 Chemical Bonding Activity
1. Ionic Bonding
2. Covalent Bonding
3. Metallic Bonding
 Ionic Bonding: chemical bonding that results
from the electrical attraction between cations
and anions.
 Electrons are completely transferred from one
atom to another.
 Covalent Bonding: results from the sharing of
electron pairs between to atoms.
 Metallic Bonding: The chemical bonding that
results from the attraction between metal
atoms and the surrounding sea of electrons.
 Bonding between atoms of different
elements is rarely purely ionic or purely
covalent.
 Depends on how strongly the atoms of
each element attract electrons
 Electronegativity
 Bond polarity helps to describe the sharing of
electrons between atoms.
 Nonpolar-covalent bond: a covalent bond in
which electrons are shared equally by the
bonded atoms, resulting in a balanced
distribution of electrical change.
 Polar-covalent bond: a covalent bond in which
the bonded atoms have an unequal attraction for
the shared electrons. If the difference is big
enough, we get an ionic bond.
 Linus Pauling makes electronegativity
scale.
 Values are unitless
 Scale: 0.7 (cesium) - 4.0 (fluorine)
 Metals are less electronegative; nonmetals
are more electronegative.
 The greater the difference in electronegativity
between two bonded atoms, the more polar the
bond.







Difference < 0.3 = nonpolar-covalent
Ex: F2 4.0 - 4.0 = 0
0.3 ≤ difference < 1.7 = polar-covalent
Ex: HCl 3.0 – 2.1 = 0.9
Difference ≥ 1.7 = ionic bond
Ex: LiF 4.0 - 1.0 = 3.0
Use + and - to show partial positive and negative
charges on atoms.
Nonpolar-covalent bond
Polar-covalent bond
+
-
+
-
Ionic bond
 Find the difference in electronegativity, the type of
bond, and the more-negative atom in the following
compounds:
 Bonding between sulfur and…
1. Hydrogen
2. Cesium
3. Chlorine
 Electronegativities: Sulfur (2.5), Hydrogen (2.1),
Cesium (0.7), and Chlorine (3.0)
 Tell which bond is more polar and indicate
in each case which atom has the partial
negative charge.
 B-Cl or C-Cl?
 B-Cl, with - on Cl
 P-F or P-Cl?
 P-F, with - on F
 Entire molecules can be polar, not just bonds within
the molecules. Ex: HF
 Polarity is important--it helps to determine many
properties of a compound!
 Many chemical compounds are composed of
molecules.
 A molecule is a neutral group of atoms that are held
together by covalent bonds.
 A chemical compound whose simplest units are
molecules is called a molecular compound.
 The composition of a compound is given by its
chemical formula.
 A chemical formula indicates the relative numbers of
atoms of each kind in a chemical compound by using
atomic symbols and numerical subscripts.
 A molecular formula shows the types and numbers of
atoms combined in a single molecule of a molecular
compound.
 Nature favors chemical bonding
 Lower potential energy for atoms when they are bonded
together
 The approaching nuclei and electrons are attracted to
each other
 At the same time…
 The two nuclei repel each other
 The two electron cloud repel each other
 As atoms come closer to each other, the attractive
forces eventually win and bonding occurs!
Covalent compounds…
1. Generally have much lower melting and boiling
points than ionic compounds.
2. Are soft and flexible.
3. Tend to be more flammable than ionic compounds.
4. Don’t conduct electricity in water.
5. Aren’t usually very soluble in water.
 The average distance between two bonded atoms is
defined as the bond length.
 Bond energy is the energy required to break a
chemical bond and form neutral isolated atoms.
Bond
C-C
Average Bond
Length (pm)
154
Average Bond
Energy (kJ/mol)
346
C-H
109
418
C-Cl
177
327
O-H
96
459
C-Br
194
285
 Octet Rule: Chemical compounds tend to form so that
each atom, by gaining, losing, or sharing electrons, has
an octet of electrons in its highest occupied energy
level.
 Exceptions: H, He, B
 Expanded valence: bonding that involves electrons in
d-orbitals as well as s- and p-orbitals.
 Lewis structures: formulas in which atomic symbols
represent nuclei and inner-shell electrons, dot-pairs or
dashes between two atomic symbols represent electron
pairs in covalent bonds, and dots adjacent to only one
atomic symbol represent unshared electrons.
 A structural formula indicates the kind, number,
arrangement, and bonds, but not the unshared pairs of
the atoms in a molecule.
 Add up the number of electrons needed for each atom
and the number of electrons available for each atom.
 Subtract values to determine the number of electrons
involved in chemical bonds.
 Write the symbols for the atoms in an arrangement
which shows their order in the bonded state.
 Try placing an “odd” numbered atom in the center of the
structure or the element that is least electronegative
 Consider number of bonds that the atom would like to
make
 Add dots or single lines between each pair of bonded
atoms to represent a shared pair of electrons. Deduct
2 electrons from the total number of electrons
available for bonding .
 Check that all atoms are surrounded by a full octet
 Add in electrons as dots (in pairs whenever possible) to
satisfy the octet or duet rule for each atom as
appropriate.
 Use multiple bonds when there are not enough
electrons to complete octets
 A single covalent bond is a covalent bond in which
one pair of electrons is shared between two atoms.
 Examples:
 Lewis dot structure: bonds represented by two
dots
 Lewis dash structure bonds represented by a
dash mark
 Each bond accounts for 2 electrons
Multiple Bonds:
 A double covalent bond is a covalent bond in which
two pairs of electrons are shared between two atoms.
 A triple covalent bond is a covalent bond in which
three pairs of electrons are shared between two atoms.
Bond
Average Bond Length
(pm)
Average Bond Energy
(kJ/mol)
C-C
154
346
C=C
134
612
120
835
O-H
96
459
C-Br
194
285
C
C
 Certain atoms bond covalently with each other to form
a group of atoms that has both molecular and ionic
characteristics.
 A charged group of covalently bonded atoms is known
as a polyatomic ion.
 Excess of electrons (negative)
 Shortage of electrons (positive)
 Polyatomic ions combine with ions of opposite charge
to form ionic compounds.
 Lewis Structures:
 Negative sign: Add electrons
 Positive sign: Subtract electrons
 Add brackets
 Resonance refers to bonding in molecules or ions that
cannot be correctly represented by a single Lewis
structure.
 Resonance structure: atoms keep same arrangement
but placement of electrons changes. Look for changes
in placement of double bonds.
 Resonance increases the stability of a compound or ion
 An ionic compound is composed of
positive and negative ions that are
combined so that the numbers of
positive and negative charges are
equal.
 Crystalline Solids
 The chemical formula of an ionic
compound shows the ratio of the
ions present in a sample of any size.
 A formula unit is the simplest
collection of atoms from which an
ionic compound’s formula can be
established.
 Complete transfer of electrons
 The ions in an ionic compound lower their potential
energy by forming an orderly, three-dimensional array
in which the positive and negative charges are
balanced.
 In an ionic crystal, ions minimize their potential
energy by combing in an orderly arrangement known
as a crystal lattice.
 To compare bond strengths in ionic compounds,
chemists compare the amounts of energy released
when separated ion in a gas come together to form a
crystalline solid.
 Lattice energy is the energy released when one mole
of an ionic crystalline compound is formed from
gaseous ions.
Compound
Lattice Energy (kJ/mol)
NaCl
CaF2
-787.5
-2634.7
LiCl
MgO
-861.3
-3760
Ionic compounds…
1. Generally have high melting and boiling points.
2. Are hard but brittle.
3. In solid state compounds cannot conduct electricity.
a.
4.
However, compounds can conduct electricity in water.
Why?
Are usually very soluble in water.
 The highest energy levels of most metal
atoms are occupied by very few electrons.
 Some sublevels may even be vacant
 Vacant orbitals in the atom’s outer energy levels
overlap.
 Allows the outer electrons of the atoms to roam freely
throughout the entire metal
 These electrons are delocalized which means that they
do not belong to any one atom but move freely about
the metal’s network of empty atomic orbitals.
 These mobile electrons form a sea of electrons around
the metal atoms, which are packed together in a
crystal lattice.
 The chemical bonding that results from the attraction
between metal atoms and the surrounding sea of
electrons is called metallic bonding.
 Different chemical bonding that ionic or covalent
compounds
 Metallic compounds…
1. Excellent electrical conductors (solid or molten)
a.
High electrical and thermal conductivity
2. Strong absorbers and reflectors of light
a.
Metallic luster
3. Malleability: the ability to be hammered into thin
sheets
4. Ductility: the ability to be drawn, pulled, or extruded
through a small opening to produce a wire
 The properties of molecules depend not only on the
bonding of atoms but also on molecular geometry- the
3D arrangement of a molecule’s atoms in space.
 Consider locations of non-bonding electron pairs and
number of orbitals involved in bonding
 VSEPR theory: Valence-Shell, Electron-Pair Repulsion
 States that repulsion between the sets of
valence-level electrons surrounding an atom
causes these sets to be oriented as far apart as
possible.
 First, we will examine molecules with no unshared
electron pairs.
 Shared electron pairs want to be as far away from each
other as possible
 Each 3D shape will be characterized by a particular bond
angle
 VSEPR Theory Activity
 Determine the value for different bond angles
 Keep track of shapes and bond angles on a separate
piece of paper
 Use the VSEPR Theory to predict the molecular
geometry of the following molecules:
 BCl3
 HI
 CBr4
 First, draw the Lewis Structure
 How many atoms are bonded to the central atom?
 A pair of unshared electrons will take up more space
than a pair of electrons in a bond. Why?
 Unshared electron pairs repel electrons more strongly
than do bonding electron pairs.
 VSEPR Theory treats multiple bonds the same way as
single bonds
 The shapes of polyatomic ions can also be
determined.
 Use the chart on pg 200 of your text to fill in
the VSEPR Theory and Molecular Geometry
table in your notes.
 Use the VSEPR Theory to predict the shape of the
following molecules:
 CO2
 CF4
 PCl3
 SeF6
 I2S
 How do orbitals in a molecule become rearranged?
 Hybridization is the mixing of two or more atomic
orbitals of similar energies on the same atom to
produce new hybrid atomic orbitals of equal
energies.
 Example: Carbon
 Hybrid orbitals are orbitals of equal energy produced by
the combination of two or more orbitals on the same
atom
 The number of orbitals involved in hybridization
determines the geometry of a molecule
 The forces of attraction between molecules are known
as intermolecular forces.
 Intermolecular forces vary in strength but a generally
weaker than bonds that join atoms in molecules or
compounds.
 Intermolecular forces exist between the
following molecules…
Polar-Polar
Polar-Nonpolar
Nonpolar-Nonpolar
 Polarity is a result of the following two factors:
1. Existence of polar covalent bonds
2. The nature of the molecule’s geometry
 Dipoles within a molecule can
cancel each other out, resulting
in a nonpolar molecule.
 The strongest intermolecular forces exist between
polar molecules.
 Polar molecules act as tiny dipoles
 A dipole is created by equal but opposite charges that
are separated by a short distance.
 Dipole-Dipole forces are the forces of attraction that
exist between polar molecules
 Induced-Dipole forces occur when a polar molecule
induces a dipole on a nonpolar molecule by
temporarily attracting its electrons .
 Some hydrogen containing compounds have unusually
high boiling points.
 Presence of strong type of dipole-dipole forces
 In compounds containing H-F, H-O, or H-N bonds,
the large electronegativity differences between
hydrogen and these atoms make the bonds connecting
them highly polar.
 The intermolecular force in which a hydrogen atom
that is bonded to a highly electronegativity atom is
attracted to an unshared pair of electrons of an
electronegative atom in a nearby molecule is known as
hydrogen bonding.
 Even nonpolar atoms experience a weak
intermolecular attraction.
 Electrons in continuous motion
 At any time, electron distribution may be slightly
uneven
 This temporary dipole may induce a dipole in an
adjacent atom or molecule.
 The intermolecular attractions resulting from the
constant motion of electrons and the creation of
instantaneous dipoles are called London Dispersion
Forces.