Unit 5 : Soap
5.1 Bonding
Why do atoms bond?




Atoms are most ______________ when they’re outer shell of electrons is full
Atoms bond to fill this outer shell
For most atoms, this means having 8 electrons in their valence shell: _____________
Rule
Common exceptions are Hydrogen and Helium which can only hold 2 electrons:
____________________
Valence Electrons


___________________ electrons are found in the highest energy levels
and play the game of bonding. Valence electrons are displayed in
Lewis dot diagrams.
oxygen
Short Cut Rule: the _______________ next to the letter A of the Representative
elements represents valence electrons.
What is a BOND and why do they form?
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

A bond is like ___________ holding atoms together
It is really __________________ of attraction between VALENCE ELECTRONS holding
atoms together
They form because it ____________________ the potential energy of the atoms and
creates ____________________.
3 Types of Bonding
A. IONIC:
o Metals have fewer valence electrons and much lower ___________________ than
non-metals
o Therefore, metals tend to __________ their electrons and non-metals _________
electrons
o Metals become __________________ (positively charged)
o Non-metals become _______________ (negatively charged)
o There is a __________________ of electrons
o The cation & anion are electrostatically attracted because of their charges—forming
an ionic bond
Potassium
+
Barium
+
Chlorine
Fluorine


Potassium Chloride
Barium Fluoride
B. COVALENT
o When two non-metals bond, neither one loses or gains electrons much more
easily than the other one.
o Therefore, they __________electrons
o 2 Identical Non-metals that share electrons evenly form ___________________
covalent bonds
o 2 Different Non-metals that share electrons un-evenly form
___________________ covalent bonds
NONPOLAR COVALENT BONDING
Chlorine
+
Chlorine

Chlorine gas
POLAR COVALENT BONDING
Hydrogen
+
Fluorine

Hydrogen Fluoride
C. METALLIC:
o Metals form a pool of electrons that they share together.
o The ______________________ are free to move throughout the structure—like a
sea of electrons
o Atoms are bonded as a network
***********Bonding is never purely ionic nor purely covalent*************
Using Electronegativity
Use electronegativity values to predict the primary type of bonding that exists
If the electronegativity difference is:
 Greater than 1.7
 Between 0-.2
 Between .3 and 1.6
IONIC
NONPOLAR COVALENT
POLAR COVALENT
0--------.3-----------------------------1.7------------------------------------------------3.3 EN difference
0%
5%
50%
100%
IONIC CHARACTER
Examples
Elements Bonded
Electronegativity
Difference
Bond Type
CH4
F2
H2O
NaF
Bond Type Affect Properties
State of Matter and Melting Point /Boiling Point
 Ionic bonds tend to have very high melting/boiling point as it’s hard to pull apart
those electrostatic attractions.
They’re found as ___________________under normal conditions
 Polar covalent bonds have the next highest melting/boiling points.
Most are solids or liquids under normal conditions
 Non-polar covalent bonds have lower melting/boiling points.
Most are found as _____________________________.
Solubility in Water (polar)
 lonic & polar covalent compounds tend to be ___________________ in water.
 Non-polar & metallic compounds tend to be _______________________
 Like Dissolves Like Rule of Thumb
Electrical conductivity
In order to conduct electricity, charge must be able to move or flow
 _____________________bonds have free-moving electrons—they can conduct
electricity in solid and liquid state
 Ionic bonds have free-floating ions when dissolved in _________________ or in
liquid form that allow them conduct electricity
 Covalent bonds _____________________ have charges free to move and therefore
cannot conduct electricity in any situation
5.2 Drawing Molecules
Rules for drawing Lewis Dot Diagrams
o In general, write out the atoms in the same order as they appear in the chemical
formula
o Hydrogen & Halogens (F, Cl, Br, I) can only bond with one other atom. Always put them
around the ______________________________
o The least electronegative atom is usually in the middle; _____________________always
goes in the middle
Steps
1. Decide how many valence electrons are around each atom
2. Arrange the atoms in a skeletal structure and connect them with a bonding pair of
electrons
3. Place remaining electrons around atoms so they each acquire 8 electrons. Exception is H.
Examples with Single Bonds:
CH4
CH3I
PCl3
Bonding Pair: pair of electrons shared by two atoms…they form the “bond”
Multiple Covalent Bonds
 Are needed when there is not enough electrons to complete an _______________
 To satisfy: move ________________________ in between atoms to satisfy the
duet/octet rule
Examples with Double Bonds:
CH2O
Double Bonds & Lone Pair
 _____________________ bonds are when 2 pairs of electrons are shared between the
same two atoms
 ___________________pairs are a pair of electrons not shared—only one atom “counts”
them
 A _____________________occurs when two atoms share 3 pairs of electrons.
Examples with Triple Bonds:
C2H2
You Try! Single, Double, or Triple Bonding
HCN
CO2
Properties of Multiple Bonds
Single Bond
Double Bond
Triple Bond
Bond Length (distance between atoms)
Longer
Shorter
Bond Strength
Weaker
Stronger
Bond Energy (energy needed to break a bond)
Lower
Higher
Drawing Polyatomic Ions
Polyatomic Ions: they are a group of atoms bonded together that have a _________________
charge
Example:
Carbonate, CO3-2
5.3 Molecules in 3D
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Bonds are electrons. Electrons are ________________________ charged
Negative charges repel other negative charges
Bonds repel each other
Molecules arrange themselves in 3-D so that the bonds are as
______________________ as possible
VSEPR Theory
 VSEPR stands for _______________________________________________________________.
 It is the theory used to predict the three dimensional shape of a molecule.
Coding for a Shape
A code can help you connect the molecule to its shape. An “A” stands for the central atom and
a “B” stands for the number of bonding atoms off the central atom. An “E” stands for the
number of lone pair coming off the CENTRAL atom.
What shapes do molecules form?
The 5 basic shapes are
Shape
Code
Formula
Molecular Structure
2 bonds, no lone pair or any 2 atom molecule
LINEAR
AB2
BeCl2
AB
HCl
3 bonds, no lone pair
TRIGONAL PLANAR
AB3
BF3
4 bonds, no lone pair
TETRAHEDRAL
AB4
CH4
 Lone pairs are electrons, too…they must be taken into account when determining molecule
shape since they repel the other bonds as well.
 But only take into account lone pairs around the CENTRAL atom, not the outside atoms!
2 bonds, 1 lone pair or 2 bonds, 2 lone pair
BENT or V-SHAPED
AB2E
SO2
AB2E2
H2O
AB3E
NH3
3 bonds, 1 lone pair
TRIGONAL PYRAMIDAL
 Lone pairs aren’t “controlled” by a nucleus (positive charge) on both sides, but only on one side.
This means they “__________________________” more than a bonding pair.
 They distort the angle of the molecule’s bonds away from the lone pair.
Compare CH4 (Tetrahedral)
bond angles…………….
vs
H2O ( bent)
bond angles…………….