Unit 5: Soap
Unit 5 Section 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.
Short Cut Rule: the _______________ next to the letter A of the Representative elements
represents the valence electrons.
Oxygen
What is a BOND and why do they form?



A bond is like ___________ holding atoms together
However, it’s 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.
o
o
o
o
o
o
IONIC:
Metals have fewer valence electrons and much lower ___________________ than non-metals
Therefore, metals tend to __________ their electrons and non-metals _________ electrons
Metals become __________________ (positively charged)
Non-metals become _______________ (negatively charged)
There is a __________________ of electrons
The cation & anion are electrostatically attracted because of their charges—forming an ionic bond
Potassium + Chlorine  Potassium Chloride
Barium + Fluorine  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 evenly share electrons form ___________________ covalent bonds
o 2 Different Non-metals that unevenly share electrons form ___________________ covalent bonds
NONPOLAR COVALENT BONDING
Chlorine
+
Chlorine

POLAR COVALENT BONDING
Hydrogen
+
Fluorine
Chlorine gas

Hydrogen Fluoride
C. METALLIC:
o
o
o
Metals form a pool of electrons that they share together.
The ______________________ are free to move (mobile) throughout the structure—like a sea of electrons
Atoms are bonded as a network
**Bonding is never purely ionic nor purely covalent**
Using Electronegativity to determine Bond Type
One can use electronegativity values to predict the primary type of bonding that exists but it is not 100% accurate
If the electronegativity difference is:
 Between 0-.40
NONPOLAR COVALENT
 Between .41.and 1.69
POLAR COVALENT
 Greater than 1.7
IONIC
0--------.4-----------------------------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
o Ionic bonds tend to have very high melting/boiling point as it’s hard to pull apart those electrostatic attractions.
They’ tend to form hard ___________________under normal conditions
o Polar covalent bonds have the next highest melting/boiling points.
Most are soft solids or liquids under normal conditions
o Non-polar covalent bonds have lower melting/boiling points.
They tend to form ______________________.
Solubility in Water (polar)
o
o
o
Ionic & polar covalent compounds tend to be ___________________ in water.
Non-polar & metallic compounds tend to be __________________ in water.
Remember the “Like Dissolves Like” Rule of Thumb
Electrical conductivity
o
In order to conduct electricity, charge must be able to move or flow
o _____________________bonds have free-moving electrons—they can conduct electricity in solid and liquid
state
o Ionic bonds have free-floating ions ONLY when dissolved in _________________ or in liquid form that allow
them conduct electricity
o Covalent bonds _____________________ have charges free to move and therefore cannot conduct electricity in
any situation
Structure
o Ionic compounds form a __________________________ framework. Alternating positive & negative ions
are closely packed.
o Covalent compounds form neutral discrete _______________________
o Metals are represented as a framework of metals atoms closely packed with a sea of
__________________. All the nuclei of the cations are closely packed with delocalized electrons moving
throughout the solid.
Unit 5 Section 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; HOWEVER, _____________________always goes
in the middle
Steps to Drawing a Molecule
1. Count how many valence electrons should be around each atom
2. Arrange the atoms in a skeletal structure and connect them with a bonding pair of electrons
represented by a dash
3. Place any remaining electrons around each atom 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
o _____________________ bonds form when 2 pairs of electrons are shared between two atoms
o A ____________ pair is a pair of electrons not shared—only one atom “counts” them
o 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
Bond Length (distance between atoms)
Longer
Bond Strength
Weaker
Bond Energy (energy needed to break a bond)
Lower
Triple Bond
Shorter
Stronger
Higher
Drawing Polyatomic Ions
Polyatomic Ions: they are a group of atoms bonded together that have an _________________ charge.
 Follow same rules as before but add the number of electrons of the negative charge or subtract the
number of electrons of the positive charge
Example:
Carbonate, CO3-2
Unit 5 Section 3 Molecules in 3D
o
o
o
o
Bonds are pairs of electrons. Electrons are ________________________ charged
Negative charges repel other negative charges
Thus, bonds repel each other
So, molecules arrange themselves in 3-D so that the bonds are as ___________as possible from each other
VSEPR Theory
o VSEPR stands for
______________________________________________________________________________.
o 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.
 “A” stands for the central atom
 “B” stands for the number of bonding atoms off the CENTRAL ATOM [A].
 “E” stands for the number of lone pair coming off the CENTRAL ATOM [A].
What shapes do molecules form? The 5 basic shapes are
Linear: 2 bonds, no lone pair – or – any 2 atom molecule
Code
AB2
AB2
Formula
BeCl2
HCl
Degrees
between Bonds
Draw Both Shapes Here
Trigonal Planar: 3 bonds, no lone pair
Code
AB3
Formula
BF2
Degrees
between Bonds
Draw Shape Here
Tetrahedral: 4 bonds, no lone pair
Code
AB4
Formula
CH4
Degrees
between Bonds
Draw Shape Here
Trigonal bi-pyramidal: 5 bonds, no lone pair {not on test}
Code
AB5
Formula
PCl5
Degrees
between Bonds
Draw Both Shapes Here
Octahedral: 6 bonds, no lone pair {not on test}
Code
AB6
Formula
SF6
Degrees
between Bonds
Draw Both Shapes Here
How is shape affected by lone pair of electrons on central atom?
 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 the lone pairs around the CENTRAL atom, not the outside atoms!
BENT or V-SHAPED: 2 bonds, 1 lone pair – or – 2 bonds, 2 lone pair
Code
AB2E
Formula
SO2
Degrees
between Bonds
Draw Both Shapes Here
Code
AB2E2
Formula
H2O
Degrees
between Bonds
Draw Both Shapes Here
Trigonal Pyramidal: 3 bonds, 1 lone pair
Code
AB3E
Formula
NH3
Degrees
between Bonds
Draw Both Shapes Here
 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.
Examples: Compare CH4 vs H2O
CH4 (Tetrahedral)
H2O ( bent)
bond angles…………….
bond angles………