Transcript Chemical Bonding

Chemical Bonding
Chemical Bond
• The forces that hold groups of atoms together
and make them function as a unit
• Bonding involves only the valence electrons
• There are 2 types of bonds:
– Ionic: Transfer of electrons from a metal and to a
nonmetal
– Covalent: Sharing of electrons between 2
nonmetals
– Note: When 2 metals bond an alloy is formed
• Electrons are transferred or shared to give
each atom a noble gas configuration (stable
octet)
– This is known as the octet rule
Lewis Diagrams
• Valence electrons involved in bonding can
be represented by Lewis dot diagrams
• A chemical symbol represents the nucleus
and the core electrons (not involved in
bonding).
• Dots around the symbol represent valence
electrons.
Drawing Lewis Diagrams
Cl
1. Write the element symbol.
2. Draw dots, one for each valence
electron
3. Dots should be spread over 4 sides
•
It does not matter what side the dots are
placed, but do not start to pair dots until there
is one on each side
The number of valence electrons is equal to
the group number. With one exception.
• Lewis diagrams for the first 20 elements
Ionic Bonding
Ionic Bonding
• Metals
– Electron donors
– Donate their valence electrons to become a
positive ion (cation)
• Nonmetals
– Electron acceptors
– Accept valence electrons to become a negative
ion (anion)
Ionic Bonding
Ionic Bonding
The two oppositely charged
ions are attracted to each
other by a force called an
ionic bond
Properties of Ionic
Compounds
Structure:
Crystalline solids
Melting point:
Generally high
Boiling Point:
Generally high
Electrical
Conductivity:
Solubility in
water:
Excellent conductors,
molten and aqueous
Generally soluble
NaCl Crystal Lattice
Ionic compounds form solids
at SATP.
Ionic compounds organize in
a characteristic crystal
lattice of alternating positive
and negative ions.
All lattices are arranged so
that each ion has the
greatest possible number of
oppositely charged ions close
by, while keeping similarly
charged ions as far away as
possible
Representing Ionic Compounds
Lewis Diagrams
• Formation of sodium chloride:

]


 Na+ [ Cl

Cl

Na  +



Lewis Structures for Ionic
Compounds
• O•
••
2+ ••
Ba
O
••
2-
••
Ba•
••
••
Ba and O
•
BaO
••
••
••
2 Cl
• Cl
••
MgCl2
••
-
••
Mg
••
••
Mg •
2+
••
Mg and Cl
•
• Cl
••
Representing Ionic Compounds
Criss-Cross Method
For monatomic ions:
Take the absolute value of the ionic charge for the
cation and make it the subscript for the anion and vice
versa.
Example: Al3+ and ClThe 3 becomes the subscript for the chloride ion and
the 1 becomes understood for aluminum.
Forming aluminum chloride: AlCl3
Representing Ionic Compounds
Criss-Cross Method
For polyatomic ions:
Additional step of including brackets around the
polyatomic ion if it has a subscript other than one.
Example: Mg2+ and OHThe 2 becomes the subscript for the hydroxide ion, but
brackets are needed to indicate 2 of each the O and the H.
The 1 becomes the understood subscript for Mg.
Forming magnesium hydroxide: Mg(OH)2
Polyatomic Ions
NICK the CAMEL ate a CLAM for SUPPER in PHOENIX
• Underlined letter represents the symbol of the
element.
• The consonants represent the number of oxygen
• The vowels represent the negative charge.
Eg.
Underlined letter= N
Number of consanants= 3
Number of vowels= 1
NO3-
Nitrate
represents oxygens
represents charge
Covalent Bonding
Lone pairs, valence electrons
not involved in covalent bond
• Formation of hydrogen chloride:
H Cl






Cl

H +


 H - Cl



Covalent bond, shared electrons
Structural Formula: H-Cl
(lone pairs are not drawn)
Lewis Structures
H  +  H  H H or H H




Cl
Cl
 
Structural Formula: Cl-Cl

or

Cl
Cl






Cl
+
Cl



Cl2:




H2:
Double and Triple Bonds
• Atoms can share 4 electrons to form a double bond
or 6 electrons to form a triple bond.


N 2:
N N

O
=O



O 2:
• The number of shared electron pairs (covalent
bonds) that an atom can form is the
bonding capacity.
Multiple Covalent Bonds
•
•
••
••
N N
•
•
•
N N
••
••
•
N N
••
•
••
•N
•
N•
••
•
••
•
Multiple Covalent Bonds
•O
••
••
•
•
•
••
•
••
••
O C O
•
•
•
••
•
••
O C O
••
• C•
•
••
••
O•
•
••
••
••
••
O C O
••
•
••
•
Drawing Lewis Structures
1. Arrange the element symbols.
•
Central atoms are generally those with the highest bonding capacity.
•
Carbon atoms are always central atoms
•
Hydrogen atoms are always peripheral atoms
2. Add up the number of valence electrons from all atoms.
•
Add one electron for each negative charge and subtract one for each
positive charge.
3. Draw a skeleton structure with atoms attached by single
bonds.
4. Complete the octets of peripheral atoms.
5. Place extra electrons on the central atom.
6. If the central atom doesn’t have an octet, try forming
multiple bonds by moving lone pairs.
Structural Formula
• From the Lewis structure, remove dots
representing lone pairs
• Replace bond dots with a dash
Draw Lewis structures and the structural formula for:
CH4:



H N H

H
H

H C
H

H

or
or
or
or



H O
H


NH3:


H2O:
H 
F

HF:


H F


H O
H


H N H
H
H
H C H
H
Coordinate Covalent Bonds
A covalent bond in which both of the shared electrons come
from the same atom.
E.g. NH3 (ammonia) and H+ (hydrogen ion) to form NH4
(ammonium)
Drawing Lewis Structures





14 ve’s


H O
Cl





HOCl
24 ve’s

COCl2
O


Cl
C
Cl


CH3OH
26 ve’s
14 ve’s

O
Cl
O




ClO3



O
H

H C O
H

H