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Chemical Bonding 2: Types of Chemical Bonding
Prior knowledge:
Continue from the previous section, we should all remember the trend DOWN and
ACROSS the periodic table to learn the different chemical bonding.


(a)
When going DOWN a family in the periodic table, properties are affected by the
INCREASING SIZE of the atoms and the INCREASING DISTANCE between
the nuclei and the valence electrons.
When going ACROSS a period in the periodic table, properties afre affected by
the DIFFEREING VALENCE, NUCLEAR CHARGE and CHARGE ON THE
SPECIES
Ionic Bonding
Definition: an IONIC BOND is formed by the attraction of positive ions to negative ions
And ionic bond from when an electron from one atom is transferred to another atom, so
as to create one positive ion and on negative ions.
Definition: a LEWIS STRUCTURE, also called an “electron dot digram”, is a diagram
showing how the valence electrons are distributed in an atom , ion, or molecule
Example: Li+ and F- formed by Li and F atoms
Example: Ions formed by Ca and O atoms
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A simple fact makes it easier to deal with charged sphere.

If a sphere possess a charge, the carge can be assumed to be concentrated at the
centre of sphere, rather than being spread over the sphere’s surface or
throughout its volume.
To see the importance of the above fact, assume that two carge sphere touch one
another. All the electraostatic attraction now depends on the distance separateing the
central charges: The greater the distance, the smaller the attraction.
Important:
How to predict when an ionic bond will form.
IONIC BONDS are formed when element from opposite side of the periodic table are
combined. That is, when a METAL and a NONMETAL are combined
Question:
Which one of the following contain ionic bonds?
NaCl
BrCl
ELECTRONEGATIVITY
The concept of ELECTRONEGATIVITY explains why Li and F form an ion
Definition: the ELECTRONEGATIVITY of an atom is the tendency of the atom to attract
electrons from a neighboring atom.
Recall from the previous class, when we go ACROSS the periodic table, the radius of
atoms get
. The attraction between nucleus and the outer electrons
becomes
. Therefore the electronegativity
across the
periodic table
When we go down the periodic table, the radius of the atoms get
The attraction between nucleus and the outer electron becomes
Therefore the electron negativity become smaller.
.
.
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Question:
Ignoring the noble gas, which atom is the most electronegative and least electronegative
on the third period?
If an atom has a HIGH electronegativity, it strongly attracts electrons from a
NEIGHBOURING ATOM and may completely remove an electron from the neighbouring
atom. This high attraction to a neighbouring atom’s electrons also means that atoms with
a high electronegativity STRONGLY ATTRACT THEIR OWN VALENCE ELECTRONS.
As a results, these valence electrons are difficult to remove and the atom has a HIGH
IONIZATION ENERGY.
If an atom has LOW electronegativity, it has little attraction to the electrons of
NEIGHBOURING ATOM and little tendency to remove electrons from a neighbour. Such
an atom also has a relatively small attraction to its OWN VALENCE ELECTRONS.
Therefore, these valence electrons relatively easy to remove and the atom has a low
ionization energy
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Explain the formation of ion between Li and F
How to measure the strength of ionic bonds
In reality, the strength of ionic bond is measured by the MELTING TEMPERATURES of
ionic compounds.
For example:
Ion
Melting
Temperature
LiF
845 OC
NaF
993 OC
KCl
770 OC
LiCl
605 OC
These high melting temperature support the following conclusion.

IONIC BONDS are VERY STRONG, so that compounds held together by ionic
bonds have HIGH MELTING TEMPERATURES.
In other words, a great deal of energy must be added to break ionic bonds.
Note: the formula for solid lithisum fluoride is written as LiF(s), whihch might suggest
that the lif and F are neutral atoms to make LiF. In fact, solid LiF is actually an “ionic
lattice” in which ions are packed together in an orderly arrangement to make a crystal of
LiF.
Going across the periodic table, a simple rule governs the charge created when an atom
forms an ion
In general, when an atom forms an ion the atom loses or gains sufficient electron to attain
a closed shell
The most common charges found when going across the periodic table are shown below.
[Elements in column 14 (C, Si, Ge, Sn, and Pb) are not includes because C, Si, and Ge do not
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form simple ionic compounds and Sn and Pb are metals which most readily form +2 ion and
only rarely, if ever, form +4 ions]
Column
in
periodic
table
Charge
on ion
1
2
13
+1
+2
+3
14
15
16
17
18
-3
-2
-1
0
Something to think about
(b)

Compare the size of O and O2-

Compare the size of Na and Na+
Covalent Bonding
Definition: a COVALENT BONDis a bond which involves the equal sharing of electrons.
A covalent bond is formed when two atoms having less than full shells of electrons are
able to share one or more of their electrons with each other to attain full electrons shells
Definition: the OCTET RULE states that atoms in columns 14 to 17 of the periodic table
tend to form covalent bonds so as to have eight electrons in their valence shells.
The Lewis structure below shows the covalent bond between fluorine atom to covalent
bonds
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Important:
How to predict when a covalent bone will form?
COVALENT BOND are formed when a nonmetal combines with a nonmetal.
Covalent bonds form when both atoms involved have relatively large electronegativities.
In generals, nont metal have large to very large electronegativities, attrat each other’s
electrons strongly and will not let go of their own electrons. Since both F atoms (above)
have the same ELECTRONEGATIVITY, they attract the shared electrons equally and
neither wins the “tug of war” for the electrons. Therefore, they MUST SHARE the
electrons in the bond.
Predicting the formula of covalent compounds
Predict the formula of the compound formed by N and F
Predict the formula of the compound formed by C and O
(c)
Intermolecular forces
So far, both IONIC BOND and COVALENT BOND are considered to be Intramolecuar
FORCES. “Intra-” means within. They are strong forces holding molecular together
WITHIN.
In addition to the bonds holding atoms together into molecular units, there are weak
forces which hold one complete, neutrally-charged molecule next to another such
molecule. These INTERMOLECULAR FORCES (“inter” means between) are called van der
Waals Forces
There are two types of intermolecular forces: 1) Dipole-Dipole forces and 2) London
forces
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1) Dipole- Dipole forces
 Formed with polar covalent molecules
 Partially positive (δ+) and partially negatively (δ-) attract each other
 Example:

Special case: Hydrogen bonding (occurs when H bonds with F, O, and N- large
electronegativity differences)
2) London forces
 Formed by temporary (instaneous) charges on an atom when electrons move to
unsymmetric position around the nucleus
 London forces are present between any molecules (including polar molecules when
they are close together)
 The weakest type of bond (unsymmetric electron location isn’t always occurring to
cause the attraction)
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Exercise
Ionic bond
1. which of the following atom pair would you expect to form ionic bond when they
join?
a) Ba and S
c) Ca and O
e) O and H
b) P and Cl
d) Rb and I
f) S and O
2. Which compound has the smallest distance between the nuclei of the two ions
involved? NaCl or KBr?
3. Following question 2, what happens to the forces of electrostatic attraction
between the two ions in an ionic bond as the ion get smaller?
4. Following question 3, what happens to the ionic bond as the ions involveds get
smaller? What happen to melting temperature?
5. Mg2+ and Na+ have roughly the same ionic radius. O2- and F- have roughly the same
ionic radius. Which substance should have a higher melting point: NaF or MgO?
Why?
Covalent bond
1. Which of the following atom pair would you expect to form covalent bonds why they
join?
a) S and O
c) Fe and Cl
e) H and S
b) Ba and O
d) N and O
f) C and H
2. Predict the formula of the
a) P and Cl
e)
b) B and O
c) C and S
f)
d) P and O
g)
compound formed by bonding together the following
H and
h) N and I
l) Si and
Se
i) B and C
S
F and O
j) C and Cl
H and O
k) Si and P
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3. When the distance between two covalently-bonded atoms increase, what happens to
the electrostatic attraction of their nuclei to the shared electrons in a covalent
bond?
4. What would you expect to happen to the strength of the covalent bond between
two identical halogen atoms when going down the halogen family from F2 to I2?
5. What would you expect to happen to the strength of a covalent bond when the
number of shared electrons increases?
Mixed Concepts
1. Ionic bond, Covalent Bond, Hydrogen Bond, and London forces. Which two of them
are intermolecular forces? Which two of them are intramolecuar forces? Explain
2. What type of bond would be expected to form between atom of each of the
following
a) H and O
d) P and O
g) I and F
b) K and I
e) Ca and Br
c) Ni and Cl
f) N and F
3. Which of the following have a higher electronegativity
a) F or I
c) Si or Cl
b) Rb or Li
d) Na or Si
e) Sn or S
f) N or Sb
4. Which of the following would be expected to form a stronger bond?
a) NaCl or KBr
c) Kf or CaO
e) C-C or Si-Si
b) C-O or Si-S
d) CsI or MgO
f) B-F or
5. Which is expected to be larger?
a) Na+ or Nad) Na+ or Cs+
b) I or I+
e) Se- or S
c) As or As3f) S2- or S
g) Ca2+ or Se2h) O2- or S2i) K+ or Cl-
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6. (Advanced) Which of ionic bonding, covalent bonding or london forces are most
important in holding each of the following substance together in solid phase
a) N2(s) melts at -210 OC
b) RbBr(s) melts at 693 OC
c) CS2(s) melts at -111 OC
d) Ge(s) a network structure melts at 938 OC
e) CaCl2(s) melts at 782 OC
f) Xe(s) melts at -112 OC