Bonding
Chapter 7
Electrons
Valence Electrons = electrons in the highest occupied
energy level
# Valence Electrons = _______________ number
What abt transition elements?
- it’s not that simple: they don’t always
bond with the same amt every time!
Electron Dot Structures: diagrams that show how many eare available to bond:
E- Dot Structure
Li Be B C
N O
F Ne
The OCTET Rule
The OCTET RULE:
In forming a compound, atoms tend to achieve the
electron configuration of a noble gas
Aka…they want ______ electrons in their outer energy
level: ns2np6
Ex: Neon: 1s22s22p6
**There are exceptions, but in general, this rule holds**
OCTET Rule (cont…)
The OCTET RULE:
Metals  lose e- to achieve an octet in a lower energy
level
Nonmetals  gain to get 8 if combining with a metal
- Or, they will share with other nonmetals
Metal cation + non-metal anion = IONIC BOND (ch 7)
2 non-metals = COVALENT BOND (ch 8)
2 metals = METALLIC BOND (ch 7)
Ions
Formation of CATIONS:
______________ electrons
Most common cations: formed by the loss of valence
electrons by metal atoms
Predict the ions:
K
Ca
Sr
Ra
Ba
Ions
Formation of ANIONS:
_______________ of electrons
Formed by: Nonmetals & Metalloids
Predict the Anion:
Cl
Br
S
P
N
O
F
Ions: Apply Your Knowledge
What is the ionic charge on the following:
NaCl
CaO
CsCl
Ionic Bonds & Ionic Compounds
Ionic compounds = compounds composed of cations
& anions
As a whole, the compounds are electrically neutral:
Na___ + Cl___ 
** The total positive charge of the cations is
_____________ to the total negative charge of the anion
**
Ionic Bonds
Ionic bond – refers to the
electrostatic force that holds ionic
compounds together!
If the difference between the
electronegativities of the atoms
is greater than 2.0, the bond is
IONIC
Metal + Nonmetal
(Cation + Anion)
Chemical Formulas
Chemical Formula = shows the kinds of atoms and
numbers of atoms in one molecule of a substance
i.e. NaCl
Formula Unit = lowest whole number ratio of ions in an
ionic compound
NaCl ( 1: 1 ration ; 1 Na1+ to Cl1- )
Example: What is the Formula Unit & Chemical Formula
for a reaction between magnesium ions & chlorine
ions?
What about aluminum ions & bromine ions?
Predicting Formulas
Examples:
1. Magnesium & Nitrogen
1. Calcium & Chlorine
1. Calcium & Fluorine
Stop: Ionic Bonding
Worksheet & Activity
Properties of Ionic Compounds
 Most are crystalline solids at room temp.
arranged in repeating three-dimensional patterns
very stable structure (large attractive forces btw
molecules)
… this results in…
 Generally have high melting points
NaCl – melts around 800 C
Which pairs of elements are
most likely to form ionic
compounds?
a. Cl, Br
b. Li, Cl
c. K, He
d. I, Na
Metallic Bonds
Metallic Bonds = attraction of free floating valence
electrons with close packed metal cations
Think…sea of electrons
- Free floating and mobile
- Allows for conductivity, ductility, and
malleability
- shifting & rearrangement of both
cations and electrons possible
Arranged in a compact and orderly structure
Alloys
 Mixtures!!
Two or more elements – at least 1 is a metal
Why are alloys important?
 their properties are often superior to the
elements that make up the alloys
Most important alloys todays are steels
STOP!! Questions??
Molecular Compounds (ch 8)
 COVALENT BONDS




Sharing of e- to hold atoms together
Forms between 2 nonmetals
Neutral
Makes molecules
Bond Dissociation Energy: the amt of energy
required to break a covalent bond
a large bond dissociation energy corresponds
to a strong covalent bond
Double & Triple bonds are strong than singles
= _______________ bond dissociation energies
Molecular Compounds (ch 8)
Chemical Formula = Molecular
Formula
Molecule = neutral group of atoms
joined together by covalent bonds
How are molecular compounds
different that ionic compounds?
Ionic compounds = salts
(crystalline solids w/ high mp)
Molecular compounds =
chemically linked together
Molecular Compounds (ch 8)
Molecular Formula  NOT lowest whole
number, but the actual number of atoms of
each element present
Properties of Covalent Bonds
2 or more non-metals
Not good conductors
Lower mp & bp than ionic compounds
Diatomic Molecules
Molecule of 2 identical atoms

Covalent Bonding
**Atoms in a covalent bond obey the OCTET RULE
 electron sharing occurs so that atoms
electron config of noble gases (8 valence e-)
- Single Bond
 Share one pair of electrons (one e- from each atom)
Ex. H – H
- Double Bond
 Share two pairs of electrons (2 e- from each atom)
Ex. O=O
- Triple Bond
 Share three pairs of electrons (3 e- from each atom)
Ex. N≡N
- Coordinate Bond
 One atom contributes both bonding atoms
Ex
:C: + :Ö:  :C≡O: (often drawn :C=O: )
attain the
Structures
23
 Dot structures can be made for covalent bonds
 Place shared electrons between the 2 symbols either as
pair of dots or dashes
 Ex
H-H
or
:Ö=Ö:
or
:N≡N:
 Each dash represents a single covalent bond
 Unshared pairs / lone pairs / non-bonding pairs =
valence e- pairs not involved in covalent bond
 Structural formulas use dashes, but also show spatial
arrangement of atoms
 Draw the Lewis dot diagrams for all the atoms in CH4
 Draw 1 carbon diagram
 Draw 4 hydrogen diagrams
How Covalent Bonds Form
Bonding in Methane
C
H H H H
Bonding in Methane
C
H H H H
Bonding in Methane
C
H H H H
Bonding in Methane
H
C
H H H
Bonding in Methane
H
C
H H H
Bonding in Methane
H
C
H H H
Bonding in Methane
H
C H
H H
Bonding in Methane
H
C H
H H
Bonding in Methane
H
C H
H H
Bonding in Methane
H
C H
H H
Bonding in Methane
H
C H
H H
Bonding in Methane
H
H C H
H
Draw the following compounds:
PCl3
H 2O
SF2
CCl4
 Two electrons are shared between two atoms
 1 electron pair
 Represented by 1 line between the atoms
Single Bond
F
Let’s Make F2
F
F F
Let’s Make F2
F F
Let’s Make F2
 Four electrons are shared between two atoms
 2 electron pairs
 Represented by 2 lines between the atoms
Double Bond
O
Let’s Make O2
O
O O
Let’s Make O2
O O
Let’s Make O2
O O
Let’s Make O2
Draw the Lewis dot diagram for
O3
CO2
CF2S
Double bonds Try It!
 Six electrons are shared between two atoms
 3 electron pairs
 Represented by 3 lines between the atoms
Triple Bond
N
Let’s Make N2
N
N N
Let’s Make N2
N N
Let’s Make N2
Let’s Make N2
N N
Let’s Make N2
N N
Let’s Make N2
Draw the Lewis Dot Diagram for
HCN
C2H2
Triple bonds Try It!
General Steps to Lewis Structures
Sum
◦Polyatomic ion?
Divide
Determine
◦Always ____ , never ____________ , _______ “bully”
Make
Complete
◦Except
Give
◦Not enough pairs?
◦Ionic compound b/c ____________ is central?
Making Sense
General Steps to Lewis Structures
Sum total val e◦Polyatomic ion? Add e- if anion or Subtract e- if cation
Divide
by 2 for # of pairs
Determine central atom
◦Always Carbon
Make
single bonds
Complete octet on external atoms
◦Except Hydrogen
Give left over pairs to the central atom for octet
◦Not enough pairs? Make double or triple bonds
◦Ionic compound b/c metal is central? No octet needed
Making Sense
Most Common Bonding Patterns
After following all of the rules, you need to make sure the
common bonding patterns are satisfied, if possible.
Otherwise, your Lewis Structure is probably incorrect.
Polyatomic Ions
 Even though they have an overall charge, polyatomic
ions are covalently bonded together
 Overall charge comes from extra electrons either
gained or lost in order to achieve an octet for all
constituent atoms
 Ex: SO32O
S O
O
 Use brackets to denote total charge
58
 Covalently bonded
 With additional or removed electrons
 Cations lose and anions gain in e- count
NO31-
SO32-
PO43-
CN1-
Polyatomic Ions
Resonance structures
60
 Occur when its possible to draw two or more valid
electron dot structures for a molecule
 Originally thought that electrons would flip back and forth
(resonate) between the options
 Actual bonding is a hybrid (mixture) of the options
 All valid structures are separated by double headed arrow
 Ex: O3 (ozone)
:Ö = Ö – Ö:
˙˙
↔ :Ö - Ö = Ö:
˙˙
Exceptions to Octet Rule
 If total # of valence e- for compound is odd, the octet rule
cannot be satisfied
 Nitrogen dioxide
 Phosphorus pentachloride
 Sulfur hexafluoride
 *Boron trifluoride
 readily reacts with NH3
61
STOP
We have reached the end of Part 2.
Any Questions?
VSEPR Theory
Valence Shell Electron-Pair Repulsion
Theory
States: the repulsion between electron
pairs causes molecular shapes to
adjust so that the valence-electron
pairs stay as far apart as possible
**Non bonding pairs are important
because they will repel other electron
pairs
- they are not being shared, so the
nucleus holds on tighter, pushing
everything else out of the way
Common Molecular shapes
# bonds
(σ)
# lone pairs
Shape
Bond < s
2
0
Linear
180
2
2
Bent
105
3
0
Trigonal planar
120
3
1
Pyramidal
107
3
2
T shaped
90/180
4
0
tetrahedral
109.5
4
2
Square planar
90/180
5
0
Trigonal bipyramidal
120/90
6
0
octahedral
90
Determining Molecular Shape:
VSEPR Rules
1. Draw the Lewis structure for the molecule or ion
2. Count the total number of regions of high e- density
 This includes BOTH bonds & unbonded pairs
- Double & Triple bonds count as: ONE REGION
- An unbonded electron or pair of electrons
counts as: ONE REGION
- For resonance structures, you may use either
structure
 Examples
 Test My Knowledge
Determining Molecular Shape:
VSEPR Rules Continued…
3. Identify the most stable arrangement of regions
of high e- density as ONE of the molecular
shapes
Examples
4. Then, using the determined starting shape,
determine final shape based on ATOMS and not
the unpaired electrons
We Need Practice…
 Practice:
 CO2
 Linear (180)
 H2O
 Bent (105)
 NH3
 Pyramid (107)
 SF6
 Octahedral (90)
 ICl3
 T-shaped (90/180)
BF3
trigonal planar (120)
CH4
tetrahedral (109.5)
PCl5
trigonal bipyramidal (90/120)
XeF4
square planar (90/180)
Polar Bonds and Molecules
68
 Think back to electronegativity…
 The amount of pull one atom has over the shared
electrons in a bond (please consult new table with
values)
 If the electronegativity values are the same or really
close (ex: two identical atoms) then the bond is said to
be non-polar covalent because they are sharing the
electrons equally
 However, if the atoms are sharing unequally due to one
atom having a higher electronegativity value, the bonds
are said to be polar covalent
 The more electronegative atom pulls the electrons
closer to its own nucleus, thus forming an uneven
electron cloud
Bond Types
Table 8.3 on page 238 describes the
bond types based upon
electronegativity value differences
 0.0-0.4  non-polar covalent
 0.4-2.0  polar covalent
 2.0-4.0  ionic bond
69
“Slight” charges
 The more electronegative
atom has a slightly negative
(δ-) charge
 The less electronegative atom
has a slightly positive (δ+)
charge
 The slight charge is denoted by
the lower case Greek letter delta
(δ) above the element symbol
 An arrow with a vertical line (+→)
points away from the δ+ atom,
towards the δ- atom above the
bond dash to show the emovement in a bond
70
Polar Molecules
 A molecule can only be polar if its constituent
bonds are polar.
 A symmetrical molecule with polar bonds is
non-polar since the bonds will negate each
other.
 If not symmetrical, there will be a general δ+
and δ- region to the molecule.
 A polar molecule with 2 polar regions is called
a dipole.
 Use your table of electronegativity to decide if
each of your molecular geometry compounds
have polar or non-polar bonds and if they are
polar or non-polar molecules.
71
Polar v. Nonpolar Molecules