Chemistry Lecture Notes
Ionic and Molecular Compounds
New evidence for inorganic origin of oil
The origin of petroleum has long been an issue of debate. Although
many people believe crude oil formed from plant and animal material in
near-surface sedimentary rocks at high temperatures and pressures, a
abiogenic origin has also been proposed. Proposed in the 1950s, the theory
attributes oil formation to inorganic carbonate rocks at high temperatures
and pressures found only at great depths.
In the latest development in this debate, researchers at Gas Resources,
Corp., Houston, and the Joint Institute of Earth Physics, Moscow, have
predicted the thermodynamic conditions under which the hydrocarbons
found in crude oil form and tested those conditions in the lab. They reacted
iron oxide, marble (CaCO3), and water at condition reaching 1,500ºC and
50,000 atm. Hydrocarbons ranging from methane to decane were formed
in proportions that mirror naturally occurring petroleum. With the
exception of methane, hydrocarbons did not form at pressures below
30,000 atm, which corresponds to about 100 km below the Earth’s surface.
Taken together, the theoretical and experimental results make the biogenic
theory untenable, the researchers conclude.
Lewis (electron-dot) symbols: show the valence electrons of atoms or ions
Li
Be
B
C
N
O
Be
Ne
2
2
Li
F
O
F
octet rule: atoms tend to gain, lose, or share electrons until they reach
the nearest noble gas configuration
Ionic compounds

+
Na
+
_
Na+Cl
Cl
Ionic “bond”: electrostatic attraction between positive ions (cations)
and negative ions (anions).
metal + nonmetal  ionic compound (e.g., NaCl, MgF2, K2S, CaO, AlBr3)
Na + Cl
low IE
low EA
gives up e–
high IE
high EA
takes e–
Na + Cl
transfer of electrons
electrolytes!
Metals tend to lose electrons to form
cations:
alkali metals:
+1
alkaline earths: +2
group 3A:
+3
transition metals: variable
Nonmetals tend to gain electrons to
form anions:
halogens:
-1
chalcogens: -2
group 5A
-3
P3–
We can usually predict the formulas of ionic compounds
 empirical formula (formula unit) only
Rules:
1. cation written first
2. sum of all charges = 0
3. empirical formula has smallest set of whole numbers
Li, Br 
K, P 
Na, O 
Al, F 
Mg, Cl 
Ca, O 
Be, N 
Be+2 N-3
Be3N2
but Ca, C  Ca+2 C-4
Ca4C2
Ca2C
Molecular compounds: covalent bonds

+
C
O2
CO2
Two or more atoms bound tightly together.
nonmetal + nonmetal  molecular compound
(covalent bonds)
e.g., N2O, CO2, H2S, PCl3, SO3
nonelectrolytes!
Can’t always predict formula: N2O, NO, NO2, N2O5, etc.
Covalent bond: attraction of two nuclei to one or more shared pairs of
electrons
nonmetal + nonmetal  covalent (molecular) compound
H + H
equal attraction
for electrons
Cl + Cl
H H = H H = H2
covalent bond; shared electrons
each atom “sees” two electrons (He)
Cl Cl or Cl
each atom
has an octet
H + Cl
H Cl
duet octet
2H + O
H O H
Cl
lone pair electrons
bonding electrons
Multiple bonds
O2
2 O
O O
O O
double bond
N2
2 N
N N
N N
triple bond
octets
Electronegativity and bond polarity
Electronegativity: ability of an atom in a compound to attract electrons to
itself
• ~ follows IE, EA
• arbitrary scale (Pauling)
most
C
increases
least
+
H H
equal sharing
of electrons
-
H Cl
dipole moment
polar bond
unequal sharing
of electrons
H—H
C = 2.1 2.1
Cl—Cl
C = 3.0 3.0
DC = 0  equal sharing of electrons
= nonpolar covalent bond
+ –
H—Cl
C = 2.1 3.0
DC = 0.9  unequal sharing of electrons
= polar covalent bond
Na+Cl–
C = 0.9 3.0
DC = 2.1  transfer of electrons
= ionic bond
generally:
when DC < 1.9 covalent
> 1.9 ionic
(approximate!)
Empirical, molecular, and structural formulas
Empirical formula: smallest whole number ratio of atoms in a compound
e.g. CH2O
Molecular formula: actual number of atoms in a molecule of the compound
CH2O
C2H4O2
C6H12O6
formaldehyde
acetic acid
glucose
all have same empirical formula
Structural formula: shows how atoms are bonded together
O
H
C
H
H
formaldehyde
H O
H C
C
H O O
O H
H
acetic acid
H C
C
H
H
glyceraldehyde
C
O
H
H C
H C
OH
C
H C
OH
H C
OH
H C
OH
H C
OH
H C
OH
H C
OH
H C
OH
H C
OH
H
glucose
isomers
CH3CO2H
HOCH2CHO
OH
Condensed formulas
O
H
fructose
Drawing Lewis structures
1. Sum all valence electrons.
•add one electron for each negative charge
•subtract one electron for each positive charge
2. Draw single bonds from the central atom to the outer atoms.
3. Place pairs of electrons on the outer atoms until they have octets.
•exception: H (duet)
4. Place any remaining electrons on the central atom.
5. If the central atom has less than an octet, form multiple bonds with
pairs from the outer atoms.
NI3
SO42–
CH2O
PCl5
SeF6
C2H4
SO2
NO+
C2H5OH
NO3–
HCN
CH3CO2H
Formal charge
Divide the electrons in each bond equally between the two atoms,
then compare each atom to its normal valence.
H
e.g., OH-
O H
7 e–
1 e–
(redo previous page examples)
NH4+
H N H
4 e–
H
When more than one Lewis structure is possible, the better (more stable)
one will have:
1. fewer (or smaller) formal charges.
2. the negative charge on the more electronegative atom, and vice versa.
COCl2
NCO-
O
O
Cl C Cl
Cl C Cl
better
N C O
or
O
Cl C Cl
N C O
2
N C O
best
worst
Exceptions to the octet rule
1. odd number of electrons
ClO
NO
Cl O
N O
2. less than an octet: Groups 1, 2, and 3
AlCl3
Cl Al Cl
Cl Al Cl
Cl
Cl
better
very bad
3. more than an octet: 3rd row and below can “expand” their octets
Cl
PCl5 Cl
P Cl
Cl Cl
SO42-
O
O
2
O S O vs O S O
O
O