 Ionic Bonds/Compounds are compounds built from positive ions (cations) and negative ions (anions).
Atoms that can transfer electrons. In the extreme case where one or more atoms lose, electrons and
other atoms gain them in order to produce a noble gas electron configuration. {Ionic compounds are
usually formed when metals bond to nonmetals.}
 Electron Transfer creates anions & cations, which attract because of their opposite charges.
electron acceptor (Cl)
meets electron donor (Na)
ions attract to form neutral
pair
 A covalent compound is a compound in which the atoms that are bonded share electrons rather
than transfer electrons from one to the other. Covalent compounds are formed when two
nonmetals bond to each other.
 Metallic Bonding: Metals lose electrons and cannot normally accept them. This means that, in
a metallic bond, there are no atoms to accept the electrons. Instead, the electrons are given up
to a "sea" of electrons that surrounds the metal atoms.
The cations within a metallic solid are known as Kernels.
The mechanism that holds a metallic bond together is the attraction between the positive kernels and
the negative electron sea. The strength of the metallic bond is derived primarily from the charges in
the system.
 Characteristics of Metal Bonding:
 The larger the magnitude of the positive charges on the metallic nuclei, the greater the
strength of the metallic bond.
 The greater the number of valence electrons contributed to the electron sea, the greater the
strength of the metallic bond.
 Metals are also known as being good conductors of heat, or thermal conductors. Heat is kinetic
energy. In order for a substance to conduct heat, it must be able to transmit kinetic energy.

If heat is applied to one side of a piece of metal, then the kernels will start to vibrate.
Because they are so loosely held into the crystal structure, they will be able to vibrate freely.

With the increase in the amount of their vibration, they will run into the kernels located
next to them. That will start more kernels to vibrate.

In this way, the process continues until all of the kernels in the system are vibrating.
 Structure: smallest building blocks are ions- not molecules!
large numbers of ions can attract to form clusters and eventually crystals
an ion pair
an ion cluster
an ion crystal
Bonds
There are three types of bonds: metallic, ionic, and covalent. The mechanism is an electrostatic
force of attraction between areas of positive and negative charge. These differences are responsible
for the various types of bonds generally found in compounds. Bonds form as an attempt to stabilize
a chemical system by releasing energy. The greater the amount of energy released during the
formation of a bond, the more stable the bond will be. All bond formation processes involve the use
of valence level electrons.
Bond Length is the average distance between the centers of two bonded atoms. Because bonded
atoms experience some vibration, moving towards and away from each other, the distance between
bonded atoms will vary slightly over a period of time.
Why Atoms Combine?
Diatomic Molecules - a molecule made up of two atoms of the same element. Examples include: Bromine
(Br2), Chlorine (Cl2), Fluorine (F2), Hydrogen (H2), Iodine (I2), Nitrogen (N2), & Oxygen (O2)
Lewis Dot Structures - represents the outer most electrons surrounding a central element.
Step 1. Let the symbol of the element represent the nucleus and all electrons except those in their outer
shell.
Oxygen O
Step 2. Write the electron configuration of the outer most electrons.
Oxygen - [He] 2s2 2p4
Step 3. Draw dots around the central element filling in the outer most shells. For example: if the s
orbital is filled completely, the element will show a pair of electrons around it. Place the remaining
electrons around the element that shows the outer most shell.
Oxygen - [He] 2s2 2p4 =
O
Rules for writing Lewis structures:
1. Count up the total number of valence electrons in the species (molecule, ion or atom).
2. Write the symbol of the central atom. This is usually the element of which there is only one
atom. In other situations, you'll need to be told which is the central atom.
3. Write the symbols for the other atom(s) around the central atom any way you please. Precisely
where you choose to put the surrounding atom(s) does not matter since Lewis structures do
not specify the 3-D shape.
4. Put a bond (pair of electrons) between the central atom and each of the surrounding atoms.
Represent the bond with a line.
5. Put dots representing the remaining valence electrons around the atoms so as to attain a noble
gas configuration of a total of 8 electrons (2 in the case of hydrogen). This is called the "octet
rule," but it is often violated (see step 6).
6. If step 5 proves difficult, try multiple bonds. If step 5 still proves difficult, exceed the octet
rule on atoms from elements in the 3rd period or higher. For B or Be atoms, one can have
fewer than 8 electrons.
More than one Lewis structure is often possible for a molecule. To choose the most reasonable, one
must consider formal charge on each atom. Formal charge should always be spread out as much as
possible.
Main properties of Ionic Compounds (salts)
 All ionic compounds form crystals. Salts like to form crystals because when you have electrical positive
and negative charges, they like to be stuck together. The arrangement is referred to as the "unit cell".
 Ionic compounds tend to have high melting and boiling points. So, why are these temperatures so high?
Well, it has to do with the way that ionic materials are held together. To break the positive and negative
charges apart, it takes a huge amount of energy.
 Ionic compounds are very hard and very brittle. Again, this is because of the way that they are held
together. Ionic compounds are hard - so they don't bend at all. This also explains the brittleness of ionic
compounds.
 Ionic compounds conduct electricity when they dissolve in water. If we take a salt and dissolve it in
water, the water molecules pull the positive and negative ions apart from each other. (This is because of the
unusual properties of water). Instead of the ions being right next to each other, they wander all over the
water.
Main Properties of Covalent Compounds
 Covalent compounds generally have much lower melting and boiling points than ionic compounds.
Covalent compounds form distinct molecules, in which the atoms are bound tightly to one another. These
molecules don't interact with each other much (except through relatively weak forces called "intermolecular
forces"), making them very easy to pull apart from each other.
 Covalent compounds are soft and squishy. Covalent compounds have low melting and boiling points.
Covalent compounds have molecules that can easily move around each other, because there are no bonds
between them. Covalent compounds are frequently flexible rather than hard.
 Covalent compounds tend to be more flammable than ionic compounds. Covalent compounds
that do not contain carbon or hydrogen tend not to burn; they burn by mechanisms other than
combustion reaction. The other exception comes with ionic compounds referred to as "organic
salts". These organic salts are ionic compounds in which the anion is basically a big covalent
molecule containing carbon and hydrogen with just a very small ionic section. As a result, they
burn.
 Covalent compounds do not conduct electricity in water. Electricity is conducted in water from
the movement of ions from one place to the other. Since there are no ions in a covalent compound,
they do not conduct electricity in water.
 Covalent compounds are not usually very soluble in water. There is a saying that, "Like
dissolves like". This means that compounds tend to dissolve in other compounds that have similar
properties (particularly polarity). Since water is a polar solvent and most covalent compounds are
fairly nonpolar, many covalent compounds do not dissolve in water. However, some covalent
compounds dissolve quite well in water.
Common Properties of Ionic and Covalent Bonds
Ionic Compounds
1. Crystalline solids
(made of ions)
2. High melting and
boiling points
3. Conduct electricity
when melted
4. Many soluble in
water but not in
nonpolar liquid
Covalent Compounds
1. Gases, liquids, or solids
(made of molecules)
2. Low melting and
boiling points
3. Poor electrical
conductors in all
phases
4. Many soluble in
nonpolar liquids but
not in water
Molecular compounds
Ionic compounds
smallest particles
molecules
cations and anions
origin of bonding
electron sharing
electron transfer
forces between
particles
strong bonds between
atoms
weak attractions
between molecules
strong attractions between
anions and cations
strong repulsions between
ions of like charge
elements present
close on the periodic
table
widely separated on the
periodic table
metallic elements
present
rarely
usually
electrical
conductivity
poor
good, when melted or
dissolved
state at room
temperature
solid, liquid, or gas
solid
melting and boiling lower
points
higher

Elements which are close together in electronegativity tend to form covalent bonds and can
exist as stable free molecules. Carbon dioxide is a common example.

Elements from opposite ends of the periodic table will generally form ionic bonds. They
will have large differences in electronegativity and will usually form positive and negative ions.
The elements with the largest electronegativities are in the upper right of the periodic table, and
the elements with the smallest electronegativities are on the bottom left. If these extremes are
combined, such as in RbF, the dissociation energy is large. Hydrogen is the exception, forming
covalent bonds.
VSEPR Theory:
VSEPR stands for Valence Shell Electron Pair Repulsion. It is a theory, which allows
the user to predict the shapes of simple polyatomic molecules by applying a set of
straightforward rules.
Designation of Similar Regions by Geometry
Taking into Account Lone Pairs and/or Single Unpaired Electrons
You must be able to construct a correct Lewis structure for a given molecule in order to determine
its geometry.
σ-bonded
electron
pairs
2 (180)
Unshared
electron
pairs
0
3 (120)
Stereoactive
pairs
Shape (linked to a
CHIME model)
2
linear (sp)
0
3
trigonal planar
(sp2)
2 (120)
1
3
bent (sp2)
4
(109.5)
0
4
tetrahedral
(sp3)
3
(107.3)
1
4
trigonal pyramidal
(sp3)
2
(104.5)
2
4
bent
(sp3)
5
(90&120)
0
5
trigonal bipyramidal
(dsp3)
4
(90&120)
1
5
see-saw
(dsp3)
3
(90&120)
2
5
T-shaped
(dsp3)
2
(90&120)
3
5
linear
(dsp3)
Sketch
6
(90)
0
6
octahedral
(d2sp3)
5
(90)
1
6
square pyramidal
(d2sp3)
4
(90)
2
6
square planar
(d2sp3)
Rules for Assigning Formal Charges
First you must have a Lewis structure that shows all the lone pairs (nonbonding
electrons). This tutorial is not meant to show you how to draw Lewis structures, but
how to assign formal charges, after you have one to look at.
1. Examine each atom in the Lewis structure, one at a time.
2. Count both electrons in a lone pair (nonbonding electrons) and one electron per
bond.
3. Compare this number with the Group # (from the periodic table), which tells you
how many the atom is supposed to have.
4. The Group # tells you how many electrons a neutral atom would have. So, if you
have one more electron than the Group # indicates, the atom has a formal charge of
−1. If you have two more electrons, the charge is −2, and so forth.
5. If you have one less electron than the Group # indicates, the atom has a charge of
+1. If you have two less electrons, the charge is +2, and so forth.
6. If you prefer to memorize an equation, it would be Formal Charge = Group
Number − (number of nonbonding electrons + number of bonds)
Types of Intermolecular Forces
The physical properties of melting point, boiling point, vapor pressure, evaporation, viscosity, surface
tension, and solubility are related to the strength of attractive forces between molecules. These attractive
forces are called Intermolecular Forces. The amount of "stick togetherness" is important in the
interpretation of the various properties listed above.
There are four types of intermolecular forces. Most of the intermolecular forces are identical to bonding
between atoms in a single molecule. Intermolecular forces just extend the thinking to forces between
molecules and follows the patterns already set by the bonding within molecules.
1. IONIC FORCES:
The forces holding ions together in ionic solids are electrostatic forces. Opposite charges attract each other.
These are the strongest intermolecular forces. Ionic forces hold many ions in a crystal lattice structure.
Ionic bonding is best treated using a
simple electrostatic model. The electrostatic model
is simply an application of the charge principles
that opposite charges attract and similar charges repel.
An ionic compound results from the interaction of a
positive and negative ion, such as sodium and
chloride in common salt.
The IONIC BOND results as a balance between the
force of attraction between opposite plus and minus
charges of the ions and the force of repulsion
between similar negative charges in the electron
clouds. In crystalline compounds this net balance
of forces is called the LATTICE ENERGY.
Lattice energy is the energy released in the formation
of an ionic compound.
DEFINITION: The formation of an IONIC BOND is the result of the transfer of one or more
electrons from a metal onto a non-metal.
Metals, with only a few electrons in the outer energy level, tend to lose electrons most readily. The energy
required to remove an electron from a neutral atom is called the IONIZATION POTENTIAL.
Energy + Metal Atom ---> Metal (+) ion + eNon-metals, which lack only one or two electrons in the outer energy level have little tendency to lose
electrons - the ionization potential would be very high. Instead non-metals have a tendency to gain
electrons. The ELECTRON AFFINITY is the energy given off by an atom when it gains electrons.
Non-metal Atom + e- --- Non-metal (-) ion + energy
2. DIPOLE FORCES:
Polar covalent molecules are sometimes described as "dipoles", meaning that the molecule has two "poles". One
end (pole) of the molecule has a partial positive charge while the other end has a partial negative charge. The
molecules will orientate themselves so that the opposite charges attract principle operates effectively.
In the example on the left, hydrochloric acid is a polar molecule with the partial positive charge on the hydrogen
and the partial negative charge on the chlorine. A network of partial + and - charges attract molecules to each
other.
Bonding between non-metals consists of two
electrons shared between two atoms. In covalent
bonding, the two electrons shared by the atoms
are attracted to the nucleus of both atoms. Neither
atom completely loses or gains electrons as in
ionic bonding.
There are two types of covalent bonding:
1. Non-polar bonding with an equal sharing of
electrons.
2. Polar bonding with an unequal sharing of
electrons. The number of shared electrons
depends on the number of electrons needed to
complete the octet.
POLAR BONDING results when two different
non-metals unequally share electrons between
them. One well known exception to the identical
atom rule is the combination of carbon and
hydrogen in all organic compounds.
The non-metal closer to fluorine in the Periodic Table has a greater tendency to keep its own electron and also
draw away the other atom's electron. It is NOT completely successful. As a result only partial charges are
established. One atom becomes partially positive since it has lost control of its electron some of the time. The
other atom becomes partially negative since it gains electron some of the time.
3. HYDROGEN BONDING:
The hydrogen bond is really a special case of dipole forces. A hydrogen bond is the attractive force between the
hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different
molecule. Usually the electronegative atom is oxygen, nitrogen, or fluorine.
The hydrogen on one molecule attached to O or N that is attracted to an O or N of a different molecule.
The hydrogen bond is really a special case of dipole forces. A hydrogen bond is the attractive force between the
hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different
molecule. Usually the electronegative atom is oxygen, nitrogen, or fluorine, which has a partial negative charge.
The hydrogen then has the partial positive charge.
Comparison of Bond Lengths:
The graphic on the left shows a cluster of water molecules in the liquid state. Water is a polar
molecule, with the oxygen (red) being the negative area and the hydrogen (white) being the more
positive area. Opposite charges attract.
The bond lengths give some indication of the bond strength. A normal covalent bond is 0.96
Angstroms, while the hydrogen bond length is 1.97 A.
4. INDUCED DIPOLE FORCES:
Forces between essentially non-polar
molecules are the weakest of all
intermolecular forces. "Temporary dipoles"
are formed by the shifting of electron clouds
within molecules. These temporary dipoles
attract or repel the electron clouds of nearby
non-polar molecules.
The temporary dipoles may exist for only a
fraction of a second but a force of attraction
also exist for that fraction of time. The
strength of induced dipole forces depends on
how easily electron clouds can be distorted.
Large atoms or molecules with many
electrons far removed from the nucleus are
more easily distorted.
Introduction to Covalent Bonding:
Bonding between non-metals consists of two electrons shared between two atoms. In covalent bonding,
the two electrons shared by the atoms are attracted to the nucleus of both atoms. Neither atom
completely loses or gains electrons as in ionic bonding.
There are two types of covalent bonding:
1. Non-polar bonding with an equal sharing of
electrons.
2. Polar bonding with an unequal sharing of
electrons. The number of shared electrons
depends on the number of electrons needed to
complete the octet.
NON-POLAR BONDING results when two
identical non-metals equally share electrons
between them. One well known exception to
the identical atom rule is the combination of
carbon and hydrogen in all organic
compounds.
IODINE:
Iodine forms a diatomic non-polar covalent
molecule. The graphic on the top left shows
that iodine has 7 electrons in the outer shell.
Since 8 electrons are needed for an octet, two
iodine atoms EQUALLY share 2 electrons.
OXYGEN:
Molecules of oxygen, present in about 20%
concentration in air are also a covalent
molecules . See the graphic on the left the
Lewis symbols.
There are 6 electrons in the outer shell,
therefore, 2 electrons are needed to complete
the octet. The two oxygen atoms share a total
of four electrons in two separate bonds, called
double bonds.
The two oxygen atoms equally share the four
electrons.
Three types of force can operate between covalent molecules:
> Dispersion Forces
also known as London Forces (named after Fritz London who first described these forces theoretically
1930) or as Weak Intermolecular Forces or as van der Waal's Forces (named after the person who
contributed to our understanding of non-ideal gas behavior).
> Dipole-dipole interactions
> Hydrogen bonds
Relative strength of Intermolecular Forces:
Intermolecular forces (dispersion forces, dipole-dipole interactions and hydrogen
bonds) are much weaker than intramolecular forces (covalent bonds, ionic bonds or
metallic bonds).
Dispersion forces are the weakest intermolecular force (one hundredth-one thousandth the strength of a
covalent bond), hydrogen bonds are the strongest intermolecular force (about one-tenth the strength of a
covalent bond).
dispersion forces < dipole-dipole interactions < hydrogen bonds
Dispersion Forces (London Forces, Weak Intermolecular Forces, van der Waal's Forces)
 are very weak forces of attraction between molecules resulting from:
a. momentary dipoles occurring due to uneven electron distributions in neighboring molecules as they
approach one another
b. the weak residual attraction of the nuclei in one molecule for the electrons in a neighboring molecule.
The more electrons that are present in the molecule, the stronger the dispersion forces will be.

Dispersion forces are the only type of intermolecular force operating between non-polar molecules, for
example, dispersion forces operate between hydrogen (H2) molecules, chlorine (Cl2) molecules, carbon
dioxide (CO2) molecules, dinitrogen tetroxide (N2O4) molecules and methane (CH4) molecules.
Hybridization:
CB20 The
 bond (molecular orbitals)
The  bond is formed when side-on overlap of orbitals
is possible. This is the case for p atomic orbitals but not
for s atomic orbitals. Note that d and f orbitals also
permit side-on overlap, but these are not dealt with in
this series. The side-on overlap of p atomic orbitals to
form a bonding  molecular orbital is shown as a
double bond.
The -bond can only be formed when the two atoms
are already bonded by a -bond. The -bond is weaker
than the -bond because side-on overlap of the lobes
of the porbital is less efficient. The total bond strength
of a double bond is therefore greater than that of a
single -bond, but is less than double this value.
If the atoms also have full py orbitals then a second
overlap is possible along the y axis and this gives rise
to a second -bond. This is shown as a triple bond.
Aim: To show the spatial difference between the
- and the -bond.
C – C single bond 0.154 (Bond distance 0.154 nm)
C = C double bond 0.134 (Bond distance 0.134 nm)
C  C triple bond 0.120 (Bond distance 0.120 nm)
CB21a The formation of sp and sp2 hybrid orbitals
The two sp-orbitals lie on the x axis and are at an
angle of 180° to one other (linear ). The interaction
of the more voluminous lobe of each orbital with
another orbital leads to an effective overlap. The
energy gain during bonding is the source of the
hybridization phenomenon.
sp-hybridisation: 2 sp-hybrid orbitals per atom.
E.g. for C:
Aim: • To explain the term hybridisation
• To explain the various possible combinations
of hybridization
If we had no knowledge of hybridisation or
recombination of orbitals we would expect for the
compound BeH2 two different
Be-H bonds: one using the 2s orbital of Be and the
other using the 2p orbital. However, the molecule is in
fact linear with two equivalent bonds.
The same is true for carbon, which has two unpaired
valency electrons and should therefore only be able to
form two bonds. Carbon can in fact make four bonds.
sp2-hybridisation:
If one s-orbital and two p-orbitals  combine , then
three identical sp2 hybrid orbitals  are formed. These
lie in the same plane at an angle of 120° (trigonal) .
sp2-hybridisation : 3 sp2-hybrid orbitals per atom.
E.g. for C:
Hybridisation, or the combination of atomic orbitals can
explain this. Three types of hybridisation are described
here: sp, sp2 and sp3.
sp hybridisation:
One s-orbital and one p-orbital  are combined  to
form two identical orbitals . These are called sphybrid
orbitals. This transition happens without any energy
change, involving only a simple linear combination of
orbitals or wave functions.
Important :
Hybrid orbitals are atomic orbitals. They can, of
course, overlap in the same way as s- and p-atomic
orbitals to form molecular orbitals. The lobes of the
hybrid orbitals are simply differently oriented in space,
to make more effective overlap possible.
CB21b The formation of sp3 hybrid orbitals
These are oriented in such a way that the lobes are as
far away from each other as possible, i.e. they point
toward the corners of a tetrahedron .
sp3-hybridisation: 4 sp3-hybrid orbitals per atom.
E.g. for C:
Aim: • to explain the term hybridisation
• to explain the various possible combinations of
hybridization
sp3-hybridisation
If three p-orbitals and one s-orbital  combine
four identical sp3 hybrid orbitals arise .
Note :
The hybrid orbitals are represented here with one large
and one small lobe (asymmetrical p-orbital). It is clear
that in this type of combination the sp3-orbital will have
more p-character than the sp-orbital. For this reason
the difference in size between the large and small lobe
will be smaller for the sp3-orbital than for the sp2orbital and largest for the sp-orbital.
, then
For the sake of clarity this relative difference in size is
not given on the illustration
CB22 The molecular orbitals of ethane
Aim: To show a worked example: the molecular
orbitals of ethane.
In ethane the two carbon atoms adopt an sp 3hybridisation. The C-C and C-H bonds are all  types.
The angle between these bonds is approximately 109°,
as in a tetrahedron.
CB23 The molecular orbitals of ethene
Aim: To show a worked example: the molecular
orbitals of ethene
In ethene the two carbon atoms adopt an sp 2hybridisation. The two C-atoms and the four H-atoms
are situated in the same xy-plane. The -bond between
the two 2pz orbitals lies at right angles to this xy-plane.
The two C-H bonds and the C-C -  bond are at an
angle of 120° to each other.
CB24 The molecular orbitals of ethyne
Aim: To show a worked example: the molecular
orbitals of ethyne.
In ethyne the two carbon atoms adopt and sphybridisation. Ethyne is a linear molecule. The two C
atoms and two H atoms are to be found in the xy plane.
The -bond between the two 2pz orbitals is at right
angles to this xy plane (in the xz plane). The second bond between the two 2py orbitals is at right angles to
the xz plane, that is in the xy plane.
The C-H bond and the C-C bond are at an angle of 180°
to each other.