Naming Molecules
Ch. 9, Section 2: pg. 248
Naming Binary Molecular
Compounds
1. The first element is always named
first, using the entire element name.
2. The second element is named using
the root of the element and adding
“-ide”. (Just like ionic cpds.)
3. Use prefixes on the names to
indicate the number of atoms of each
type that are present in the
compound.
Common Prefixes in Covalent
Compounds
# of atoms
Prefix
# of atoms
Prefix
1
mono-
6
hexa-
2
di-
7
hepta-
3
tri-
8
octa-
4
tetra-
9
nona-
5
penta-
10
deca-
What is the name of P2O5?
Step 1: phosphorous
Step 2: oxide
Step 3: diphosphorous pentoxide
EASY!!!!!!
What are the names of the
following binary molecular
compounds?
1. CCl4
2. As2O3
3. CO
4. SO2
5. NF3
Answers:
1. carbon tetrachloride
2. diarsenic trioxide
3. carbon monoxide
4. sulfur dioxide
5. nitrogen trifluoride
Naming Acids
There are two common types of acids:
*binary acids
*oxyacids
Binary acids contain HYDROGEN
and one other element.
Steps to naming binary acids:
1. Use the prefix hydro- to name the
hydrogen part of the compound.
2. Use a form of the root of the second
element plus the suffix –ic.
3. End with the word acid.
Example: HCl
hydrochloric
acid
Name these binary acids.
1. HI
2. HF
3. H2S
4. H2Se
5. HBr
Answers
1. hydroiodic acid
2. hydrofluoric acid
3. hydrosulfuric acid
4. hydroselenic acid
5. hydrobromic acid
Oxyacids are acids that contain
hydrogen and an oxyanion.
**Remember: An oxyanion is a
polyatomic ion that contains oxygen.
EX: PO43SO42-
Steps to naming oxyacids:
1. Identify the oxyanion present in the
acid. (You can use your ion chart.)
2. Use a form of the anion and
**if the oxyanion ends in “-ate”,
change to “-ic”.
**if the oxyanion ends in “-ite”,
change to “-ous”.
3. Add the word “acid” at the end.
(We do NOT use “hydro” with
oxyacids.)
“-ate”
“-ite”
“-ic”
“-ous”
Example: HClO3
Step 1: ClO3 is the chlorate anion.
Step 2: Change chlorate to chloric.
Step 3: Add the word “acid”.
ANSWER:
chloric acid
Name the following oxyacids
1. HClO2
2. H2SO4
3. H2SO3
4. H3PO4
5. HNO3
Answers
1. chlorous acid
2. sulfuric acid
3. sulfurous acid
4. phosphoric acid
5. nitric acid
ASSIGNMENT
**Page 250: #18 – 22
**Page 251: #23 – 29
Lewis Structures
Lewis structure: when electron-dot
diagrams are used to show how electrons
are arranged in molecules
EX:
We use a line to show bonding pairs of
electrons in a Lewis structure and a pair
of dots to show electrons that are NOT
being shared (called “lone pairs”).
Drawing Lewis Structures
1. Count ALL the valence electrons for the
molecule.
EX: CCl4 C = 4 e-, Cl = 4x7 = 28 eTotal valence e- = 32 e2. Determine the central atom.
*H is NEVER a central atom.
*The halides are NEVER a central atom.
*The element with the lowest electronegativity
is the central atom.
*The only single element would be the central
atom
EX: C is the only single element, so it would be
the central atom in the Lewis structure.
3. Place two electrons in each bond by
drawing a line to represent the bond.
Cl
l
Cl – C – Cl
l
Cl
4. Complete the octet of the atoms
attached to the central atom by
adding electrons in pairs.
(See board for example.)
5. Place any remaining electrons on the
central atom in pairs.
*****REMEMBER!
The total electrons
in the Lewis structure MUST equal the
number of electrons in step #1.******
(See board for example.)
6. If the central atom does not have an octet
after step #5, form double or triple bonds,
as needed, between the central atom and
one or more of the terminal atoms.
(See board for example.)
Draw the Lewis Structures for the
following molecules
1.
2.
3.
4.
5.
PH3
H2S
HCl
CCl4
SiH4
Multiple Bonds
 Many molecules attain a noble-gas
configuration by sharing more than
one pair of electrons between two
atoms, forming a multiple covalent
bond.
 C, N, O, and S most often form
multiple bonds.
Lewis Structures
O=C=O
Which of these is a
**single bond?
**triple bond?
**double bond?
Multiple Bonds: Draw the Structural
Formulas for the following:
1. O2
2. N2
3. CO2
Lewis Structures for Polyatomic
Ions
1. Determine the number of valence
electrons in the atoms present
EX: PO4
1 P = 5 electrons
=5
4 O = 6x4 electrons = 24
29 electrons
2. Draw the Lewis structure for the ion.
3. Count the total electrons in the Lewis
structure.
4. For negative ions: Subtract ions in
the Lewis structure from the valence
electrons found in step 1 and this is
the charge.
For positive ions: Subtract the
valence electrons in step 1 from the
total ions in the Lewis structure
See next slide for examples……………………
1. ClO4-1
2. NH4+1
Resonance Structures
 If a molecule or polyatomic ion has
BOTH a double bond AND a single
bond, it is possible to have more than
one correct Lewis structure.
 RESONANCE is a condition that
occurs when more than one valid
Lewis structure can be written for a
molecule or ion.
 Resonance structures differ only in
the position of the electron pairs.
Examples of Resonance Structures
1. BCl3
2. SO2
3. SO3
See board for answers.
Exceptions to the Octet Rule
1. Molecules with an odd # of valence
electrons
EX: NO2
ClO2
NO
2. Compounds with fewer than 8
valence electrons present (this is
very rare)
EX: BH3
3. Compounds in which the central
atom has more than 8 valence
electrons -- called an EXPANDED
OCTET ; occurs in energy levels of
elements in period 3 and up
**Extra lone pairs are added to the
central atom OR more than 4
bonding atoms are present.
EX: IF4
PCl5
VSEPR Model : Valence Shell
Electron Pair Repulsion model
 Based on an arrangement that
minimizes the repulsion of shared and
unshared pairs of electrons around
the central atom
 READ pages 259 – 261.
 Look at Table 9-3 as you read.
 Complete #49 – 53 on pg. 262.
Electronegativity and Polarity
 Electronegativity indicates the
relative ability of an atom to attract
electrons in a chemical bond.
 It generally increases as the atomic
number increases ACROSS A
PERIOD and generally decreases as
you go DOWN A GROUP.
Polar or Nonpolar???????????
 Identical atoms, like N2, have an
electronegativity difference of zero, and
the electrons in the bond are equally
shared between the two atoms. This is a
NONPOLAR COVALENT BOND, or a
PURE COVALENT BOND.
 A covalent bond between atoms of
different elements does not have equal
sharing of the electron pair, due to the
difference in electronegativity.
 Unequal sharing results in a POLAR
COVALENT BOND. The shared electrons
are pulled toward one of the atoms and
spend more time around that atom than
the other atom. Partial charges occur at
the ends of the bond. This bond is often
referred to as a dipole (two poles).
 To determine if the bond is polar or
nonpolar, you must look at the shape of
the molecule. Draw the molecular
structure, using what you know from
Table 9-3.
 SYMMETRIC MOLECULES ARE USUALLY
NONPOLAR AND ASYMMETRIC ARE POLAR
AS LONG AS THE BOND TYPE IS POLAR.
Do #60 – 63 on page 266.
CLASSWORK: Pg. 273
Complete #94, 95, 96, 97, 98,
99: a,c,d,
100: a,d,e