Download Lone pairs

Chemical Bonding:
Bonding Theory and Lewis
Formulas
Periodicity

There are certain trends within the periodic
table which affect reactivity and the ability to
form bonds.
Periodic Law
When elements are arranged in order of
increasing atomic #, elements with similar
properties appear at regular intervals.
250
Atomic Radius (pm)

Cl
200
F
150
Br
100
50
0
0
5
10
Atomic Numbe r
15
20
Chemical Reactivity
Families

1
2
3
4
5
6
7
Similar valence e- within a group result in
similar chemical properties
Chemical Reactivity



1
2
3
4
5
6
7


Alkali Metals
Alkaline Earth Metals
Transition Metals
Halogens
Noble Gases
Properties that affect Reactivity

Atomic Radius

size of atom
Atomic Radius
250
K
Atomic Radius (pm)
Halogens
200
Na
Li
150
100
Ar
Ne
50
Noble Gases
0
0
5
10
Atomic Number
15
20
Atomic Radius
Atomic radius increase as you move top to bottom.
Atomic radius increase as you move right to left.
1
2
3
4
5
6
7
Ionization Energy
Again
Ionization
whenenergy
you compare
is the amount
the ionization
of energy
energy
neededoftoelements
turn a stable
thereelement
is a trend..
into an ion.
He
1st Ionization Energy (kJ)
2500
Ne
2000
Ar
1500
1000
500
Li
Na
K
0
0
5
10
Atomic Number
15
20
Ionization Energy
Ionization energy increases as you move right and
down.
1
2
3
4
5
6
7
Ionization Energy

Why are atomic radius and ionization energy
opposite?

In small atoms, the negative electrons (e-)
are closer to the positive nucleus where the
attraction is stronger.
50 mμ
20 mμ
Small Atoms
Large Atoms
Valence Electrons, Valence Energy
Levels & Valence Orbitals

Periodic trends are related to the number of
valence electrons an element has.

Valance electrons are those electrons
occupying the highest energy level of an atom
(the outside shell).
How Many Valence Electrons?
First Level: 2eSecond Level: 8eThird Level: 8e-
***You can tell how many valence electrons are in each orbital
by counting the number of elements on each row of the periodic
table.***
 Ex.
The magnesium atom has 12 protons and
12 electrons.

The maximum number of electrons in the first
energy level is 2 e-
 That
leave 10 electrons left to place.
 The
next energy level cam
only hold 8 electrons.
 That
only leaves 2 electrons
that are TRUE valence
electrons
Orbital: a region in space in which an electron
with a given energy is likely to be found.
 There are four valence orbitals within the
valence energy level of an atom (1s and 3p’s)
 There are few exceptions to this rule

Hydrogen
Helium
Electrons will occupy all valence orbitals
before forming electron pairs. Empty bus
seat rule
 Normally a maximum of 8 electrons may
occupy a valence energy level. This is

known as the octet rule.

urinaltest
2 electrons in
first energy
level.
You place the
other 4
electrons in
each of the
four orbital.
\Leaving
you with 4
more electrons to
place.
Start at the top
and go
clockwise.
Electron Dot Diagrams (Lewis Symbols)

Electron dot diagrams can represent atoms
(neutral) or ions (charged).

ONLY show the atom’s valence electrons! –
These are the only electrons involved in a chemical
reaction (the electrons in the outer most ring!)
Dot Diagram Trends
Follow these steps:
1) Write the atomic symbol for the atom. This
symbol represents the nucleus and the core
electrons that do not participate in the
chemical bonding.
2) Dots () represent the electrons in the valence
energy level of the atom. Arrange these dots
around the atomic symbol.
3) One dot must be placed in each of the four
orbitals before any electron pairing occurs.
4) Begin with the fifth electron to make lone
pairs. (if you have to)
5) There is a maximum of 8 electrons that can
be drawn.
Lets try some…

Calcium
Ca

Oxygen
O

Bromine

Carbon
Br
C
Bonding Electrons versus Lone Pairs
Bonding electrons are unpaired electrons
that are involved in bond formation.
 Paired electrons are called lone pairs and are
generally not involved in bond formation.

Bond Types

There are 3 types of bonds that can be
formed

These are determined by which elements
combine
Types of Bonds
1. IONIC
2. COVALENT
Bond
Formation
e- are transferred from
metal to nonmetal
e- are shared between
two nonmetals
Type of
Structure
crystal lattice
true molecules
Physical
State
solid
liquid or gas
Melting
Point
high
low
Solubility in
Water
yes
usually not
Electrical
Conductivity
yes
(solution or liquid)
no
Other
Properties
odorous
Metallic Bonding




2 metals share electrons but no chemical
reaction occurs
Valence electrons are free to move about
between the atoms
Positive ions surrounded by a “sea” of mobile
electrons
Allows metals to be formed into any shape
1. Ionic Bonding

A complete transfer of electrons occurs in an
ionic bond.
Valence Electrons
1. Ionic Bonding
1
7
8
8
2. Covalent Bonds
2. Covalent Bonding

Results in a mutual sharing of electrons
between the two non-metals.

This Sharing can be:
 Equal
= nonpolar covalent
 Unequal = polar covalent
Nonpolar Covalent Bond
 e-
are shared equally between both
nucleus.
 Electron “Cloud” is symmetrical.
Polar Covalent Bond
 e-
are shared unequally
 asymmetrical e- density
 results in partial charges (dipole)
+


Polar Covalent Bond
Example:
H2O(l)
H
+

O
H

3.4 - 2.2 = 1.2
POLAR
Bond Polarity

Most bonds are a
blend of ionic and
covalent
characteristics.

Difference in
electronegativity
determines bond
type.
Electronegativity
Electronegativity:
 The attraction an atom has for a shared pair
of electrons….aka….the strength an atom
has to hold onto or take electrons.
Trends in Electronegativity
Name That Bond!!!!
Electronegativity
Difference
Between Two
Atoms.
Type of Bond
Between the
Atoms
 1.7
Ionic
0.4 < 1.7
Polar covalent
< 0.4
Nonpolar
covalent
Descriptions of
Electrons in the
Bond
Transfer of
electrons
between metal
and nonmetal
Electrons shared
unequally
between unlike
atoms
Electrons shared
equally between
identical
Bond Polarity
Non-Polar
(+)
Ionic
(-)
Polar
Bond Polarity
Examples:

Cl2
3.2-3.2=0.0
Nonpolar

HCl
2.2-3.2=1.0
Polar

NaCl 0.9-3.2=2.3
Ionic

Covalent bonds are classified as:
1. single (sigma bond),
2.
double (1sigma and 1pi bond),
or 3. triple bonds (1sigma and 2 pi bonds)
depending on the number of electrons shared
between the two nuclei.
Electron Dot Diagrams for Ionic Compounds

Electrons are transferred from the metal to the
nonmetal.

results in a net negative charge for the nonmetal.

And a net positive charge for the metal.
Electron Dot Diagrams for Ionic Compounds
Drawing
Na Cl
Drawing
Mg
Cl
Cl
[Mg]2+
-[Na]+ [ [Cl
]
Cl ]
Lewis Dot
Diagrams
For Molecular
Compounds
Octet Rule

Remember…

Most atoms form bonds in order to have 8
valence electrons.
Octet Rule

Exceptions:

Hydrogen  2 valence e-

Helium  2 Valence e-
HOH
Drawing Lewis Diagrams
1. Find total # of valence e-.
2. Arrange atoms - singular atom is usually in
the middle.
3. Form bonds between atoms (2 e-).
4. Distribute remaining e- to give each atom an
octet (recall exceptions).
5. If there aren’t enough e- to go around, form
double or triple bonds.
Drawing Lewis Diagrams

CF4
1 C × 4e- = 4e4 F × 7e- = 28e32e- 8e24e-
How many bonds can
Carbon (C) form?
F
F C F
F
Drawing Lewis Diagrams
BeCl2
1 Be × 2e- = 2e2 Cl × 7e- = 14e16e- 4e12e
How many bonds can
Beryllium (Be) form?
Cl Be Cl
Drawing Lewis Diagrams
CO2
1 C × 4e- = 4e2 O × 6e- = 12e16e- 4e12e
How many bonds can
Carbon (C) form?
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-
How many bonds can
Carbon (C) form?
O
O Cl O
O
Polyatomic Ions

NH4+
1 N × 5e- = 5e4 H × 1e- = 4e9e- 1e8e- 6e2e-
How many bonds can
Nitrogen (N) form?
H
H N H
H
Resonance Structures



Molecules that can’t be correctly represented by
a single Lewis diagram.
Actual structure is an average of all the
possibilities.
Show possible structures separated by a doubleheaded arrow.
1.
2.
Resonance Structures
 SO3
O
O S O
1.
3.
2.
O
O S O
O
O S O
Don’t Panic!!!!



Lewis dot diagrams can take some time to figure
out.
If you cannot immediately determine the right
configuration, take a deep breath and try again.
Think of these as puzzles, keep on working with
the pieces until they fit together
You Want What?
So what is the difference between a Lewis dot
diagram?
 A chemical formula?
 And a structural diagram?

CH4
Lewis Dot Diagram
___________________
Chemical Formula
___________________
Structural Diagram
___________________
3.3 Molecular Shapes and Dipoles
I
II
III
VSEPR Theory
 Valence Shell Electron Pair Repulsion Theory
 Stereochemistry: The study of 3-D spatial configuration
 Electron pairs (bond pairs and lone pairs) orient themselves
in order to minimize repulsive forces.
VSEPR Theory
In non science speak:
The interaction of electrons is what forms bonds.
Electrons have a negative charge, so they want to be AS FAR AS POSSIBLE AWAY FROM EACH
OTHER.
VSEPR Theory
 Types of e- Pairs
 Bonding pairs - form bonds
 Lone pairs - nonbonding e-
Lone pairs repel
more strongly than
bonding pairs!!!
VSEPR Theory
 Lone pairs reduce the bond angle between atoms.
Bond Angle
Determining Molecular Shape
 Draw the Lewis Diagram.
 Count how many things attached to central atom
 double/triple bonds = ONE
 Lone pairs = ONE
 Shape is determined by the # of bonding
pairs and lone pairs.
Know the base shapes
& their bond angles as well as the derived
shapes
Base vs. Derived shapes
 Base Shapes: AKA electron pair geometry. Include all lone
pairs as a bond
 Derived Shapes: AKA molecular geometry. Only worry
about actual bonds.
1. Base Shape: Linear (AX2)
2 total electron pairs
2 bond pairs
0 lone pairs
180o bond angle
BeH2
2. Base Shape: Trigonal Planar (AX3)
3 total electron pairs
3 bond pairs
0 lone pairs
120o bond angle
BH3
3. Derived Shape: Bent/Angular (AX2E1)
3 total electron pairs
2 bond pairs
1 lone pair
< 120o bond angle
SO2
4. Base Shape: Tetrahedral (AX4)
4 total electron pairs
4 bond pairs
0 lone pairs
109.5o Bond angle
CH4
5. Derived Shape: Trigonal Pyramidal (AX3E1)
4 total electron pairs
3 bond pairs
1 lone pair
< 109.5o bond angle
NH3
6. Derived Shape: Bent/angular/v-shaped
(AX2E2)
4 total electron pairs
2 bond pairs
2 lone pairs
<109.5o bond angle
H2O
7. Derived Shape: Linear (AX1E3)
 4 electron pairs
 1 bond pair
 3 lone pairs
 180o bond angle
HF
8. Derived Shape: Linear (AX2E3)
 5 total electron pairs
 2 bond pairs
 3 lone pairs
 180o bond angles
XeF2
Examples
 PF3
4 total
3 bond
1 lone
AX3E
F P F
F
TRIGONAL
PYRAMIDAL
<109.5°
Examples
 CO2
2 total
2 bond
0 lone AX2
O C O
LINEAR
180°
Examples (You Try)
 NO2-
 C 2H 2
 CH3OH
Molecular Polarity: Dipole Theory
Polar Molecule: The negative (electron charge) is not distributed symmetrically
among the atoms
Dipole Moment
 Direction of the polar bond in a molecule.
 Arrow points toward the more e-neg atom.
+

H
Cl

Determining Molecular Polarity
 Depends on:
 dipole moments
 molecular shape
Nonpolar Molecules
 Dipole moments are symmetrical and cancel each other
out.
F
Non-Polar
BF3
B
F
F
Polar Molecules
 Dipole moments are not symmetrical.
 The bonds are between elements with an electronegativity
difference of between 0.3-1.7.
3.4
***Presence of lone pairs is
a good trick to decide if
its polar*****
H2O

O
Polar
2.2
H
2.2
+

H
Group Activity
VSEPR shape and Polarity Flashcards.
In groups of three.
Find yourself a place to work in the
room with a flat surface.
Decide who is A, B, and C.
Group Activity
VSEPR shape and Polarity Flashcards.
A’s  Come up to the front and get
three markers of different colours.
B’s  Come up to the front and get 5
coloured flashcards.
C’s  You just sit tight…for
now…..your time is coming.
Group Activity
VSEPR shape and Polarity Flashcards.
When all the group members are back
and the group has all the materials
you may do the following:
A’s: On the back (side without lines) of
two cards you will draw the base
shapes of Linear & Trigonal Planar.
Group Activity
VSEPR shape and Polarity Flashcards.
B’s: On the back (side without lines) of
three cards you will draw the base
shapes of Bent, derived shape of
trigonal planar and Tetrahedral.
C’s: On the back (side without lines) of
three separate cards you will draw the
two derived shapes of Tetrahedral.
A,B & C: You’re the “Editors” to make
sure their all correct pictures.
Group Activity
VSEPR shape and Polarity Flashcards.
LINEAR
180°
BENT
<109.5°
BENT
<120°
TRIGONAL PLANAR
120°
TETRAHEDRAL
<109.5°
TRIGONAL
PYRAMIDAL
<109.5°
You have about 10
minutes.
Group Activity
VSEPR shape and Polarity Flashcards.




Once all three group members have finished
drawing the base and derived molecules
shapes on the flashcards….
Each group member will find a actual
compound that is a match to the shapes they
drew earlier
Each person will then draw the structural
formula of the compound they found on the
other side of the flashcard, along with the
chemical formula of the compound.
You have 10 minutes to complete pictures.
Group Activity
VSEPR shape and Polarity Flashcards
EXAMPLE
FRONT:
TETRAHEDRAL
<109.5°
BACK:
CH4
VSEPR LAB
CH2Cl2
CCl4
CH3Cl
BH3
C2H2
PCl3
NF3
AlCl3
HClO
C2Cl2
C2H4
C2H6
C2H2Cl
NH3
HCN
SiF4
Cl2O
H2O2
Group Activity
VSEPR shape and Polarity Flashcards.

Once every member of your group has
completed their flashcards, share them in
your group, everyone must be able to do ALL
of the shapes by themelves!!

HINT: There are assignments in your
assignment bundle that are very close to what
you are doing…..

You have about 10 minutes to share in your
group.
Meniscus?
Capillary action?
Is water attracted to a charged object?
 Intramolecular bonds: attractions
within a molecule.
 Intermolecular forces: The weak
forces or bonds between molecules.
*** Think of a highway
in the states.
***What’s the highway
that goes between
different states called?
The Interstate.
•Intermolecular bonds involve the electrostatic
attractive forces between molecules.
•Ionic substances do not form molecules.
• Therefore, intermolecular bonding only occurs
in substances that form covalent bonds.
(molecular Compounds)
S
T
R
O
N
G
W
E
A
K
3.4
Cl2.2
2.2
Types of Intermolecular Bonding

Van der Waals forces can be
divided into three different
types



London dispersion forces
dipole-dipole forces
hydrogen bonding
WEAKEST
STRONGEST
London Dispersion Forces
 Weak attractive forces that result when the electrons of one
molecule are attracted to the positive nuclei of a nearby
molecule (random chance)
 The movement of the electrons in a molecule generates
temporary positive and negative regions in the molecule
 A temporary dipole
***Temporary***
Let us consider the Chlorine molecule, Cl2(g):
•At a particular instant,
we may find that the
two electrons that form
the bond may be closer
to one nucleus than the
other.
•Results in a
temporary dipole with
one end more
negative than the
other.
Dipole-Dipole Forces
- The positive end of one polar molecule will be
attracted to the negative end of a
neighbouring polar molecule. (Permanent)
Hydrogen Bonding




Occurs when hydrogen is bonded to a highly
electronegative element (fluorine, oxygen and nitrogen) –
chemistry is FON!!!
The hydrogen end of the bond takes on a strong positive
charge because of the exposed positive nucleus, while the
other element takes on a strong negative charge
This positive hydrogen will be attracted to nearby negative
atoms.
It appears as though the hydrogen atom bonds to different
molecules.
Predicting Boiling Points
 What affects boiling point?
1. Number of Electrons
2. Number of Carbons
3. Polarity…electronegativity difference
Predicting Boiling Points
In order to cause a substance or compound to boil,
You must provide enough energy to break the bonds in the
compound.
More energy needed, higher temperature needed.
In order to describe how a
solid looks and behaves,
you first need to
determine what class of
substance it is.
Ionic Crystals
Metallic Crystals
Molecular Crystals
Covalent Network Crystals
 Very hard (harder than ionic and molecular crystals)
 Brittle with very high melting and boiling points
 Higher than ionic and molecular
 Insoluble nonconductors of electricity
Carbon based covalent network crystals
GROUP ACTIVITY :
THINGS TO WATCH FOR

SATP: Standard Ambient (room) Temperature and
Pressure.

“Compound” refers to the joining of elements (MetalNonmetal).
Ion charges of metals DO NOT have to be memorized,
there on the periodic table.
Ion charges of non-metals must be memorized, as
they are not on the periodic table.
“Molecule” refers to a molecular joining of elements
(Nonmetal-Nonmetal).



INDIVIDUAL----->PAIR------>GROUP---->SHARE

FIRST
Form pairs of two….quickly and quietly.
 Switch desks so that pairs are sitting next to each
other.
 Each pair decides who is going to be A and who is
going to be B.

INDIVIDUAL----->PAIR------>GROUP---->SHARE

A: Reads Intro and all Ionic “Stuff” (p 119-123).

B:Reads Intro and all Molecular/Covalent “Stuff”
(p. 119(Intro) and 124-126).
As each partner reads over their pages, summarize
the information (write it down).
Summarize in such a way that when you are done
summarizing you can explain it to your partner.
First part is individual, so it should be pretty quiet.
You have about 25 minutes.



INDIVIDUAL----->PAIR------>GROUP---->SHARE
Second
 Now each pair has the next 15 minutes to explain
their sections to each other…BOTH must have the
summarized notes.


It should be a little louder now, but you should only
be talking to your partners.

When you are done, bring up your summaries to
me and so I can have a look.
INDIVIDUAL----->PAIR------>GROUP---->SHARE

Third
The pairs should now join up with one other pair
and make a “Place Mat” of their summarized
ideas.
 4 People in each
group.
 Divide into A, B, C,
D, etc.
 WAIT!

PLACE MAT ACTIVITY

B’s get the markers. At the front.

C’s get the Paper. At the back.

A’s Divide paper into section with shared area
in middle.

WAIT!
PLACE MAT ACTIVITY

THIS TIME!

Each group member summarizes the
information they have in their “spot” on the
place mat.

You have 5 minutes

WAIT!
OK, NOW YOU CAN GO.
Four
 You then rotate the “Place Mat” and check over
and add (if necessary) information to the next
persons summary.
 You will keep rotating the “Place Mat” and
checking/adding until you have rotated 4 times
and are back to where you started (your own).


Use your time wisely. You have 15 minutes.

Summarize the information your group found in
the middle.

Sort of a final copy..

Summarize the information your group found in
the middle.

Sort of a final copy..