Chapter 6
Chemical Bonding
Sect. 6-1: Introduction to
Chemical Bonding
Chemical bond – electrical attraction
between nuclei and valence electrons of
different atoms that binds them together
Being bonded lowers potential energy and
creates more stable arrangement of matter
Types of Chemical Bonding
Ionic bonding – electrical attraction
between cations and anions
Electrons are transferred from one atom to
another forming ions
Covalent bonding - sharing of electrons
between two atoms
Difference in Electronegativities of the 2
atoms determines what type of bond it will
have (chart on pg. 151)
If difference is greater than 1.7, it is
considered ionic, less than 1.7, covalent
Atoms from same element always
covalently bond
Nonpolar-covalent bond – electrons are
shared equally between atoms (difference
of about 0-0.3)
Polar-covalent bond – unequal sharing of
electrons (difference of 0.3-1.7)
One atom is partially negative (δ-)and the other
partially positive (δ+)
Sect. 6-2: Covalent Bonding and
Molecular Compounds
Molecule – group of atoms held together
by a covalent bond
Chemical formula – indicates relative
number and type of atoms in a compound
Molecular formula – chemical formula for a
molecular compound
Diatomic molecule – only 2 atoms of same
element
For a covalent bond to form, repulsive
forces between electrons or protons must
balance out the attractive force between
protons and electrons
Bond length – average distance between 2
bonded atoms
Bond energy – amount of energy required
to break a chemical bond (equal to amount
released when bond was formed)
Pg. 168
Octet rule – atoms gain, lose, or share
electrons in order to have 8 valence
electrons
Exception – boron likes to only have 6
Exception – some elements are surrounded by
more than 8 when paired with highly
electronegative elements such as F, O, & Cl
(called expanded valence)
Electron-dot Notation – shows valence
electrons as dots around the element
symbol, which represents nucleus and
inside electrons
Lewis-structures – uses pairs of dots for
unshared (lone pair) electrons and dashes
for shared electron pairs
Structural formula – similar to Lewisstructure, but does not show lone pair
electrons
Steps to drawing Lewis-structures
1. Determine # and type of atoms in
compound
2. Write electron dot notation for each type
of atom in compound
3. Determine total number of valence
electrons in compound
4. Arrange atoms to form skeleton
structure. Carbon, or least
electronegative element goes in center.
5. Connect atoms by shared electron pairs.
6. Add unshared pairs of electrons until
each atom has 8 electrons (except
hydrogen which only wants 2)
7. Count # electrons used and be sure it
matches # available, if not adjust to form
multiple bonds.
Multiple Covalent Bonds
Double bond – sharing of 2 pairs of
electrons
Triple bond – sharing of 3 pairs of
electrons
The more electron pairs shared, the
shorter the bond length and higher the
bond energy
Resonance – bonding that allows for more
than one correct Lewis structure for a
compound
Represented by double-headed arrow between
the 2 possible structures
Sect. 6-3: Ionic Bonding and
Ionic Compounds
Ionic compound – composed of cations &
anions, paired so that charges balance
Formula unit – simplest ratio for which an
ionic compound’s formula can be written
Use dot structure to show how electrons
transfer to form ionic compounds
Ionic compounds are arranged in a 3-D
structure called a crystal lattice
Lattice energy – energy released when
one mole on an ionic compound is formed
Negative values show that energy is released
Comparison of Covalent/Ionic
Covalent – low melting/boiling because
weaker bonds; soft solids, liquid, or gas;
do not conduct electricity
Ionic – high melting/boiling because
stronger bonds; very brittle & hard;
conduct electricity in liquid state or
dissolved in water, but not in solid state
Polyatomic ion – a group of covalently
bonded atoms with a charge
Will bond with ions of opposite charge to form
ionic compounds
Lewis structures are drawn with brackets
around it and the charge written outside the
brackets
Sect. 6-4: Metallic bonding
 Metallic bonding – chemical bond that results
from the attraction between a sea of electrons
and metal atoms
 Properties
Excellent conductor of electricity & heat
Lustrous - shiny
Malleable – hammered in to thin sheets
Ductile - drawn into wires
 Heat of vaporization – used to measure strength
of bonds in metals
Sect. 6-5: Molecular Geometry
VSEPR Theory (Valence Shell Electron
Pair Repulsion Theory) – repulsion
between valence electron pairs causes
them to spread as far apart as possible
Central atom with 2 atoms around it and no
lone electron pairs will spread apart with 180˚
bond angles and a straight configuration
See pg. 186 for full list
Unshared electron pairs act like an atom,
but bond angles are slightly different
Double and triple bonds act like single
bonds
Hybridization – mixing of 2 or more orbitals
of similar energies on the same atom to
produce new orbitals of equal energy
A 2s and three 2p orbitals can combine to form
four sp3 orbitals that have an energy higher
than the 2s, but less than the 2p
Intermolecular forces – force of attraction
between molecules
Dipole-dipole
Hydrogen bonding
London dispersion forces
Dipole – created by equal, but opposite
charges separated by a small distance
Direction is from positive to negative,
represented by an arrow pointing toward
negative and tail crossed at positive end
Dipole-dipole – forces of attraction
between polar molecules
Induced dipole
Hydrogen bonding – very strong dipole
between hydrogen bonded to a highly
electronegative atom (N, O, F) attracted to
an unshared pair of electrons of an
electronegative nearby atom
London Dispersion Forces – weak
intermolecular forces caused by
instantaneous dipoles
The only type of intermolecular forces acting
between noble gases and non-polar molecules
Forces increase with increasing atomic mass