Ch. 5: Molecules and
Compounds
Dr. Namphol Sinkaset
Chem 200: General Chemistry I
I. Chapter Outline
I.
II.
III.
IV.
V.
VI.
Introduction
Representing Compounds
Lewis Dot Symbols
Ionic Bonds and Nomenclature
Covalent Bonds and Nomenclature
Problems Involving Chemical Formulas
I. Introduction
• When elements form compounds, the
original properties of the elements are
lost.
I. To Lower Potential Energy
• A chemical bond is the force that holds
atoms together in a compound.
• But why would atoms want to join with
other atoms?
• It all comes back to positive-negative
attractions between particles in the
atom which lead to lower PE!
I. Two Main Ways to Lower PE
I. Metal + Nonmetal = Ionic
I. Nonmetal + Nonmetal =
Covalent
• Instead of transferring e-’s, covalent
bonding occurs via sharing of e-’s.
• Attraction to two nuclei lowers PE.
II. Chemical Formulas
• There are three types of
formulas.
 molecular: gives the actual
number of atoms of each
element in a molecule of a
compound (e.g. H2O2)
 empirical: gives the relative
number of atoms of each
element in a compound (e.g.
HO)
 structural: uses lines to
represent covalent bonds and
shows interconnectivity
II. Chemical Models
• Formulas lead to models which give an idea
of the 3-D shape of a molecule.
II. From Names to Models
III. Lewis Dot Symbols
•
•
•
Valence e-’s are the most important e-’s
in bonding.
Lewis dot symbols are a way to depict
the valence e-’s of atoms.
Lewis dot symbols have two parts:
1) element symbol: represents nucleus and
core e2) dots around symbol: represent valence e-’s
III. Lewis Dot Symbols
• The number of valence e- is given by the
element’s group number!!
III. The Octet Rule
• Noble gases are known for their lack of
reactivity – what do their e- configs have in
common?
• Lewis generalized bonding behavior by
observing that when atoms bond, they lose,
gain, or share e- to obtain 8 valence e-.
• Known as the octet rule (duet for H and He).
IV. Ionic Bonding
• In ionic bonding, metal transfers e- to
the nonmetal.
• Transferring e- achieves octet.
• Resulting ions attracted by +/- charge.
IV. Depicting Ionic Bonding
1. Draw Lewis dot structures for atoms
involved.
2. Use harpoons to indicate e- transfer.
3. Fill octet of nonmetal, drawing additional
Lewis dot structures as needed.
4. Use bracket notation on ions formed.
Charges should cancel.
5. Write formula of ionic compound.
IV. Some Practice
• e.g. Show the formation of the bonding
that occurs between magnesium and
chlorine using Lewis dot symbols.
IV. Energetics of Ionic Bonds
• Although transfer of e- achieves an
octet, ionic compounds are really stable
because of +/- attractions.
• The e- transfer process needs energy.
 1st IE of Na = 496 kJ/mole
 1st EA of Cl = -349 kJ/mole
• However, 411 kJ/mole of heat evolves
upon NaCl formation.
IV. Lattice Energy
IV. Electrical Conductivity
IV. Ionic Compounds Melt at
High Temperatures
• Why are such high
temperatures needed?
IV. Ionic Nomenclature
• Ionic compounds are named
systematically, broken into two groups.
IV. Type I Compounds
• Type I compounds are ionics that have a
metal from Groups 1 or 2 and a nonmetal
from Groups 14-17.
• Examples:
 NaCl = sodium chloride
 MgBr2 = magnesium bromide
 K2S = potassium sulfide
IV. Type I Compounds
• To get a formula from a name,
remember that a compound must be
neutral.
• Ion charges can be found by locating
the element on the periodic table.
• “The charge on one becomes the
subscript of the other.”
IV. Type I Compounds
• e.g. What are the formulas for sodium
nitride, calcium chloride, potassium
sulfide, and magnesium oxide?
IV. Transition Metals
• Transition metals are found in the
“Valley,” Groups 3-12, of the periodic
table.
• Transition metal cations often can carry
different charges, e.g. Fe2+ and Fe3+.
• Thus, a name like “iron chloride” is
ambiguous.
IV. Type II Compounds
• Type II compounds are ionics that have a
transition metal (Groups 3-12) and a
nonmetal (Groups 14-17).
• Examples:
 FeCl2 = iron(II) chloride
 FeCl3 = iron(III) chloride
IV. Type II Compounds
•
e.g. Give the correct name or formula
for the compounds below.
a)
b)
c)
d)
e)
MnO2
copper(II) chloride
AuCl3
molybdenum(VI) fluoride
Hg2Cl2
IV. Type II Compounds
• An archaic naming system uses
common names for transition metal
cations of different charge.
 Higher charge given –ic suffix
 Lower charge given –ous suffix
• FeCl3 = ferric chloride
• FeCl2 = ferrous chloride
IV. Additional Complications
• To make naming ionic compounds
harder, sometimes polyatomic ions are
involved.
• polyatomic ion: an ion composed of two
or more atoms
IV. Common Polyatomic Ions
IV. Oxyanion Families
• Oxyanions are anions that contain
oxygen and another element.
• There are families of oxyanions, and
they have a systematic naming system.
• Have either two- or four-member
families.
 e.g. NO2- and NO3 e.g. ClO-, ClO2-, ClO3-, and ClO4-
IV. Two-Member Families
• For a two-member family, oxoanion with
fewer O atoms is given the “–ite” suffix
while the one with more O atoms is
given the “–ate” suffix.
 e.g. NO2- = nitrite
and NO3- = nitrate
IV. Four-Member Families
• For the four-member families, the
prefixes “hypo-” and “per-” are used to
indicate fewer or more oxygen atoms.
• e.g. the chlorine oxoanions




ClO- = hypochlorite
ClO2- = chlorite
ClO3- = chlorate
ClO4- = perchlorate
IV. Hydrated Ionic Compounds
• Ionics with trapped
waters are are called
hydrates.
• Greek prefixes are
used to indicate #’s of
trapped waters.
• e.g. cobalt(II) chloride
hexahydrate.
IV. Naming Practice
•
e.g. Give names or formulas for the
following compounds.
a)
b)
c)
d)
e)
f)
Na2CO3
magnesium hydroxide
CuSO4·5H2O
CoPO4
nickel(II) sulfate
NaClO2
V. Covalent Bonding
• In covalent bonding, nonmetals share
some (or all) of their valence electrons
to achieve an octet.
• Nonmetals can share two, four, or six
electrons in what are known as single,
double, or triple bonds.
V. Single Bonds
• In H2O, two unpaired e-’s on each H
combine with the two unpaired e- sites
on O.
V. Bonding vs. Lone Pairs
• Bonding pairs are shared; lone pairs are
not.
• Bonding pairs are often represented
with a line between the two atoms.
V. Lewis Model & Diatomics
V. Double Bonds
• When more e-’s need to be shared to
reach an octet, a double bond is
possible.
V. Triple Bonds
• When even more e-’s need to be
shared, a triple bond is possible.
 e.g. molecular nitrogen, :N≡N:
• As the bond order increases, the bond
gets stronger and shorter.
V. Covalent Nomenclature
• For covalent compounds, many different
compounds can exist from the same
two elements.
 e.g. NO, NO2, N2O, N2O3, N2O4, N2O5!
• Therefore, we need a systematic
naming method.
V. Type III Compounds
•
•
Type III compounds are covalent (nonmetal
bonded to nonmetal).
Naming rules:
1) Element w/ lower group # is named 1st using the
normal element name EXCEPT when halogens
are bonded to oxygen.
2) If elements are in the same group, lower element
named first.
3) Second element is named using its root and the
“-ide” suffix.
4) #’s of atoms indicated with Greek prefixes
EXCEPT when there is only one atom of the first
element.
V. Greek Prefixes
V. Type III Compounds
• Some examples:




ClO2 = chlorine dioxide
N2O5 = dinitrogen pentoxide
S2Cl2 = disulfur dichloride
SeF6 = selenium hexafluoride
V. Naming Practice
•
e.g. Give the correct formula or name of the
compounds below.
a)
b)
c)
d)
e)
f)
g)
h)
i)
j)
k)
l)
CoCl3
dichlorine heptaoxide
SrO
magnesium hydroxide
carbon tetrachloride
MgSO4·7H2O
sodium hydride
V2O5
Ru(ClO4)3
NI3
titanium(IV) oxide
N2F2
VI. Composition of Compounds
• The ratio of elements in a compound is
given by its formula.
• We can calculate composition of
specific elements in different ways.
 Mass percent
 Inherent conversion factors
• We can also do the opposite: given
composition, determine formulas.
VI. Masses of Compounds
• Atomic masses are readily accessible via
the periodic table, e.g. H = 1.008 amu.
• Molecular masses or molecular weights
are calculated by adding up the masses
of each atom in the compound.
• Thus, molecular mass = sum of atomic
masses.
VI. Molecular Mass of Water
• The formula for water is H2O, so it is
comprised of 2 H atoms and 1 O atom.
VI. Mole Calculations
• Of course, we can use molar masses
and Avogadro’s number to calculate the
number of particles in a sample.
• e.g. How many water molecules are in a
sample of water that weighs 2100 g?
VI. Mass Percent
• As we know, elements account for a set
amounts by mass in a compound.
VI. Sample Problem
• e.g. Calculate the mass percent of
nitrogen in ammonium nitrate.
VI. Inherent Conversion Factors
• Formulas have conversion factors within
them to allow calculation of their
composition.
VI. Sample Problem
• How many grams of carbon are present
in 7.25 mL of butane (C4H10) if the
density of butane is 0.601 g/mL?
VI. Finding Formulas from
Mass Data
• Given elemental mass data of a
compound, it’s possible to find the
formula of the compound.
• Elemental analysis is a common test
performed on newly synthesized
compounds.
VI. Subscripts are Mole Ratios
• When finding formulas from mass data,
always go to moles of each element.
• Write a temporary formula using these
mole numbers.
• Divide by the smallest mole number to
get to empirical formula.
VI. Adjusting Subscripts to
Whole Numbers
VI. Empirical to Molecular
• The molecular molar mass is always a
whole number multiple (n) of the
empirical molar mass.
• Use this n to convert empirical to
molecular formula.
molecular molar mass
n
empirical molar mass
VI. Sample Problem
• e.g. The carcinogen benzo[a]pyrene
(MW = 252.30 g/mole) is found to be
95.21% C and 4.79% H by mass. What
are its empirical and molecular
formulas?
VI. Combustion Analysis
• Empirical formulas of compounds containing
C, H, and one other element can be found via
combustion analysis.
• When sample is burned in O2, all the C
becomes CO2 and all the H becomes H2O.
VI. Sample Problem
• A 12.01 g sample of tartaric acid
(comprised of only C, H, and O) was
analyzed via combustion. If 14.08 g
CO2 and 4.32 g H2O are produced, find
the empirical formula of tartaric acid.