Chemical Bonding
When two or more atoms come together to form a molecule, they are held together by
chemical bonds (electrical charges).
There are three types of chemical bonds.
1. Ionic bonds
2. Covalent bonds
3. Polar covalent bonds
1. Reactive Metals:

Have low ionization energies: Therefore a small amount of energy is required to
cause them to loose an electron.
2. Reactive Non-metals:

Have large electron affinities: Therefore, they tend to gain electrons readily.
When an active metal combines with an active non-metal, electrons are readily
transferred from metal to non-metal.
Metal+
+
Looses electron
become a
Positive ion
Non-Metal-
→
Ionic
Compound
gains electron Neutral Compound
to become a
negative ion
Ionic Bonding
It is a chemical bond that is the result of the attractions between the oppositively charged
metal and non-metal ions.
Therefore bonds/compounds are formed via the transfer of electrons between metals and
non-metals.
Ex.
Na+ +
Cl----- NaCl
Na + Energy ---- Na+ + electron
Cl
+ Electron --- Cl- + energy
Properties/Structures of Ionic Compounds
All ionic compounds have similar properties
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1. Solids at SATP
2. They have high melting point
3. All are electrolytes, i.e. they all conduct electricity.
These properties are the result of:
1. The type of bonds that are formed between the metals and non-metals (i.e.
Anions and Cations)
2. Ionic compounds are locked in a regular structure, held by the balance of
attractive bond and electrical repulsion
3. The structure(s) that result is a crystal lattice
Ionic Structure
The anions and cations in an ionic compound are locked together in a structure know as
crystal lattice.
Crystal Lattice:
 A 3-D pattern created by regularly repeating cations and anions.
 It looks like the scaffolding used when workers are repairing bridges (series of
struts joined to one another in a repeating pattern.
All lattices are arranged so that each ion has the greatest possible number of oppositely
charged ions close while keeping ions with the same charge as far away as possible. In
theory, this arrangement creates strong attractions.
THE FORMATION OF IONIC COMPOUNDS








Most simple compounds containing metallic elements are classified as ionic
Elements in same chemical group tend to participate in the same chemical
reaction producing ionic compound with same general formula
Examples: Elements in groups 1 and 2 form ionic compounds with Oxygen called
Oxides
Oxides formed by elements in group 1 has a general formula M2O, while those of
group 2 form oxides with formula MO; M represents the metal ion
In the same way metals in groups 13 and 15 except mercury, will form ionic oxide
when burned in air
Groups 1 metals also react readily with group 17 elements to form ionic
compounds with general formula MX called ionic halides
while group 2 metals form ionic compound with halogens with a general formula
MX2
In general, the addition of metals from groups 1 and 2 to water produce hydrogen
gas and a basic ionic compound
1. Formation of Ionic Bonds
Bonds formed between cations and anions
2
Cations-----give up é
Anions------ accept é
In order to become isoelectric to a noble gas
OR to have the maximum number of valence
electrons
Ex.
Sodium
+1
Na
Na
+
+
+
Fluorine
-1
F ----- NaF
+
F ----- Na+
+ F-
Isoelectric to
Na+
Ne
F
Ne
Try these:
a) Calcium and fluorine
b) Potassium and oxygen
Representing Ionic Bonds:
Lewis Symbols (or Electron Dot diagram)
The electron dot diagrams are a simple method used to represent atoms. In this notation,
the symbol represents the nucleus and all the electrons in the inner shells. Dots are then
placed around the symbol showing the number of electrons in the valence shell.
e.g. Chlorine ------- 17 Electrons (10 inner electrons and 7 valence
electrons)
Cl
3
PREDICTING COMMON IONS OF ATOMS




Noble gases are stable and inert because they have eight valence electrons
Other elements tend to form ions in order to attain this eight valence electron to
have special stability
This arrangement is called Stable Octet
To attain this, Groups 1, 2 and 3 elements will lose electrons to form Cations;
while elements in groups 15, 16, and 17 will gain electrons to form Anions
IONS AND THE HUMAN BODY
1. Five metals: Calcium, potassium, sodium, magnesium and iron, which
form positive ions in solution are essential for maintaining good health in
the body
2. Mg2+, Na+, and K+ are major component of blood plasma
3. Ca2+ is important in bone formation
4. Negative ions essential to life include Cl- ( an important component of
blood), I- ( Prevents goitre)
Covalent Bonding
A bond formed by 2 or more non-metal atoms sharing one or more pairs of electrons
equally.
Many non-metals exist as covalently bonded diatomic molecules
Most compounds do not contain ions. Instead, they contain neutral groups of atoms called
molecules.
Water is a molecule made up of the non-metals, hydrogen and oxygen.
Non-metals form stable arrangements by sharing a pair of electrons.
Diatomic Molecules: A molecule made of two identical atoms.
Common Diatomic Molecules
Oxygen gas
O2
Hydrogen gas
H2
Nitrogen gas
N2
Fluorine gas
F2
Chlorine gas
Cl2
Polyatomic molecules: This is when the molecule contains more than two atoms, such as
ammonia, NH3
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Formation of Covalent Bonds
Bonds formed when non-metals combine. These bonds are the result of the sharing of one
or more pairs of é between two atoms.
Ex.
H
+
H --------- H2 (g)
H
+
H ------ H H
Atoms that bond in this manner usually follow the octet/duel rule.
Octet/Duel Rule: A rule that states when atoms combine to from covalent molecules, the
bonds is formed in such a way that each atom achieves 8 valence electrons (2) in the case
of Hydrogen).
Ex. Carbon Dioxide (CO2)
O
C
O
------
O
O
C
C
O
O
Lone Pair
Lone Pair: A pair of valence electrons not involved in bonding.
Try nitrogen gas N2)
The bond distance is
 Shorter in a triple bond compared to a double bond
 Shorter in a double bond compared to a single bond
Rules for drawing Lewis structure for Molecular Compounds
Rule1: Identify the central atom (the element with the highest bonding capacity). The
central atom is the one that is the least electronegative. Arrange the other atoms
symmetrically around the central atom.
O
S
O
Rule 2: Determine the # of bonds formed in the molecule.
Compare the total # of valence é in the molecule with the # of é each atom would
acquire for a full valence shell.
SO2
#é
:
6+2(6) = 18 é
5
Full Octet
8+2(8) = 24 é
6é
(3 pairs of é. 3 bonds)
If the structure is a polyatomic ion
 Add 1 é for each unit of negative charge (anions gain é)
 Subtract 1 é for each unit of positive charge (cations loose é)
Rule 3: Place the number of é to be shared between the atoms, a pair at a time. Use as
many pairs that remain to make double or triple bonds.
O
S
O
(bonds, 6 é)
Rule 4: Add remainder of the available é to complete the octet/duets of all atoms. These
should be just enough if the molecule or ion follows the octet rule.
O
S
O
Examples:
a) Dot structures for single bonds:
H2, H2O, NH3, BF3, N2F5 , O22b) Dot structure for double bonds
CO2, SO2, O2 CS2
c) Dot structure for triple bonds
N2, HCN
Coordinate covalent bond
Coordinate Covalent Bond: A covalent bond in which one atom supplies both the é
that are being shared.e.g. [NH4]+
# é avail: 5 + 1 + 1 + 1 + 1(-1) = 8 é
Octet electron reqd: 8 + 2 + 2 + 2 + 2
= 16 é
8 é (4 bonds)
Polar Covalent Bonding
In many molecules, bonding is neither fully ionic nor fully covalent.
i.e.
H
+
Cl
--------
6
H Cl
The molecule as a whole is electrically neutral; however, chlorine has a stronger pull for
electrons over hydrogen. Thus, the shared electrons will be closer on average to the
chlorine.
H
Cl
Although, the chlorine atom has a stronger pull on the shared electrons, it does not attract
strongly enough to gain complete possession of it. Therefore the bond is not ionic.
Instead, we say that it is a covalent bond in which there is an unequal sharing of
electrons. This is known as a Polar Covalent Bond.
The unequal sharing of the bonded electron pair results in a compound where one end is
slightly positive (δ+) and the other end are slightly negative (δ-).
H
Cl
It is almost as if Cl has gained an electron and become negatively charged while
hydrogen has lost an electron to become positively charged.
Polar Molecule: A molecule that is δ+ on one end and δ- on the other end.
H2O
H
2H
+
O--------
O
H

Each oxygen-hydrogen bond is a polar covalent bond.

Because the hydrogen atoms position themselves around the oxygen in this way,
one side is δ+ and the other side of the molecule is δ-.
Therefore, water had the following shape.
H
O
H
H2O would be non-polar because it would not have opposite δ+ and δ- sides.
When determining whether a molecule is POLAR OR NON-POLAR, we must consider
the following:
1. Polarity of the bond (ionic, covalent, polar covalent
2. Bond/Atom arrangement in a molecule.
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Electronegativities
Electronegativities are used to identify the character of a bond: ionic, covalent, polar
covalent.
Elecronegativity: A measure of the electron attracting ability of the atoms in a molecule
Electronegativities are represented numerically for each atom on the periodic table. The
electronegative values are such that the bigger its number, the greater that atom ‘s ability
to attract electrons.
F------------ 4.0 (has the highest electronegativity)
Ca + Fr ---- 0.7 (have the lowest electronegativity)
By determining the difference in electronegativities between 2 atoms, we can determine
what type of bond is formed.
Covalent
Polar Covalent
Ionic
________________________________________________________
0
Br
.03
+
1.7
Cl
4.0
------
Br
Cl
Br --------- electronegativity of 2.8
Cl---------- electronegativity of 3.0
Properties of Molecular Compounds
1. They are liquid, gases, or soft solid at STP.
2. They are normally non-electrolytes in solution.
3. They have low-melting and boiling points.
4. Covalent bonds are strong intramolucular forces, but weak intermolecular forces
Intramolecular Forces: Bonds that hold atoms together are strong
O------H-------O
Strong
Intermolecular Forces: Bonds that hold the molecules to other molecules
(VanderWaals Forces)
H2O-------H2O------H2O
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Weak bonds
1.5: Chemical Nomenclature
Binary Compounds
A binary compound is a compound that is composed of only two types of monatomic
ions. Examples include ionic compounds and molecular compounds.
Examples of binary ionic compound include sodium chloride, Lithium chloride,
Magnesium oxide etc.
Writing and Naming Binary Compounds
(see P. 90- 91 including multivalent metals)
Some binary compounds of hydrogen (mainly H with the halogen family atoms) dissolve
in H2O to form acid solutions. These are called Binary Acids
When these solutions are not dissolved in water, they should be called by their proper
names.
The aqueous solutions are names as follows:

The binary acid begins with the prefix “hydro” and ends with the suffix “ic”.
Binary Compound
Hydrogen fluoride
Hydrogen chloride
Hydrogen bromide
Hydrogen Iodide
Formula
HF
HCl
HBr
HI
Binary Acid
Hydrofluoric acid
Hydrochloric acid
Hydrobromic acid
Hydroiodic acid
If we want to represent the formula that refers to the acid, we write (aq).
Ex.
HCl-- hydrogen Chloride
HCl (aq) --- hydrochloric acid
Polyatomic Compounds
Tertiary Compounds: Compounds that are usually composed of a metal and a
polyatomic ion. i.e. MgSO4.
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Polyatomic Ion: Groups of atoms that tend to stay together and carry an overall ionic
charge. i.e. NO3-----. Nitrate
Oxyanions: Polyatomic ions that include oxygen
Some Common Polyatomic Ions
Name of Polyatomic Ion
Nitrate
Chlorate
Carbonate
Sulfate
Phosphate
Bicarbonate
Hydroxide
Acetate
Ion Formula
NO3ClO3CO-2
SO4-2
PO4-3
HCO3OHC3H3O2-
Ionic Charge
-1
-1
-2
-2
-3
-1
-1
-1
Note:
The form of the polyatomic ion that occurs most commonly ends with”ate”.
If there is more oxygen --------
If there is less oxygen ----------
If there is even less oxygen----
“per” stem “ate”
stem “ate”
hypo stem “ate”
Examples:
1) Nitrate
NO3NO2-
2) Carbonate
CO3-
3) Chlorate
ClO4ClO3ClO-
4) Sulfate
SO4-2
SO3-2
5) Phosphate
PO4-3
PO3
 Most polyatomic ions are negative but some are positive. Example: NH4+
(Ammonium).
Writing Polyatomic Ions
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Example: Sulfate and Sodium
Rule #1:
Write the symbol of the cation followed by the anion.
Rule #2:
Write the ionic charge/oxidation #’s above each.
Rule #3:
Use the crossover rule to determine the subscripts.
Rule #4:
Reduce if possible
Example2: Aluminum and Carbonate
Naming Polyatomic Compounds
Name the cation followed by the anion. If the polyatomic ion is the anion, name the
polyatomic ion 2nd without changing its name.
Example:
Na + SO4-2
Al + CO3-2
Pb + CO3-2
NH4+ + NO3-
Hydrates
Hydrate: A tertiary compound that contains H2O molecules within its crystal structure
E.g. CuSO4 • H2O (s)
Compound that contains both the tertiary compound, and the water molecule.
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When naming hydrates, we name:
1. Polyatomic compound
2. Water (using the suffix “hydrate”)

You must indicate the amount of H2O using prefixes. i.e. Mono, di, tri etc.
Therefore, CuSO4 • 5H2O
Copper (II) sulfate pentahydrate
When a hydrate is heated, it will decompose to form
1. Water vapour
2. Ionic compound
Therefore, H2O is weakly held in a hydrate.
When the water is removed by heating, the product is referred to as anhydrous.
Anhydrous compound: a tertiary compound without the water
Naming Bases
Most bases are aqueous ionic hydroxide. Therefore, they are composed of a metal (+ve
ion) and an “OH” (-ve ion).
Eg.
NaOH (aq) -------------- Sodium hydroxide
Ba(OH)2 (aq)------------- Barium hydroxide
Name bases much like polyatomic ions.
 Name the caition first followed by the hydroxide.
 The name “hydroxide” does not change.
Acids/Bases
An easy way to distinguish between acids and bases.


Acids ------ Have an “H” at the beginning of the formula and (aq) following.
Bases ---- End with an “OH” and have (aq) following.
Writing Formulas for Molecular Compounds
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The number of electrons that a non-metal need to share indicates how many covalent
bonds it can form. The number of electrons that the non-metal needs to share can be
given by the combining capacity.
Rules for writing molecular compounds
Rule 1: Write the symbol of the atom with the smallest atomic # first, then write the
symbol of the second atom.
Rile2: Write the combining capacity on top of each symbol.
Rule 3: Cross over the combining capacities and write them as subscripts.
Ex.
Carbon and Hydrogen
4
C
1
H
CH4
Naming Molecular Compounds
1.
2.
3.
Write the name of the atom with the smallest atomic number firs. An appropriate
prefix should be used. (If the prefix is “mono”, this is left out for the name of the
first atom.)
Write the name of the second atom with its appropriate prefix. (Here, the prefix
“mono’ is always written.)
Change the ending of the second atom to “ide”.
Number
1
2
3
4
5
6
7
8
9
10
Prefix
Mono
Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca
Relationship between atomic notation, Bohr diagram, é=configuration, and Lewis
Dot diagram
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Atomic Structure
Bohr diagram
é-configuration
1s1
H
He
Lewis
H
Li
Try these: Be, B, C, N, O, F, Ne
Electron dot diagrams can also be used to show the formation of bonds
Chemical Bonding and Quantum Mechanics
Quantum Mechanics can also be used to determine both the number of bonds formed and
the resultant bond angles in a chemical reaction.
Determining the # and type of bonds formed using é-configuration.
This method of bond determination is based on the Valence Bond Theory.
Valence Bond Theory:
 A covalent bond is formed when 2 orbitals overlap to produce a new
combined orbital containing 2 é .
 This arrangement results in a decrease in energy between the bonded
atoms.
 Therefore increased stability.
This approach starts with a central atom and then builds the molecule using overlapping
atomic orbitals to form covalent bonds.
F
+
F
--------- F2 (gas)
Lewis Diagram
F
+
F
---------
Orbital Diagram
F


F
1 covalent bond is formed by combining (overlapping) both half filled 2p orbitals
Any two half-filled orbitals can combine in this way.
Rules for Bonding using Orbital Diagrams
1. Determine the # of bonds formed in the molecule.
2. Identify the central atom and give the orbital diagram for its valence electrons.
14
3. Share é between half-filled orbitals so as to form the predicted # of bonds. (Remember,
a new bonding orbital can only contain 2 é maximum.)
4. Determine the types of bonds that have formed: single, coordinate covalent, lone pairs.
Therefore, the Lewis diagram for H2O:
H
O
H
Example 2:
P2H4
1. # of electrons: 2(8) + 4(2)=24 é
full octet/duet: 2(2) + 4(1) = 14 é
10 é (bonds)
3.
P
P
3. Central Atom: P2
4. – Single bond between each phosphorus
- single bond between P and H
- each P has a lone pair
Therefore Lewis Diagram
H
P
P
H
H
Example 3.
ClO41.
H
# of electrons: 7 + 4(6) + 1 = 32 é
Full octet/duet: 8 + 4(8) = 40 é
3.
Cl
4. – 4 coordinate covalent bonds
formed
2. Central Atom: Cl
Therefore, Lewis Diagram:
Example 4:
SO3
1. # of electrons:
Full Octet/duet:
3.
15
2. Central Atom:
4.
Therefore, Lewis Diagram
Summary of Valance Bond Theory
 Each atom that makes up a molecule behaves as if it were isolated. Therefore
inner é of each atom in a molecule or compound are under the influence of
ONLY that atom’s nucleus.
 A half-filled orbital in one atom can overlap with another half-filled orbital of a
2nd atom to form a new bonding orbital.
 The new, bonding orbital contains a maximum, number of electrons.
Hybridization
The approach we have taken so far in determining the # and types of bonds formed by
atoms work well with simple molecules. The bonds formed are explained well using the
Valence Bond Theory and the overlap of atomic orbital model. However, we can identify
many molecules that fall to fit this model.
Example: Methane CH4
#é
octet
: 4 + 4(1) = 8 é
: 8 + 4(2) + 16 é
8é
C
é-configuration for C: 1s22s22p2
How so we form 4 bonds?
To explain the bonds in molecules such as CH4, we =must examine the ways atomic
orbitals in the same atom interact in order to form bonds----- Hybridization.
Hybridization: The blending of atomic orbitals in the same atom to create a new set of
orbitals that take part in bonding.
Formation of Hybrid Orbitals
Experimental evidence indicates that when atoms form bonds, their simple atomic often
mix to form new orbitals we call hybrid orbitals.
i.e. hybrid orbital designations: sp, sp2, sp3, sp3d1, sp3d2
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