LOGO
Lecture 6:
Ionic Vs Covalent bonding
International University of Sarajevo
Course lecturer :
Jasmin Šutković
8th April 2015 n
Contents
1.
2.
3.
4.
5.
6.
Ionic bonding
Lewis electron structures
Lewis acids and bases
Covalent bonds
Models of covalent bonding
Valence Bond Theory
International University of Sarajevo
Overview of Chemical Bonding
 Chemical bond — the force that holds atoms together in
a chemical compound
 Covalent bonding — electrons are shared between atoms in a
molecules (nonmetals+nonmetals)
 Ionic bonding — positively and negatively charged ions are
held together by electrostatic forces (metal + nonmetals)
 Ions : charged atoms, cations (atoms that loses electrons) and
anions (atoms that gains electrons)
 Ionic compounds — dissolve in water to form aqueous
solutions that conduct electricity!
 Covalent compounds — dissolve to form solutions that
do not conduct electricity!
Chemical bonds share 3 features :
1. Interaction between ATOMS creates aggregates, such as compounds and
crystals, which lower the total energy of the system (because they are
byproducts of the chemicals equations)
2. Energy is required to separate bonded atoms or ions into isolated atoms or
ions.
a. In ionic solids
– ions form a three-dimensional array called a lattice;
–energy is called lattice energy, the enthalpy change that occurs when a
solid ionic compound is transformed into gaseous ions.
b. In covalent solids
– energy is called the bond energy, the enthalpy change that occurs when
a given bond in a gaseous molecule is broken.
3. Each chemical bond is characterized by a particular optimal internuclear
distance called the bond distance.
1.Ionic Bonding
 Ionic bonds are formed when positively and
negatively charged ions are held together by
electrostatic forces (lecture 2).
 Energy of the electrostatic attraction (E) is a
measure of its strength and is inversely
proportional to the distance between the
charged particles ( r ) and directly proportional to
the magnitude of the charges on the ions.
Ionic Bonding cont …
Q is the charge 1 and 2, r is the distance
and k is the proportionality constant ,
equal to : 2.31x10 -28 J*m
NaCl
If Q1 and Q2 have opposite signs ( NaCl)
then E is negative – energy is released
Energy is released when the ionic bond is formed !
Ionic Bonding cont …
 Ionic compounds characteristics
1. Usually rigid, brittle, crystalline substances with flat
surfaces
2. Not easily deformed
3. Melt at relatively high temperatures
These properties are the results of regular ion
arrangement in the crystalline lattice and due to the
strong electrostatic attractive forces between ions with
opposite charges
Lattice Energies in Ionic Solids
 Lattice energy
1. Formation of ion pairs from isolated ions releases
large amounts of energy
2. More energy is released when these ion pairs
condense to form an ordered three-dimensional array
Calculating Lattice Energy
 Lattice energy, U, of an ionic solid can be calculated by
the equation
U > 0.
•
U, a positive number, represents the amount of energy
required to dissociate a mole of an ionic solid into the
gaseous ions:
AB (s) → A + (g) + B - (g)
∆H = U
Q1 and Q2 are the charges on the ions;
ro is the internuclear distance.
The Relationship between
Lattice Energies and Physical Properties
1. Melting point
a. Temperature at which the individual ions have
enough kinetic energy to overcome the attractive forces
that hold them in place
b. Temperature at which the ions can move freely and
substance becomes a liquid
c. Varies with lattice energies for ionic substances that
have similar structures
The Relationship between Lattice Energies
and Physical Properties
2. Hardness
a. Resistance to scratching or abrasion
b. Directly related to how tightly the ions are held
together electrostatically
3. Solubility of ionic substances in water:
a. The higher the lattice energy, the less soluble the
compound in water
2. Lewis Electron Structures
 Lewis dot symbols
1. Used for predicting the number of bonds formed
by most elements in their compounds
2. Dots represent the atoms valence
electrons
3. A single electron is represented as a single dot
4. Number of dots in the Lewis dot symbol is the same as
the number of valence electrons, which is the same as the
last digit of the element’s group number in the periodic
table.
5. Unpaired dots are used to predict the number of bonds
that an element will form in a compound.
The Octet Rule
 Atoms tend to lose, gain or share electrons to reach a
total of eight valence electrons, called an octet.
 The octet rule explains the stoichiometry of most
compounds in the s and p blocks of the periodic table.
 Number eight corresponds to one ns and three np
valence orbital's, which together can accommodate a
total of eight electrons.
Creating a Lewis Dot Symbol
Example Fluorine….has 7 valence electrons
LETS DO AN EXAMPLE WITH N and F ….
How do we proceed …?
Example of covalent
compound – covalent bond
 It is NF3…why ?
 First find the valence electrons of each
element…
 N has 5 and F has 7 ….
 Let arrange them around the elements ….
 First lets look at the N…
What about F ?
To satisfy the
octate rule
we need to
have 3 x F
Each Fluorine atom has only one gap in its outer shell, so it can bond in only one place. Nitrogen has
three gaps in its outer shell so it makes sense to place the Nitrogen in the middle of the molecule and
attach the Fluorine atoms to it. Note that the eight electrons in the outer shell can be moved around to
suit the build of the molecule, though they must be evenly spread across the four quarters of the atom as
outlined in the electron dot structure page.
From Lewis Dots To Lines
Ionic bond example …NaF
Covalent Bonding example
 The H2 molecule
 Two identical neutral atoms
 Contains a purely covalent bond with each
hydrogen atom containing one electron and
one proton and with the electron attracted to
the proton by electrostatic forces
Covalent Bonding in H2
H2 molecule
Lewis electron structures
construction
1. Arrange the atoms to show specific connections – place the
central atom
 The central atom is usually the least electronegative element in the
molecule or ion; hydrogen and the halogens are usually terminal.
2. Determine the total number of valence electrons in the
3.
4.
5.
6.
molecule or ion.
Place a bonding pair of electrons between each pair of
adjacent atoms to give a single bond.
Beginning with the terminal atoms, add enough electrons to
each atom to give each atom an octet (two for hydrogen).
If any electrons are left over, place them on the central atom.
If the central atom has fewer electrons than an octet, use lone
pairs from terminal atoms to multiple (double or triple) bonds
to the central atom to achieve an octet.
What is electronegativity?
Represented with the symbol χ,
is a chemical property that describes the
tendency of an atom or a functional group
to attract electrons .
How t determine electro
negativity ….
H2O example
Example with CH2O
Book page 534 ..
Formal Charges
It is Possible to write more than one Lewis structure for a
substance, and it does not violate the octet rule, but not
all of the Lewis structures may be equally reasonable
and stable.
To determine the most stable Lewis structure we consider
the formal charge on the atoms.
Formal Charges cont ..
Formal charge represents the difference between the
number of valence electrons in the free atom and the
number assigned to it in the Lewis electron structure
Facts regarding the formal charge:
a) the sum of the formal charges on the atoms within a molecule or
ion must equal the overall charge on the molecule or ion
b) the formal charge is NOT the true charge on an atom; simply
used to predict the most likely structure when a compound has
more than one Lewis structure
Calculation of Formal Charge
 Assign electrons in the molecule to individual
atoms according to the following rules:
1. Nonbonding electrons are assigned to the atom on
which they are located
2. Bonding electrons are divided equally between the
bonded atoms
3. For each atom, the formal charge is computed by the
following equation:
formal charge = valence e- - (nonbonding e- + number of
bonds )
free atom
Atom in Lewis structure
Formal charge example
(page 538)
 BH4
Formula :
formal charge = valence e- - (nonbonding e- + number of
bonds)
B ve = 3
Number of nonbonded electrons is 0
Number of bonds arround boron is 4
formal charge of NH3 = 3 v.e – (0 e- + 4 bond. e-) = -1
Properties of Covalent Bonds
In the Lewis bonding model, the number of electron pairs
that hold two atoms together is called the Bond order
 Single bonds have a bond order of one
 Double bonds have a bond order of two
 Triple bonds have a bond order of three
Bond length decreases as bond order increases !
Relationship between bond length and bond order is not
linear!
Bond Energy
Bond energy is the energy required to break a particular
bond
For covalent bonds, bond energies and bond lengths
depend on many factors:




electron affinities,
sizes of atoms involved in the bond,
differences in their electronegativity
overall structure of the molecule.
There is a general trend in that the shorter the bond
length, the higher the bond energy!
The larger the bond energy, the stronger the bond.
The Relationship Between
Bond Order and Bond Energy
The Relationship Between
Bond Order and Bond Energy
 Bonds of different order between like atoms
a. Triple bonds between like atoms are stronger and shorter than double bonds
b. Double bonds between like atoms are stronger and shorter than single bonds
Page 560
Polar Covalent Bonds
One of the four types of Chemical bonds
1. Ionic
2. Covalent
3. Polar covalent — electrons are shared unequally between the
bonded atoms
4. Polar bond — bond between two atoms that possess a partial
positive charge (õ+) and a partial negative charge (õ-)
Polar Covalent Bonds
Bond polarity
Bond polarity is the extent to which it is polar !
Determined largely by the relative electronegativities of the bonded
atoms!
Electronegativity () — ability of an atom in a molecule or ion to attract
electrons to itself
Rules of bond polarity
 A bond is nonpolar if the bonded atoms have equal electronegativities!

If electronegativities of the bonded atoms are not equal, bond is polarized
toward the more electronegative atom!

A bond in which the electronegativity of B (B) is greater than the
electronegativity of A (A) and is indicated with the partial negative charge on
the more electronegative atom
õ+ õ(less electronegative) A --- B (more electronegative)
 To estimate the ionic character of a bond (the magnitude of the charge
separation in a polar covalent bond), calculate the difference in
electronegativity between the two atoms
Net electronegativity = B - A
Electronegativity table
Example for bond polarity
Estimate the polarity of bond in Cl2 and
NaCl ?
 Cl has  = 3.16 , so according the formula
B - A = 3.16-3.16 = 0 and it is NONPOLAR
 NaCl , Na  =0.9 and Cl  = 3.16 = 3.16-0.90
= 2.26 (The electronegativity difference)
Any difference higher than 2.0 proves that the electronegativity of
one of the elements is high enough to remove an electron from the
other atom. This proves that the atoms fulfill the octet rule by
transferring electrons and creating an IONIC bond.
Dipole moments
1. Produced by the asymmetrical charge distribution in a polar
substance ( like in HCl)
2. Abbreviated by µ
3. Defined as the product of the partial charge Q on the bonded
atoms and the distance r between the partial charges
µ=Qxr
Q measured in coulombs (C)
r measured in meters (m)
4. Unit for dipole moment is the debye (D)
1D = 3.3356 x 10-30 C x m
Reading task ….
 Pages : 511- 520, 528-575 (Book chapter
8)
Please read the book according the power
point slides (according the titles and
subtitles).
HW2 by 21th April