Chemical Bonds
Elements form bonds to be in a lower energy state
1. Ionic Bonds – transfer of electrons, between
metal and nonmetal
2. Covalent Bonds – sharing of electrons, between
two nonmetals
3. Metallic Bonds- neighboring atoms in solid
metals form bonds
Octet rule: atoms tend to gain, lose, or share
electrons until they are surrounded by eight valence
electrons to achieve a stable octet (noble gas
configuration)
Electron Dot Symbols
Valence electrons: reside in the highest occupied
energy level, reside in the outer s & p orbitals and
are the electrons involved in chemical bonding.
Electron-dot symbols are convenient way of
showing the s & p electrons & tracking them in
bond formation
-They consist of the chemical symbol for the
element plus a dot for each valence electron
Consider sulfur whose electron configuration is
[Ne]3s23p4, thus there are six valence electrons:
S
Ch. 7 – Ionic and Metallic Bonding
I
II
III
IV
B. Ionic Bonds
• Ionic Bonds – atoms transfer electrons from a
cation (positive ion) to an anion (negative ion) to
achieve an octet.
• Ionic compounds are stable due to the
electrostatic forces between unlike charges
organizing the ions of ionic substances into
a rigid, organized three-dimensional
arrangement:
• The ions are drawn together
• Energy is released
• Ions form solid lattice structure
Lattice Energy
•Lattice energy:
•The energy required to completely separate
a given quantity of a solid ionic compound
into its gaseous ions.
•Thus, in reverse, the high energy is given
off as heat and light when Na+ and Cl - is
incorporated into the NaCl salt lattice.
Steps in Ionic Bonding Process (1)
• Ionization Energy (IE) Step 1:
•The minimum energy required to remove an electron from
the ground state of an isolated gaseous atom or ion.
•First ionization energy:
•Na(g) + IE1 → Na+(g) + e- ; +∆H (positive/ endothermic)
•Second Ionization Energy:
•Be(g) + IE1 + IE2 → Be2+ (g) + 2 e- ; +∆H (positive)
•The greater the IE, the more difficult it is to remove an e-.
Steps in Ionic Bonding Process (2)
Electron Affinity (EA): Most atoms can attract e- to
form negatively charged ions
•The energy change that occurs when an e- is added to
a gaseous atom. For most atoms, the energy released
when an e- is added.
•Cl(g) + e- → Cl- (g) + electron affinity
; -∆H
(negative/exothermic)
•The greater the attraction between a given atom
and an added e-, the more negative the atom’s EA.
Halogens’ –s2p5 have the most negative EA.
Steps in Ionic Bonding Process (3)
• Lattice Energy (LE)
•The release of energy that occurs when ions of
opposite charge are attracted to each other and form a
stable ionic compound.
• Na+ + Cl- → NaCl + lattice energy ; -∆H
(negative/exothermic)
•Ionic compounds have very large LE which makes up
for endothermic ionization energy.
B. Properties of Ionic Compounds
• Most ionic compounds are crystalline
solids at room temperature
• Arranged in repeating threedimensional patterns
• Ionic compounds generally have high
melting points
• Large attractive forces result in very
stable structures
B. Properties of Ionic Compounds
• Ionic compounds can conduct an electric
current when melted or dissolved in
water
• When ionic compounds are dissolved
in water the crystalline structure breaks
down so ions are able to move freely
which results in conductivity
The positive Na ions move to the cathode and
the negative Cl ions move to the anode
Ch. 7 – Ionic and Metallic Bonding
III. Bonding in Metals
(p. 201 – 203)
I
II
III
IV
A. Metallic Bonding
Metallic bonds: Consist entirely of metal
atoms.
• Bonding is due to valence electrons which
are delocalized throughout the entire solid
• The metal is held together by the strong
forces of attraction between the positive
nuclei and the delocalized electrons.
B. Metals
• Metals are good conductors of heat and
electricity because the valence electrons
are able to flow freely
• Valence electrons of metals can be thought
of as a ‘sea of electrons’, very mobile
C. Metallic Bond
Metallic Bonding - “Electron Sea”
D. Metallic Properties
• Have luster, are ductile and malleable
• Luster = shine
• Ductile = ability to be drawn into wires
• Malleable = ability to be shaped,
pounded, etc
D. Metallic Properties
• Properties can be explained by the mobility
of electrons in metals
• When subjected to pressure , the
cations easily slide past each other like
a ball bearing immersed in oil.
B. Types of Bonds
METALLIC
Bond
Formation
Smallest Unit
Physical
State @ RT
Melting
Point
e- are delocalized
among metal atoms
“electron sea”
solid
very high
Solubility in
Water
no
Electrical
Conductivity
yes
(any form)
Other
Properties
malleable, ductile,
lustrous
Covalent Bonding
What is a covalent bond?
• A chemical bond that results from the sharing
of electrons, to form a stable octet or duet
(Hydrogen only needs 2 to be stable)
• Molecule = two or more atoms
that are held together by
covalent bonds
H 2O
• Majority of covalent bonds form between
nonmetals (CLOSE together on periodic table)
Covalent Bonding Formation
• Diatomic molecule
– molecule containing the same two atoms
• Some elements always exist this way
because they are more stable than the
individual atoms
Cl2
B. Diatomic Elements
• The Seven Diatomic Elements
Br2 I2 N2 Cl2 H2 O2 F2
H
N O F
Cl
Br
I
Bonds in all the
polyatomic ions
and diatomics
are all covalent
bonds
Single Covalent Bonds
Two atoms held together by a sharing
of one pair of electrons are joined
together by a single covalent bond.
Single Covalent Bonds
An electron dot structure represents the
shared pair of electrons of the covalent bond
by two dots.
A structural formula represents the covalent
bonds by dashes and shows the
arrangement of covalently bonded atoms
Single Covalent Bonds
A pair of valence electrons that is not
shared between atoms is called an
unshared pair, also known as a lone pair
of a nonbonding pair.
Lone pair
Double and Triple Covalent Bonds
Atoms form double or triple covalent
bonds if they can attain a noble gas
structure by sharing two or three pairs of
electrons.
A double bond involves sharing two pairs
of electrons.
A triple bond involves sharing three pairs
of electrons.
Bond Length
•From a study of various Nitrogen containing
compounds bond distance as a function of bond
type can be summarized as follows:
•N-N 147pm
N=N 124pm
N≡N 110pm
•As a general rule, the more e- that are shared:
•the stronger the covalent bond (N≡N > O=O > F–F)
•the shorter the covalent bond (N≡N < O=O < F–F)
Double and Triple Covalent Bonds
Molecular Structure
Lewis Diagrams
(p. 220 – 229)
I
II
III
Drawing Lewis Diagrams
1. Arrange atoms
•
•
•
Singular atom is usually in the center (often
Carbon)
If not Carbon, least e- neg atom is in center
Hydrogen is always terminal
2. Find total # of e- available to bond
(valence e- )
3. Place a pair of electrons between central
atom and each terminal atom
Drawing Lewis Diagrams
4. Place remaining electrons in pairs around
terminal atoms (except H) to satisfy octet
rule
• Any remaining pairs are assigned to
central atom
5. Determine whether or not central atom
satisfies octet
• If not, convert one or more lone pairs
from terminal atoms to double or triple
bonds
• Certain atoms can be exceptions to
octet rule – H, Be, B, S, P, Xe
Drawing Lewis Diagrams
CF4
1 C × 4e- = 4e4 F × 7e- = 28e32e- 8e24e-
F
F C F
F
Drawing Lewis Diagrams
CO2
1 C × 4e- = 4e2 O × 6e- = 12e16e- 4e12e-
O C O
Polyatomic Ions
To find total # of valence e-:
 Add 1e- for each negative charge
 Subtract 1e- for each positive
charge
Place brackets around the ion and
label the charge
Polyatomic Ions
ClO41 Cl × 7e- = 7e4 O × 6e- = 24e31e+ 1e32e- 8e24e-
O
O Cl O
O
Octet Rule
F
F
 Hydrogen  2 valence e
F
B
F
 Boron & F
Beryllium
getF
6 & 4 valence e
S
H
N
O
O
H
respectively
F
 Expanded
octet
 more than 8
Very
unstable!!
F
F
valence e (e.g. S, P, Xe)
Exceptions:
-
-
-
Drawing Lewis Diagrams
BeCl2
1 Be × 2e- = 2e2 Cl × 7e- = 14e16e- 4e12e-
Cl Be Cl
Drawing Lewis Diagrams
SF6
1S× 6e- =+ 6e6F× 7e- = 42e48e
48 e-12 e36 e-
F
F
F S F
F
F
Resonance Structures



Molecules that can’t be correctly
represented by a single Lewis
diagram
Actual structure is an average of all
the possibilities
Show all possible structures
separated by double-headed arrows
C. Resonance Structures

SO3
O
O S O
O
O S O
O
O S O
Bond Polarity
• Most bonds are a
blend of ionic and
covalent
characteristics.
• Difference in
electronegativity
determines bond
type.
Bond Polarity
• Electronegativity
– Attraction an atom has for a shared pair of
electrons.
– higher e-neg atom  – lower e-neg atom +
Electronegativity
Difference
• If the difference in electronegativities is between:
– 1.7 to 4.0: Ionic
– Greater than 0.3 & less than 1.7: Polar Covalent
– 0.0 to 0.3: Non-Polar Covalent
The type of bond can usually be calculated by
finding the difference in electronegativity of
the two atoms that are bonded.
Compound
F2 or F-F
CF4
LiF or Li-F
Electronegativity
4.0 - 4.0 = 0 4.0 - 2.5 = 1.5 4.0 - 1.0 = 3.0
Difference
Non-polar (strong) Polar Ionic (nonType of Bond
covalent
covalent
covalent)
Bond Polarity
• Nonpolar Covalent Bond
– e- are shared equally
– symmetrical e- density
– usually identical atoms
• Ex: H2 or Cl2
Bond Polarity
• Polar Covalent Bond
– e- are shared unequally
– asymmetrical e- density
– results in partial charges (dipole)
• Ex: H2O
+


Polar Covalent Bonds: Unevenly
matched, but willing to share.
- water is a polar molecule because oxygen is more electronegative than hydrogen, and
therefore electrons are pulled closer to oxygen.
Ch. 8 – Molecular Structure
Molecular
Geometry
(p. 232 – 236)
I
II
III
VSEPR Theory
Valence Shell Electron Pair
Repulsion Theory
Electron pairs orient themselves in
order to minimize repulsive forces
VSEPR Theory
Types of e- Pairs
 Bonding pairs – form bonds
 Lone pairs – nonbonding e Total e- pairs– bonding + lone pairs
Lone pairs repel
more strongly than bonding
pairs!!!
A. VSEPR Theory
Lone pairs reduce the bond angle
between atoms
Bond Angle
Determining Molecular Shape
Draw the Lewis Diagram
Tally up e- pairs on central atom
(bonds + lone pairs)
 double/triple bonds = ONE pair
Shape is determined by the # of
bonding pairs and lone pairs
Common Molecular Shapes
2 total → Electronic Geometry = linear
2 bond
0 lone
BeH2
LINEAR
180°
Common Molecular Shapes
3 total → Electronic Geometry =
trigonal planar
3 bond
0 lone
BF3
TRIGONAL PLANAR
120°
Molecular Polarity
 Polar
molecule = one end slightly +
and one end slightly –
 Molecule with 2 poles = dipolar
molecule or dipole
Molecular Polarity
 Shape,
symmetry and bond
polarity determines molecular
polarity
 H – O bond is polar and water
is asymmetrical, so H2O is
polar
 C – Cl bond is polar, but CCl4
is symmetrical, so molecule is
nonpolar
Molecular Polarity
 Identify
each molecule as polar or
nonpolar
• O2
Nonpolar bonds → nonpolar
• CS2
Linear → nonpolar
• CF4
Tetrahedral → nonpolar
• H2O
Bent → polar