November 25, 2014
Chemical Bonding: Ch 7 and 8
Chemical Bonding
A chemical bond is a mutual electrical attraction
between the nuclei and valence electrons of different
atoms that binds the atoms together.
Why are most atoms bound together?
Most atoms are at relatively high potential energy by
themselves and are more stable when combined with
other elements!
http://www.youtube.com/watch?v=QXT4OVM4vXI
http://www.youtube.com/watch?v=UR4eG60jjQQ
http://www.youtube.com/watch?v=QqjcCvzWwww
We will be discussing ionic vs. covalent bonding.
Chapter 7- Ionic Bonding
Chapter 8- Covalent Bonding
Chemical Bonding-Vocabulary
Chemical Bonding-Vocabulary
A chemical bond
Chemical Formula
covalent
ionic
-Attractive force between atoms or ions that binds them
together as a unit.
-bonds form in order to:
-decrease potential energy (PE)
-increase stability
Chemical Bonding-Vocabulary
Formula Unit
NaCl
Binary Compound
NaCl
H2O
Chemical Bonding-Vocabulary
Ion
Compound
2 elements
Molecular Formula
more than 2 elements
Ternary Compound
NaNO3
1 atom
Monoatomic Ion
Na+
2 or more atoms
Polyatomic Ion
NO3-
November 25, 2014
B) Types of Bonds
Ionic
Covalent
Metallic
electrons
transferred
Bond Formation
from metal to
nonmetal
electrons
shared
between two
nonmetals
electrons are
delocalized
among metal
atoms.
Type of Structure crystal lattice
true molecules "electron sea"
Physical State
solid
liquid or gas
solid
Melting Point
high
low
very high
Solubility in
Water
yes
usually not
no
Electrical
Conductivity
yes (solution
or liquid)
no
yes (any form)
odorous
malleable,
ductile,
lustrous
Other Properties
Example
NaCl, BaCl2,
CaO, etc.
1) Ionic Bonding
-Transfer of electrons from one atom to
another.
Metal + Nonmetal
Ionic Compound
Valence electrons are usually the only electrons used in
chemical bonds.
H2O, NO, NO2 Al, Fe, etc.
Example: Sodium Chloride
The chemical formula: NaCl
because Na+1 and Cl-1
Atoms of metals tend to lose their valence electrons, leaving
a complete octet in the next-lowest energy level.
Atoms of some non-metals tend to gain electrons or to
share electrons with another nonmetal to achieve a complete
octet.
Practice Problems:
Determine the charge of the ions and the final formula unit
of the following:
1. potassium and iodine
2. aluminum and oxygen
3. magnesium and chlorine
Steps:
1. Look at the charge
2. Determine the ratio
of charges
3. Criss-Cross the
charges and simplify
to the lowest ratio.
Nomenclature of Ionic Compounds
1. Keep the name of the cation the same and change the
name of the anion to an -ide ending.
a) NaCl= sodium chloride
b) CaCl2= calcium chloride
c) BaF2=
d) KCl=
e) MgO=
November 25, 2014
Coordination patterns and packing depends on the size of the
ions. Remember the trends from the Periodic Table!
http://www.youtube.com/watch?v=KNgRBqj9FS8
Body Centered Cubic
Every atom has 8 neighbors
Ex: (Na, K, Fe)
http://www.youtube.com/watch?v=ZVqocQLAEr0
Face Centered Cubic
Every atom has 12 neighbors
Ex: (Ag, Au, Al, Pb)
Oxidation Numbers (Honors)
-Can figure out from the formula
-Unpaired electrons in d orbitals
Fe: (+2 and +3)
http://www.youtube.com/watch?v=Rm-i1c7zr6Q&feature=related
Hexagonal close-packed
Every atom ALSO has 12 neighbors but is
pattern is different.
Ex: (Zn, Mg, and Cd)
Section 7.3 Bonding in Metals
*Similar to structure of ionic!
The valence electrons of metal atoms can be modeled as a
sea of electrons.
The valence electrons are mobile and can drift freely from
one part of the metal to another.
The electrons can drift because of the vacant d orbitals just
below their highest energy level.
These orbitals can overlap and the overlapping is what causes
the electrons to move about freely.
http://www.drkstreet.com/resources/metallic-bonding-animation.swf
Bonding in Metals
Metallic bonds:
consist of the attraction of the free-floating valence
electrons for the positively charged metal ions.
November 25, 2014
Crystalline Structure of Metals
Alloys:
Mixtures composed of two or more
elements, at least one of which is a metal.
*Alloys are important because their
properties are often superior to those of
their component elements.
Types of Alloys
Substitutional- atoms
of different
components are of
similar size.
Ex: Brass
3) Covalent Bonding
Electrons shared between nonmetals
Interstitial- solute
atoms occupy "gaps"
because they are
different sizes.
Ex: Carbon for steel
Potential Energy Diagram
4.
3.
2.
1.
B(r)illiant
Harvard
Nerds
Mr. BrINClHOF
Often
Find
C(l)hemistry
Interesting
Most elements can form diatomics, HOWEVER,
only at very high temperatures.
Kinetic Theory (Random movement)
1. The separated atoms do not affect each other.
2. Potential energy decreases as the atoms are drawn together.
3. Potential energy is at a minimum when attractive forces are
balanced by repulsive forces.
4. Potential increases when repulsion between like charges outweighs
attraction between opposite charges.
The "7" exist free at normal conditions.
November 25, 2014
1. Single Covalent Bond
The Nature of Covalent Bonding
In covalent bonds, electron sharing usually occurs so that
atoms attain the electron configurations of noble gases.
Why would atoms want to do that?
We will be looking at three different types of covalent
bonds:
1. Single Covalent
2. Double and Triple Covalent
3. Coordinate Covalent
2. Double and Triple Covalent Bonds
Atoms form double or triple covalent bonds if they can attain
a noble gas structure by sharing two pairs or three pairs of
electrons.
Double covalent bond: A bond that involves two shared
pairs of electrons
Triple covalent bond: A bond formed by sharing three pairs
of electrons.
Electron-Dot Notation (Lewis Dot Structures)
Why the exceptions??
Incomplete Octet: (Group 3A)
Group 3A: Boron 1s2 2s2 2p1 (only 3 valence electrons)
Expanded Octet: (3rd Period)
Beyond the 3rd period with the "d" orbitals
Sulfur: [Ne]3s2 3p4 (6 valence electrons)
Odd-Electron Molecules (Radicals)
(NO and NO2 )
Structural formula:
Represents the covalent bonds by dashes and shows the
arrangement of covalently bonded atoms. One dash represents two
electrons.
Unshared Pair:
A pair of valence electrons that are not shared between atoms.
(lone pair or nonbonding pair).
3. Coordinate Covalent Bonds
Coordinate Covalent Bond:
A covalent bond in which one atom contributes both
bonding electrons.
In a coordinate covalent bond, the shared electron pair
comes from one of the bonding atoms.
Lewis Structures-Steps
1. Find the total # of valence electrons
2. Arrange atoms-singular atom is usually in the middle
(usually carbon)
3. Form bonds between atoms (2 electrons)
4. Distribute remaining electrons to give each atom an
octet.
5. If there are not enough electrons to go around, you
will have to make double or triple bonds.
· Don't forget the exceptions.
· Boron and Aluminum only need 6
· Sulfur and Phosphorus can have an expanded octet.
November 25, 2014
Polyatomic Ions
Lewis Structures-Practice
NH3
-a tightly bound group of atoms that have a positive or
negative charge and behaves as a unit.
The charge is on the whole ion but the individual
components can be held together covalently.
CO2
Example is the ammonium ion (NH4 + )
Polyatomic Ions-Practice
Exceptions to Lewis Structures
1. Odd number of valence electrons
*Central atom will not have an octet
*Examples: ClO2 and NO
ClO4-
2. Central atom with less than eight electrons
*These compounds tend to be very reactive
*Examples: BF3, BH3
more than eight electrons
*Most common exception
*Expanded octet
*extra electrons fill in the empty d-sublevel
*Examples: SF6, XeF4
3. Central atom with
NH4+
Resonance
Resonance-Practice
O3
SO3
November 25, 2014
Bond Dissociation Energies
Nomenclature (Intro)
Bond Dissociation Energy:
The energy required to break the bond between two covalently
bonded atoms.
The larger the bond dissociation energy, the stronger
the bond
Example: C2H2 vs. C2H6
Nomenclature (Intro)
Nomenclature (Intro)
Name the following binary molecular compounds:
1. SF6 is sulfur hexafluoride
2. CO
3. CO2
4. N2 O
5. Cl2 O8
Resonance
(Honors Only)
Formal Charge is used to determine the best overall
resonance structure.
Handout