Chemical Bonding
Ms. Grobsky
What Determines the Type of Bonding in Any
Substance?
Why do Atoms Bond?
• The key to answering the first question are found in the
electronic structure of the atoms involved
• In general, bonding is the interaction between atoms
• There are two types of interactions between atoms:
▫ Energetically favored
 Electrons on one atom interacting with protons of another atom
▫ Energetically unfavorable
 Electrons on one atom interacting with electrons of another atom
 Protons on one atom interacting with protons of another atom
• A bond will form if the system can LOWER its total
energy in the process
Illustration of Bonding
Using Electronegativities to Determine
Type of Bond
• Electronegativity is a measure of an atom’s
attraction for another atom’s electrons
• It is a made-up scale that ranges from 0 to 4
▫ 0 = no electronegativity
▫ 4 = high electronegativity
• All electronegativity values are relative to fluorine,
the most electronegative element
• Generally, metals are electron givers and have low
electronegativities
• Nonmetals are electron takers and have high
electronegativities
• What do you think about the electronegativity of
noble gases?
Using Electronegativities to Determine
Type of Bond
• Can use the difference in electronegativities to determine
type of bond formed
▫ Ionic – electronegativity difference greater than 1.67
▫ Covalent – electronegativity difference less than 1.67
▫ Non-polar covalent – electronegativity differences less than
0.4
• In other words, the closer the elements are on the
periodic table, the similar their electronegativities
• More likely to form covalent bonds
• Elements that are farther away on the periodic table,
greater difference in electronegativity
• More likely to form ionic bonds
Where are Metals Located in the
Periodic Table?
• Metal are found to the left and centre of the
periodic
table!
H
Li
He
Be
B
C
N
O
F
Ne
Na Mg
Al
Si
P
S
Cl
Ar
Ni Cu Zn Ga Ge As
Se
Br
Kr
I
Xe
K
Ca
Sc
Ti
Rb
Sr
Y
Zr Nb Mo Tc Ru Rh Pd Ag Cd
In
Sn Sb
Te
Cs
Ba La
Hf Ta
Tl
Pb
Bi
Po
At Rn
?
?
?
?
?
Fr Ra Ac
V
Cr Mn Fe
W
Re Os
Co
Ir
Pt
Au Hg
Rf Db Sg Bh Hs Mt Ds Rg
?
?
Types of Bonds – Ionic Bond
• Ionic bonds are formed between a metal and
non-metal
• Electron(s) transferred from the metal to the
non-metal
▫ As a result, a cation and anion are produced
• The reason why an ionic bond is so stable is
because of the electrostatic attraction between
the cation and anion
Types of Bonds – Ionic Bonds
• Ionic compounds form huge, repeating 3-D crystal lattices
where the ions and electrons are located at fixed positions
• Strong interactions between ions have a profound effect on
melting points and solubilities
▫ Large melting points
▫ Solids at room temperature
Types of Bonds - Covalent Bonds
• Defined as a type of bond in which valence electrons are
shared between nuclei of two non-metal atoms
▫ Sharing is based on electronegativity of each atom in the
bond
• Bonds can be single, double, or triple depending on how
many electron pairs are shared:
• Single bond
• One pair of electrons shared
• Double bond
• Two pairs of electrons shared
• Triple bond
 Three pairs of electrons shared
• Because electrons are shared, physical properties of covalent
compounds vary wildly
There are 2 Types of Covalent Bonds!
• Polar covalent
▫ The electrons are shared, but one atom exerts a
greater attraction for the bonding electrons than
the other
 As a result, the electrons spend more time around
that atom and creates a partial negative charge
 The other atom now has a partial positive charge
• Non-polar covalent
▫ The electrons are evenly shared between the two
atoms
Example of a Polar Covalent Bond
How are Ionic and Covalent Bonds
Formed?
• Both of these types of bonds involve valence
electrons!
▫ As discussed previously, valence electrons are the
electrons found in the highest energy s- and p- orbitals
• Most elements want to obey the octet rule when
forming bonds
• Each atom has a TOTAL of 8 valence electrons around
it
•
Most important requirement for the formation of a
stable compound is that atoms achieve a noble gas
configuration (octet)
Bonding between Metal Atoms
• Atoms of metals are tightly packed together in a giant
lattice similar to the lattice in ionic compounds
• The outer electrons separate from their atoms and become
delocalized, creating a ‘sea of electrons’
▫ The atoms become positive ions and are attracted to these
electrons
• This attraction is called metallic bonding and is the reason
why the positive metal ions do not repel each other
Metal
atoms
Sea of free
electrons
Metallic Bonding and Its Effects on
Metal Properties
• Metal ions form a lattice which is more tightly
packed and denser than the lattices in ionic
compounds
▫ As a result, metals generally have a very high
melting and boiling point because metallic bonds
are very strong and therefore need a large amount
of energy is needed to break them
Metallic Bonding and Its Effects on
Metal Properties
• Metals are good conductors of:
▫ Heat
 The free electrons can take in heat
energy, which makes them move
faster
 They can then transfer the energy
throughout the lattice
▫ Electricity
 The free electrons can carry an
electrical charge
• Silver is the best conductor of
electricity and copper is the second
best
▫ Why do you think copper is used
instead of silver for electrical wires?
Metallic Bonding and Its Effects on Metal Properties
• Metals are usually tough, not brittle
• When a metal is hit, the layers of the lattice just slide over
each other
▫ The metallic bonds do not break because the electrons are free to
move
• This means that metals are:
▫ Malleable
 They can be bent and pressed into shape
▫ Ductile
 They can be drawn out into wires
Force
Illustrating Bonds - Lewis Dot
Structures
• Lewis Dot structures are also known as electron
dot diagrams
• These diagrams illustrate valence electrons and
subsequent bonding
Using Electronegativity Differences to
Draw Appropriate Lewis Dot Structures
• Check the electronegativity difference between the
elements to determine if electrons are transferred or
shared
• If the electronegativity difference > 1.67, the
reaction forms an IONIC BOND
• Draw the Lewis Dot structure for each atom
• Write its element symbol
• Draw dots to represent number of valence electrons
• Remove the electrons from the metal and add them
to the nonmetal
Lewis Dot Structure of Ionic
Compounds
• Write the charges of the ions formed and use
coefficients to show how many of each ion are
needed to balance the overall charge
+
2-
2Na , [ O ]
Ionic sodium oxide, Na2O
Lewis Dot Structure of Covalent
Compounds
• If the electronegativity difference < 1.67, then
the atoms will share electrons
• Unpaired (single) electrons will participate in
bonding
• Paired electrons will not participate in bonding
and are called lone pairs
Bonding and Lone Pairs
• Valence electrons are distributed
as shared or BOND PAIRS and
unshared or LONE PAIRS
••
H
Cl
•
•
••
shared or
bond pair
lone pair (LP)
This is called a LEWIS
structure!
Steps to Draw Lewis Dot Diagrams for
Covalent Compounds
• Add up number of valence electrons that can be used
• Add electrons for anions, subtract electrons for cations
• Place the least electronegative atom as the central atom
• Central atom is NEVER H because it is too small!
• Draw a single bond , -, between the central atom and each
surrounding atom
▫ Each bond takes 2 electrons!
• Add remaining electrons to form lone pairs, :, to each atom in
order to satisfy octet rule
▫ Note – H only needs 2 electrons to be satisfied!
• Check to make sure there are 8 electrons around
each atom EXCEPT H
Steps to Draw Lewis Dot Structures
• Count the number of electrons in your Lewis Dot
structure
• If you have more electrons in the drawing than you
needed, you must form double or triple bonds!
Practice!
• CH4
• F2
H
H C H
F F
H
Do atoms (except H or metals) have octets?
Carbon Dioxide, CO2
1. Central atom =
2. Valence electrons =
3. Form bonds.
C 4 eO 6 e- X 2 O’s = 12 eTotal: 16 valence electrons
This leaves 12 electrons (6 pair).
4. Place lone pairs on outer atoms.
5. Check to see that all atoms have 8 electrons around it except for H, which can
have 2.
Carbon Dioxide, CO2
C 4 eO 6 e- X 2 O’s = 12 eTotal: 16 valence electrons
How many are in the drawing?
6. There are too many electrons in our drawing. We
must form DOUBLE BONDS between C and O.
Instead of sharing only 1 pair, a double bond shares 2
pairs. So one pair is taken away from each atom and
replaced with another bond.
Double and even
triple bonds are
commonly
observed for C,
N, P, O, and S
H2CO
SO3
C2F4