Chemical Bonding
and Nomenclature
By Paul Surko
New Dimensions High School
Poinciana, FL
s
8
Bonding,
the to
way
atoms
attracted
to
I want you
meet
a are
friend
of mine?
each other to form molecules, determines
nearly all of the chemical properties we see.
And, as we shall see, the number “8” is very
important to chemical bonding.
5.1 What are Molecules?
Molecules are a
combination of atoms
bonded together. Bonding
determines the chemical
properties of the molecule
(compound).
5.5 Ionic Bonding-Being Like
the Noble Gases
All atoms want to have the same number of electrons as the
Noble Gases. The Noble Gases have very stable electron
configurations. In order to achieve the same electron
configuration as the Noble Gases metal atoms will give up
electrons to form positive ions (cations) and non-metal atoms
will receive or take additional electrons to become negative
ions (anions). IONS are charged particles.
Na becomes Na+
Mg becomes Mg+2 Al becomes Al+3
Cl becomes Cl-
O becomes O-2
N becomes N-3
The positive and negative ions are attracted to each other
electrostatically.
Opposites Attract!
Putting Ions Together
Na+ + Cl- = NaCl
Ca+2 + Cl- = CaCl2
Ca+2 + O-2= CaO
Na+ + O-2 = Na2O
Al+3 + S-2 = Al2S3
Ca+2 + N-3 = Ca3N2
You try these!
Li+ + Br- = LiBr
Mg+2 + F- =
Al+3 + I- =
NH4+ + PO4-3 = (NH4)3PO4
Not NH43PO4
+
K + Cl = KCl
AlI3
Sr+2 + P-3 = Sr3P2
MgF2
5.2 The Covalent Bond
Atoms can form molecules by sharing
electrons in the covalent bond. This is done
only among non-metal atoms.
5.3 Dot Structures-Octet Rule
(All atoms want 8 electrons around them.)
Valence electrons are those in the outermost orbitals.
They are the ones that can form bonds.
Lewis came up with a way to draw valence electrons so
that the bonding could be determined.
Rules to Write Dot Structures
Write a skeleton molecule with the lone atom in the middle
(Hydrogen can never be in the middle)
2. Find the number of electrons needed (N)
(8 x number of atoms, 2 x number of H atoms)
3. Find the number of electrons you have (valence e-'s) (H)
4. Subtract to find the number of bonding electrons (N-H=B)
5. Subtract again to find the number of non-bonding
electrons (H-B=NB)
6. Insert minimum number of bonding electrons in the
skeleton between atoms only. Add more bonding if needed
until you have B bonding electrons.
7. Insert needed non-bonding electrons around (not
between) atoms so that all atoms have 8 electrons around
them. The total should be the same as NB in 5 above.
1.
Let's Try it!
1.S
2.N
3.H
4.B
5.NB
6.E
H O H
Water H2O
2 x 2 = 4 for Hydrogen
1 x 8 = 8 for Oxygen
4+8=12 needed electrons
2 x 1 = 2 for Hydrogen
1 x 6 = 6 for Oxygen
You have 8 available electrons
12 - 8 = 4 bonding electrons
8 – 4 = 4 non-bonding electrons
..
H:O:H
●●
12 N
- 8H
- 4B
4 NB
H:O:H
..
H:O:H
●●
1.S
2.N
3.H
4.B
5.NB
6.E
H
HNH
Let's Try it!
Ammonia NH3
3 x 2 = 6 for Hydrogen
1 x 8 = 8 for Nitrogen
6+8=14 needed electrons
3 x 1 = 3 for Hydrogen
1 x 5 = 5 for Nitrogen
You have 8 available electrons
14 - 8 = 6 bonding electrons
8 – 6 = 2 non-bonding electrons
H
..
H:N:H
●●
14 N
- 8H
- 6B
2 NB
H
..
H:N:H
H
..
H:N:H
●●
Let's Try it!
1.S
2.N
3.H
OCO
Carbon Dioxide CO2
1 x 8 = 8 for Carbon
2 x 8 = 16 for Oxygen
8+16=24 needed electrons
1 x 4 = 4 for Carbon
2 x 6 = 12 for Oxygen
You have 16 available electrons
4.B
24 - 16 = 8 bonding electrons
5.NB
16 – 8 = 8 non-bonding electrons
6.E
..
..
O::C::O
●●
●●
24 N
- 16 H
- 8B
8 NB
O::C::O
..
..
O::C::O
●●
●●
1.S
2.N
3.H
Let's Try it!
O
OCO
Carbonate CO3-2
3 x 8 = 24 for Oxygen
1 x 8 = 8 for Carbon
24+8=32 needed electrons
3 x 6 = 18 for Oxygen
1 x 4= 4 for Carbon
You have 22 + 2 more available e-'s
4.B
32 - 24 = 8 bonding electrons
5.NB
24 – 8 = 16 non-bonding electrons
6.E
..
.. :O:
.. ..
O::C:
O:
●●
●●
-2
32 N
- 24 H
- 8B
16 NB
O
..
O::C:O
..
.. :O:
.. ..
O::C: O:
●●
●●
5.6 Polarity-Unequal Sharing
of Electrons
Even though all atoms want the same number of electrons as the Noble
Gases, some want to get or give them more than others. The magnitude
of this attraction for electrons is called “Electronegativity”. The more
electronegative an atom is, the more it wants the electrons.
Some atoms
want to gain
electrons so
bad, they
take them
altogether to
form negative
ions. Some
want to lose
them so bad
that they
become
positive ions.
Examples of Polar and NonPolar
Compounds
HCl The Chlorine wants the electrons more than the
Hydrogen. Thus we have +δHCl-δ.
NaCl Since Na is a metal it gives up its electron to form Na+
and Cl takes the electron completely to form Cl-.
Cl2 (Cl—Cl) The Chlorine molecules want the electrons
equally so they form a non-polar molecule with NO partial or
full charges.
H2O Water is a bent molecule. The lone pair of electrons
from the Lewis structure distorts its shape and it becomes a
very polar molecule.
..
..
..
:O:H
O::C::O
●●
●●
●●
H
CO2 Carbon Dioxide is a linear molecule. It has no lone pairs
of electrons from the Lewis structure. The two oxygen atoms
pull equally and make it a non-polar molecule.
5.7 Nomenclature
Naming of Compounds
Binary Compounds have two types of atoms (not diatomic
which has only two atoms).
Metals (Groups I, II, and III) and Non-Metals
Metal _________
Sodium
+ Non-Metal _________ide
Chlorine
Sodium Chloride NaCl
Metals (Transition Metals) and Non-Metals
Metal ______
Iron +Roman Numeral (__)
III + Non-Metal ________ide
Bromine
Iron (III) Bromide FeBr3
Compare with Iron (II) Bromide FeBr2
5.7 Nomenclature
Naming of Compounds
Binary Compounds have two types of atoms (not diatomic
which has only two atoms).
Metals (Transition Metals) and Non-Metals
Older System
Metal (Latin) _______
Ferrous + ous or ic + Non-Metal ________ide
Bromine
Ferrous Bromide FeBr2
Compare with Ferric Bromide FeBr3
Non-Metals and Non-Metals
Use Prefixes such as mono, di, tri, tetra, penta, hexa, hepta, etc.
CO2 Carbon dioxide CO Carbon monoxide
PCl3 Phosphorus trichloride CCl4 Carbon tetrachloride
N2O5 Dinitrogen pentoxide
CS2 Carbon disulfide
Let’s Practice!
Name the following.
CaF2
K2S
CoI2
SnF2
SnF4
OF2
CuI2
CuI
SO2
SrS
LiBr
Calcium Flouride
Potassium Sulfide
Cobalt (II) Iodide or Cobaltous Iodide
Tin (II) Flouride or Stannous Flouride
Tin (IV) Flouride or Stannic Flouride
Oxygen diflouride
Copper (II) Iodide or Cupric Iodide
Copper (I) Iodide or Cuprous Iodide
Sulfur dioxide
Strontium Sulfide
Lithium Bromide
Polyatomic Ions
(partial list from page 195 (193 2nd edition))
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Ammonium……………...
Nitrate……………………
Permanganate…………. .
Chlorate…………………
Hydroxide……………….
Cyanide………………….
Sulfate…………………...
Carbonate……………….
Chromate………………..
Acetate…………………..
Phosphate……………….
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NH4+
NO3MnO4ClO3OHCNSO4 2 CO32CrO42C 2 H 3 O2 PO43-
Acids (with H in front)
Binary acids (without oxygen in formula)
Hydro _________ ic Acid
HCl Hydrochloric acid HBr Hydrobromic acid
Oxy acids (with oxygen in formula)
-ate goes to –ic and –ite goes to -ous
HNO3 Nitric acid HNO2 Nitrous acid
H2SO4 Sulfuric acid H2SO3 Sulfurous acid
H3PO4 Phosphoric acid H3PO3 Phosphorous acid
Lets Practice!
HF
Na2CO3
Hydroflouric acid
H2CO3
KMnO4
HClO4
H2S
NaOH
CuSO4
PbCrO4
Sodium carbonate
Carbonic acid
Potassium permanganate
Perchloric acid
Hyrdogen sulfuric acid
Sodium hydroxide
Copper (II) sulfate or Cupric sulfate
Lead (II) chromate or Plubous chromate
H2O
Hydrooxic acid (no……just water)
NH3
Nitrogen trihydride (no..just ammonia)