Chapter 6
Manipulating Polyatomic Ions and
Chemical Bonding
Basic Polyatomics
Name
Formula
Name
Formula
Hydroxide
OH-1
Nitrate
NO3-1
Ammonium
NH4+1
Phosphate
PO4-3
Acetate
C2H3O2-1
Chromate
CrO4-2
Carbonate
CO3-2
Dichromate Cr2O7-2
Sulfate
SO4-2
Chlorate
ClO3-1
Ways to expand your polyatomics

Polyatomic ions vary in their charges, number of
oxygen atoms, and number of hydrogen atoms.

1. To change the number of oxygens:
One more oxygen
 Memorized
 One less oxygen
 Two less oxygens

ClO4-1
ClO3-1
ClO2-1
ClO-1
perchlorate
chlorate
chlorite
hypochlorite
Ways to expand your polyatomics

2. Family Members

Whatever is true for chlorine, is also
true for fluorine, bromine, and iodine.
ClO3-1

Memorized

F substitution FO3-1
fluorate

Br substitution BrO3-1
bromate

I substitution
IO3-1
chlorate
iodate
Ways to expand your polyatomics

3. If you add a hydrogen, you have to make
the ion more positive (one less negative)
and call the ion “bi________”

Memorized
CO3-2 - carbonate
HCO3-1 – bicarbonate

Memorized
SO4-2 - sulfate
HSO4-1 - bisulfate
Ways to expand your polyatomics

4. Combinations of #1, #2, and #3 are
possible:
HSO3-1 is called bisulfite
 FO2-1 is called fluorite

Rules for expanding your list of
polyatomic ions
Rule #1
To Change the number of oxygens:

Remove one oxygen = change ending of name to –ite

Remove two oxygens = change ending of name to –ite and
beginning of name to Hypo-

Add one oxygen = change beginning of name to Per-
Rules for expanding your list of
polyatomic ions
Examples:
Name
Formula
Name
Formula
Name
Formula
Name
Formula
ClO4-1
SO5-2
NO4-1
PO5-3
ClO3-1
SO4-2
NO3-1
PO4-3
ClO2-1
SO3-2
NO2-1
PO3-3
ClO-1
SO2-2
NO-1
PO2-3
Rules for expanding your list of
polyatomic ions
Rule #2 Other Family Members
 Elements near each other in the same column tend to
form similar polyatomic ions.
Name
Formula
Name
Formula
Name
Formula
Chlorate
ClO3-1
Sulfate
SO4-2
Phosphate
PO4-3
Fluorate
FO3-1
Selenate
SeO4-2
Arsenate
AsO4-3
Iodate
IO3-1
Bromate
BrO3-1
Rules for expanding your list of
polyatomic ions
Rule #3

Add a hydrogen
Add only one H = change the beginning of the name to
bi- and make the charge one less negative (due to
hydrogen’s positive one charge)
Name
Formula
Name
Carbonate
Sulfate
Bicarbonate
Bisulfate
Formula
Rules for expanding your list of
polyatomic ions
Rule #4 Combinations of 1, 2, & 3
 Combinations of #1, #2, and #3 are possible:









HSO3-1
Memorized
Lose an “O”
Add an “H”
SO4-2
SO3-2
HSO3-1
Sulfate
Sulfite
Bisulfite
HFO2
Memorized
Substitute an F
Lose an “O”
Add an “H”
ClO3-1
FO3-1
FO2-1
HFO2
Chlorate
Fluorate
Fluorite
Bifluorite
Rules for expanding your list of
polyatomic ions

Combos Cont’d

Ex1: What is the formula for hypoiodite?

Find I on the P-table (near Cl). Chlorine forms
chlorate (ClO3-1). Thus, Iodine forms iodate
(IO3-1). The –ite and hypo- in hypoiodite mean
that iodate lost two oxygens.

Hypoiodite = IO-1
Rules for expanding your list of
polyatomic ions

Combos Cont’d

Ex2: What is the formula for Biperselenate?

Find Se on the periodic table. It is near S. Sulfur
forms sulfate (SO4-2). Therefore, selenium forms
selenate (SeO4-2). The per- in biperselenate
means that selenate has gained one oxygen.
Also, the bi- means that it has gained a hydrogen
(don’t forget to change the charge!).

Biperselenate
= HSeO5-1
Monatomic Ions

For nonmetals, almost all single names that end with
–ide indicates a single charged atom.

Simply write the symbol and the charge. The periodic
table column indirectly indicates the element’s charge.
Remember, elements want to have 8 electrons in their
outer shell (Octet Rule).
Monatomic Ions






Column #1 elements have a +1 charge
Column #2 elements have a +2 charge
Column #3 = +3
Column #15 = -3
Column #16 = -2
Column #17 = -1
Monatomic Ions

Ex1: What is the formula for chloride?
Cl-1

Ex2: What is the formula for an aluminum ion?
Al+3

Ex3: What is the name of the S-2 anion?
Sulfide

Ex4: What is the name of the Mg+2 cation?
Magnesium Ion
6.1 Introduction to chemical bonding

Most elements are not found alone in nature. They
are “stuck” to other atoms.

Chemical Bond - Link between atoms that results
from the mutual attraction of their nuclei for their
electrons.

Types of chemical bonds:



Ionic - transfer of electrons (metal + nonmetal)
Covalent - sharing of electrons (2 nonmetals)
Metallic - happens in metals when there is only
one type of element
Introduction to Chemical Bonding


Covalent bonds may be polar or nonpolar

Polar - unequal sharing of electrons (HCl)

Nonpolar - equal sharing of electrons (H2)
There are two ways to predict polar vs. nonpolar ( and
covalent vs. ionic)
Introduction to Chemical Bonding

#1 Use electronegativity difference
 0 = nonpolar covalent
 0.4 - 1.7 = polar covalent
 greater than 1.7 = ionic

Examples

NaCl

Cl2
Cl = 3.16
Na= - 0.93
2.23
Ionic
HCl
Cl = 3.16
Cl = - 3.16
0 Nonpolar Covalent
Cl = 3.16
H = - 2.20
.96
Polar Covalent
Introduction to chemical bonding

#2 - There is an easier way to predict
 Ionic = metal + nonmetal

Polar Covalent = 2 different nonmetals

Nonpolar Covalent = 2 of the same nonmetals
Ionic Bonds

Ionic compound - a substance composed of
positive and neg. ions so that the charges are
equal. It involves a transfer of electrons.
 Ca+2 with Cl–1 will form the compound CaCl2.

It takes two chlorine ions to cancel out the the +2
charge on the calcium ion.

Formula unit - lowest whole # ratio of ions
 Ionic Bond = a METAL + a NONMETAL


Metals - lose e- - why? low IE
NM - gain electrons - why? high electronegativity
Ionic Bonds

Metals lose electrons until they become like a
noble gas . (8 valence e-)

Nonmetals gain e- until they do the same.

Both go to s2p6 - 8 valence e- - called a stable octet

The tendency to arrange e- so each atom has 8 is
called the octet rule or rule of 8
Ionic Bonds

The formation of an ionic bond:
Na to Cl = [Na]+1[Cl]-1
 Na

1s 2s

2p
3s
Cl
1s 2s 2p 3s 3p
Ionic Bonds

Ionic bonding picture:

Ex1: Na to Cl=
Na

[Na]+1[ Cl ]-1
Cl
Ex2: Ba to Cl = [Ba]+2 2[ Cl ]-1
Ba
Cl
Cl
Ionic Bonds

The easy way:
Find the charge of each atom
 “criss cross” the charges – charge cancels
out and you are left with a neutral
compound
 EX3: Al
N
 EX4: Na S
 EX5: Al
S

Ionic Bonds

A few more examples not in your note
packet using Polyatomic Ions

USE PARENTHESIS
Li and NO3-1
 Ca and C2H3O2-1
 Magnesium Phosphite
 Aluminum hyponitrite

Ionic Bonds


Energy is involved in all chemical reactions.

Na + Cl
yields
NaCl
+ 769 kJ

Lattice energy - energy released when an ionic compound
forms.

NaCl = - 769 kJ/mole
KCl = -718 kJ/mole
NaF = - 922 kJ/mole
smaller ions have higher lattice energies
Ionic Bonds

Properties of ionic compounds:
Hard
 Shatter
 Conduct electricity
 High melting point


Odorless
6.2 Covalent Bonding

In covalent bonding atoms share electrons. In the H2
molecule, each H atom says, "I only need one more eto be like a noble gas (helium)." Since each hydrogen
has only one electron, when two hydrogens bond they
can share their electrons.
Covalent Bonding

Molecule - smallest quantity of matter that exists by
itself and retains the properties of that substance.
Describes a covalently bonded substance.

monatomic molecules - He, Ne, Ar, (noble gases are
always monatomic)

diatomic molecules – H2 O2 N2 Cl2 Br2 I2 F2 (you must
memorize these!!)

polyatomic molecules - P4, S8, C6H12O6
Covalent Bonding

The formation of a covalent bond
=

Bond Length vs. Bond Energy

 Bond length =  Bond Energy
Covalent Bonding


Diatomic Molecules and Orbital Notation (Orbital
overlap or notation diagrams):
H2
O2
H 1s
O
1s 2s
2p
H
1s
O
N2
N
1s
2s
2p
N
1s
2s
2p
1s
2s
2p
Covalent Bonding

Why are these atoms forming bonds?
 Octet Rule- Atoms lose, gain, or share electrons to
have 8 electrons in their outer shell.

HF – orbital notation
H
F
1s
1s
2s
2p
Lewis Dot Diagrams of molecules (covalent
compounds) and polyatomic ions

Basic rules







Each atom wants 8 electrons (except H wants 2).
Each atom goes for close to the right # of bonds.
The least electronegative atoms goes in the middle
OR
The atom that makes the most bonds goes in the
middle. (H always on the outside.) OR
The “single guy” (the atom that does not have a
subscript after it) goes in the middle.
Symmetry is key!!!
Place the atoms in order (left, right, bottom, and
top) around a central atom
Lewis Dot Diagrams

To determine the number of bonds:



S=N–A
2
S = shared electrons (bonds)
N = needed e- (all elements need 8 except
for H which needs 2)
A = how many e- an atom actually has (#
of valence e-)
If using S=N-A, add the charge into the Actual amount
of electrons.
Put [ ] around the molecule and include the charge
Lewis Dot Diagrams

Examples: Draw the following Lewis structures

EX1:
CH4

Ex2:
H 2O

Ex3:
PCl3
Lewis Dot Diagrams

Ex4:
SiH2F2

Ex5:
CS2

Ex6:
C2H6
Lewis Dot Diagrams

Ex7:
C2H4

Ex8:
C2H2

Ex9:
CH2O
Lewis Dot Diagrams

Ex10:
HCN

Ex11:
FON
Drawing polyatomic ions

count electrons: if the charge is - 3, add 3 electrons to A
 EX: PO4-3






less bonds than atoms want = negative charge
more bonds than atoms want = positive charge
P wants 3 bonds, has 4: + 1 charge
Each O wants 2, has 1 so each O = -1
Total = - 3
Ex11:
PO4-3
Coordinate covalent bond

Coordinate covalent bond- 2 shared electrons
in a bond are donated by 1 atom

Examples:









NH4+
OH-1
sulfate
nitrate
nitrite
carbonate
bicarbonate
H2SO4
H3PO4
6.4 Metallic Bonding - “Sea of electrons
theory”

The nuclei are arranged in a systematic lattice.

The bond strength relies on the nuclear charge and
the number of valence e

Ex. Mg is stronger than Na
The valence electrons form a sea of free moving
electrons that are attracted to multiple positive nuclei.
Metallic Bonding

Conducts Electricity as a result of free electrons.

Malleability and ductility results from the nuclei's ability
to move passed each other
Metallic Bonding

Remember:



in ionic bonds some atoms want e- and some don’t
in covalent bonds, all atoms share –
in metals, no one atom wants the e-
6.5 The Properties of Molecular
Compounds

Valence shell electron pair repulsion theory
(VESPER) – e- pairs get as far away from each
other as possible

Because of this we can predict the shape of molecules
based on how many bonds and lone pairs are on the
central atom
Shapes
Example
Shared
Pairs
AB
1
AB2
2
Lone
Pairs
Shape
Example
Angle(s) and drawing
0
Linear
HCl
180o
0
Linear
CO2
180o
Shapes
Lone
Pairs
Shape
Example
Example
Shared
Pairs
Drawing
AB2E
2
1
Bent
SO2
119.5o
AB2E2
2
2
Bent
H2O
H2S
104.5o
AB3
3
0
Triangular
Planar
CH2O
120o
Shapes
Example
Shared
Pairs
Lone
Pairs
Shape
Example
Drawing
AB3E
3
1
Triangular
pyramidal
NH3
107.5o
AB4
4
0
Tetrahedron
CH4
109.5o
Shapes
Example
Shared
Pairs
Lone
Pairs
Shape
Example
Drawing
AB5
5
0
Triangular
Bipyramidal
PCl5
90o
1200
1800
AB6
6
0
Octahedron
SF6
900
1800
Shapes









Examples: Predict the shapes of the
following (show all work):
Ex1:
CCl4
Ex2:
HBr
Ex3:
SO3
Ex4:
SO2
Ex5:
H2S
Ex6:
NH3
Ex7:
ClO4 -1
Ex8:
PF5
Intermolecular Forces

Intermolecular forces (IMF)- forces that hold
molecules together



happens between covalent compounds
intermolecular forces - can be weak or strong
Intramolecular forces – chemical bonds (ionic,
covalent, metallic)


happens within a molecule or compound
always strong
H -------------Cl
H -------------Cl
Intermolecular Forces

Types of intermolecular forces

dipole-dipole

dipole - when electrons are unevenly distributed

Ex1: predict the IMF that occurs with HCl

Ex2: predict the IMF that occurs with H2O (***one of the
most impt. Ever!)
Intermolecular Forces
hydrogen bonding - H-bonding is a “super-duper”
dipole-dipole


H-bonding happens any time H is bonded to F, O, or N

Hydrogen bond is FON!!!
Why does this happen?



A large difference in electronegativity between F, O, or
N and H results in one
end of the molecule being very negative, while the other
end is very positive.
Intermolecular Forces

Effect of H-bonds on physical properties:

H-bonding tends to cause the following in
substances:






 Boiling Point
 Heat of Vaporization
 Vapor Pressure
 Melting Point
H-bonds causes water to expand when it freezes.
H-bonding is also responsible for the shapes of
proteins.
Intermolecular Forces

Recap:






Polar molecules have dipole-dipole IMF holding
them together.
H-bonding is “super” dipole-dipole IMF
These two types of IMF usually result in substances
being solids or liquids at room temp.
Most nonpolar covalent substances are gases at
room temp. as the forces holding them together are
not strong enough to keep the molecules attracted hence they are gases
O2, H2, N2 - straight nonpolar substances
CO2 - have dipoles, but nonpolar due to its
molecular geometry
Intermolecular Forces

Van der Waals Forces (London Forces)



Temporarily induced dipoles caused by the motion of
electrons.
more electrons = more attraction so, bigger atoms have
stronger Van der Waals forces
Occur with noble gases
Bonding

Bond Energy -what is the strength of chemical bonds?
 bond energy - energy needed to break a bond measured in kJ/mole
 bond strength and stability:



stronger bond - more stable -needs more energy to break
the bond
weaker bond - takes little energy to break the bond so the
chemical is unstable
chemical changes favor lower energy states exothermic reactions are favored
Bonding

Bond Strength
which is stronger? - single, double, or triple
bond? Triple
 which is shortest bond length? s, d, or t?
triple
 which is stronger, short or long bonds?
short
