion named by changing the
ending on the nonmetal name to -ide
Name of metal______name of nonmetal + -ide
Examples
 NaI
KCl
Potassium Chloride
 Sodium Iodide
Na3P
 CaF2
 Calcium Fluoride Sodium Phosphide
Rb2S
 Li2O
Rubidium Sulfide
 Lithium Oxide
Mg3N2
 AgCl
Magnesium Nitride
 Silver Chloride
Types of Metal Ions
 Multiple Cation Metals
 Form two or more types of ions (variable
possible charge)
 Transition metals in groups 3B to 12B,
and some 4A and 5A
 For example, Iron forms two ions (Fe2+
and Fe3+) in chemical reactions
 Determine charge by the “stock system”
for naming ions
 The metal name followed by a Roman
numeral in parentheses to indicate its
charge (see table 5.4 on page 137)
Multiple Cation Metal Compounds
 Metal listed first in formula & name (same
order as for Type I compounds)
 Determine metal cation charge from anion
charge
 Use the metal name (cation) first followed by
a Roman numeral in parentheses to indicate
its charge
 Common multiple cations in Table 5.4,
page 137
 Nonmetal anion named by changing the
ending on the nonmetal name to -ide
Determining the Charge of the Cation
from the Anion
Determine the charge of Cu in Cu2O
Write the name of the compound
1)Determine the charge of the cation from
the anion
 Cu2O - the nonmetal anion is O, since it is
in Group 6A, its charge is -2
 Since there are 2 Cu ions in the formula
and the total positive charge is +2, divide
by the number of cations so each Cu has
a +1 charge
Naming Ionic Compound with
Variable Charge Metal Ions
2) Name the cation by its element name and
use a Roman numeral in parenthesis to
indicate its charge
 Copper (I)
3) Name the anion by changing the last part of
its element name to –ide
 Oxygen
oxide
4) Write the name of the cation first and the
name of the anion second
copper (I) oxide
Examples
 FeI3
 1(?)+3(-1)=0
 Iron (III) Iodide
 Cu2O
 2(?)+1(-2)=0
 Copper (I) Oxide
 SnBr2
 1(?)+2(-1)=0
 Tin (II) Bromide
Examples
• SnI4
• 1(?)+4(-1)=0
• Tin (IV) iodide
• HgO
• 1(?)+1(-2)=0
• Mercury (II) Oxide
• MnCl2
• 1(?)+2(-1)=0
• Manganese (II) Chloride
Writing Formulas from the Name of
an Ionic Compound

Usually involves a metal and a nonmetal
1) Identify the cation and the anion
2) Balance the charges to write the formula
 If it is a multiple cation metal, the Roman
numeral determines the charge of the
cation
3) When writing the formula, take the name of
the cation first, followed by the name of the
anion
Writing Formulas from the Name of
an Ionic Compound
 Compound name is lithium chloride
 Li+ and Cl- are the ions
 Balance the charges


Li
Cl
1(1 )  1 (1)  0
 Write the formula
 LiCl is the formula using the
subscripts from the charge balance
Writing Formulas from the Name of
an Ionic Compound
 Compound name is iron (III) oxide
 Fe3+ and O2- are the ions
 Balance the charges
Fe3
O2
2 (3 )  3 (2)  0
 Write the formula
 Fe2O3 is the formula using the
subscripts from the charge balance
Polyatomic Ions
 A group of atoms covalently bonded
together into a single unit
 The unit obtains a charge
 Most PA ions are negatively charged
 Oxyions (anions): P, S, C, or N
covalently bound to one or more
oxygens
 Never occur independently, always
associated with ions of opposite charge
 Only one PA is positively charged
 ammonium ion
Naming Polyatomic Ions
 Must memorize name, formula and
charge (Table 5.6 on page 142).
Look for relationships between ions
 Oxyions: The number of oxygen
atoms bonded to the same element
(i.e. P, S, or N) will determine the
name of the ion
 ~ate is most common
 ~ite has one less oxygen bonded
Polyatomic Ions
 ~ate, ~ite pairs of ions
 The ion in the pair with the most oxygens is always the




~ate ion
The ion in the pair with one less oxygen is always the
~ite ion
Ion pair with a -3 charge
 phosphate PO43-, phosphite PO33Ion pair with a -2 charge
 sulfate (SO42-), sulfite (SO32-)
Ion pair with a -1 charge
 nitrate (NO3-), nitrite (NO2-)
Polyatomic Ions
 Group 7A elements can form more than two




types of polyatomic ions (oxyions)
Additional prefixes are used to differentiate
the ions
See page 142 (class text) and page 330 (lab
text)
The number of oxygens attached to the
central atom has an effect on the name of
the ion
e.g. Polyatomic ions of chlorine, bromine
and iodine
Polyatomic Ions
 Example: Polyatomic ions of chlorine
1) -ate ion
 chlorate = ClO32) -chlorate ion with 1 more O than chlorate, use perprefix
 perchlorate = ClO43) - chlorate ion with 1 less O, use -ite suffix
 chlorite = ClO24) -chlorite ion with 1 less O, use hypo- prefix
 hypochlorite = ClO-
Writing Formulas for Compounds
Containing Polyatomic Ions
 Formulas are written like binary ionic




compounds
Consider polyatomic ions as single units
with a certain charge
Obtain the correct ratio of cation to anion
to achieve a net charge of zero
Use parentheses if more than one of the
same PA unit is needed in a formula
Use subscripts to indicate the number of a
particular ion in a formula
Writing Formulas for Compounds
Containing Polyatomic Ions
 Compound name is magnesium carbonate
 Mg2+ and CO32- are the ions
 Balance the charges
2
2
Mg
CO3
1(2 )  1(2)  0
 Write the formula
 MgCO3 is the formula using the
subscripts from the charge balance
Writing Formulas for Compounds
Containing Polyatomic Acids
 Compound name is calcium nitrate
 Ca2+ and NO3- are the ions
 Balance the charges

Ca2
NO 3
1 (2 )  2(1)  0
 Write the formula
 Ca(NO3)2 is the formula using the
subscripts from the charge balance
Naming Compounds Containing
Polyatomic Ions
 Named the same way as binary ionic





compounds
Positive ion (metal) name is written first
Polyatomic ions name follows the metal
No prefixes are used in the name
Cation: Check to see if metal is single or
multiple cation
Use the name of the PA ion given in table
5.6 on page 142
Writing Formulas for Compounds
Containing Polyatomic Ions
 Compound name is iron (III) sulfate
 Fe3+ and SO42- are the ions
 Balance the charges
2
Fe3 
SO4
2(3  )  3(2 )  0
 Write the formula
 Fe2(SO4)3 is the formula using the
subscripts from the charge balance
Writing Formulas for Compounds
Containing Polyatomic Ions
 Compound name is ammonium phosphate
 NH4+ and PO43- are the ions
 Balance the charges

3
NH 4
PO 4
3 (1 )  1 (3 )  0
 Write the formula
 (NH4)3PO4 is the formula using the
subscripts from the charge balance
Naming Compounds Containing
Polyatomic Ions
• CaSO4
• calcium sulfate
• Ca2+ and SO42-
• Li2CO3
• lithium carbonate
• Li+, CO32-
• Al(NO3)3
• aluminum nitrate
• Al3+, NO3-
Summary of Naming Ionic
Compounds
 Summary of guidelines when writing binary ionic
compound
 The symbol of the cation always precedes the
anion
 The sum of the positive charges must equal the
sum of the negative charges:
A net charge of zero
 Whole numbers are written as subscripts to
indicate the number of each ion in the formula
Covalent Compounds and
Their Names
 Involves a bond between two nonmetals
 Bonds occur between similar or identical
atoms
 Nonmetals such as O, Br, or N do not tend
to lose electrons (tend to gain them)
 Electrons are shared and not transferred
between atoms forming covalent bonds
 Exist as individual molecule
Formation of a Hydrogen Molecule
 The simplest covalent




bonding condition
Hydrogen has one 1s
electron
H atom requires one
additional electron to
obtain the stable noble
gas configuration of
helium
Each H atom contributes
its one electron
The electron pair shared
by the two atoms,
forming diatomic
hydrogen H2
Formation of Octets in
Covalent Molecules
 Two identical nonmetal atoms
 Each atom will share valence electrons with the
other
 The shared pair of electrons allow each atom to
achieve a stable noble gas configuration
 This configuration can be achieved by a single,
double, or triple shared pair of electrons
Formation of Octets in
Covalent Molecules
 Two identical
nonmetal atoms
 H atom exists as a
diatomic molecule by
achieving a duet of
electrons
 F, Cl, Br, I, O, N exist
as diatomic molecules
by achieving an octet
of electrons in their
valence shells
Sharing Electrons Between Atoms
of Different Elements
 Two nonidentical
nonmetal atoms
 The number of
covalent bonds an
atom forms will equal
the number of
electrons needed to
form a noble gas
configuration
 Each vacancy +
unpaired electron
combination in the
valence shell can be
used to form a twoelectron bond
 Each atom will share
valence electrons with the
other forming a shared
pair of bonding electrons
(achieves a stable noble
gas configuration)
Names and Formulas of
Covalent Compounds
 Molecular binary compounds
 Composed of two nonmetal elements
 Naming a compound
 Use the full (element ) name for the first
nonmetal
 Add the –ide ending to the full name of the
second nonmetal
 Second nonmetal named like the nonmetal in
binary ionic compounds (anion)
 Indicate the number of atoms by adding
numerical prefixes
Names and Formulas of Covalent Compounds
(table 5.11, page 151)
Subscript
1
2
3
4
5
6
7
8
9
10
Prefix used
mono~
di ~
tri ~
tetra ~
penta ~
hexa ~
hepta ~
octa ~
nona ~
deca ~
(Usually omitted on the first atom)
Names and Formulas of
Covalent Compounds
 In ionic compounds the subscripts are not mentioned in
the name
BaCl2
Na2SO4
barium chloride
sodium sulfate
barium dichloride
disodium sulfate
 Many compounds can exist for many pairs of nonmetallic
elements (i.e. nitrogen and oxygen)
NO
NO2
N 2O
nitrogen monoxide
nitrogen dioxide
dinitrogen monoxide
Molecular Binary (Covalent) Compounds
 Naming binary molecular compounds from a formula
 Name the first nonmetal by its element name
 Name the second nonmetal by adding the –ide
suffix
 Add the prefixes to indicate the number of atoms
 Whenever the vowels a and o or o and o appear
together, the first vowel is dropped from the prefix
for easier pronunciation
Cl2O
P4O6
dichlorine monooxide
tetraphosphorous hexaoxide
dichlorine monoxide
tetraphosphorous hexoxide
Examples
 IF5
 iodine pentafluoride
 B2O3
 diboron trioxide
 NO3
 nitrogen trioxide
Examples
• AsCl3
• arsenic trichloride
• CO2
• carbon dioxide
• CO
• carbon monoxide
Molecular Binary (Covalent) Compounds
 When writing a formula from the name of
a binary molecular compound
 You must know definition of the
numerical prefixes used in naming
covalent compounds (see table 5.11)
 (you MUST memorize these
prefixes)
1) Write the symbols in order the
elements appear in the name
2) Identify the prefixes with the
appropriate subscripts
Acids (Section 14.1)
 Produce H+ when dissolved in water
 Composed of H+ (cation) and an
anion
 Binary acids have H+ cation and a
nonmetal anion
 Oxyacids have H+ cation and a
polyatomic anion (contain oxygen)
Naming Binary Acids
 Use the prefix hydro- before the root name
of the element
 Add the suffix -ic and the word acid to the
root name for the element
 Example: HCl
 hydrochloric acid
 Example: HI
 hydroiodic acid
Oxyacids
 Produce H+ and a polyatomic ion when dissolved in
water
 Composed of hydrogen, oxygen, and another nonmetal
 Use the root name of the polyatomic ion
 If it ends in -ate use the suffix -ic acid
 If it ends in -ite use the suffix -ous acid
 Example: H2SO4 (from SO42- ,sulfate ion)
 sulfuric acid
 Example: H2SO3 (from SO32- ,sulfite ion)
 sulfurous acid