CP NT Ch 8 & 9—Covalent Compounds
Why do atoms bond?
 Atoms want ______________—to achieve a noble gas configuration (__________)
 For _______________ bonds there is a _________________ of electrons to get an octet of electrons
 For covalent bonds there is a _______________ of electrons to get an octet
What is a covalent bond?
 Covalent bond—the chemical bond that results from sharing of _______________ electrons
— Occurs with elements _______________to each other on the periodic table
— Between a nonmetal and a nonmetal
— Molecule is two or more atoms are bonded _______________
Examples of Molecules
 F2
 H2O
 NH3 (_______________)
 CH4 (_______________)
 Notice there are no _______________, only non-metals
Diatomic molecules
 Some atoms do not exist as a _______________atom
 Atoms that exist as two:
H2, O2, N2, Cl2, Br2, I2, F2
 HONClBrIF
 Magnificent _____-don’t forget H
3 Types of Covalent Bonds:
 Single, double, triple
Strength of Covalent Bond
 Several factors control bond strength
o _______________ of shared electrons—the __________ electrons shared, the shorter the bond,
and the greater the bond strength
o _______________ of the atoms
Single Covalent Bonds
 Each atom shares one _______________ (___________) of electrons
 _______________ bond
 _________________ bond of the three
Double Covalent Bonds
 Each atom shares _______ pairs (___________) of electrons
 ________________ length bond
 ________________ strength bond
Triple Covalent Bond
 Each atom shares _____________ pairs (___________) of electrons
 ______________ bond
 _______________ bond
 Carbon, nitrogen, oxygen, and sulfur can form double and _______________covalent bonds
Covalent Molecule Properties
 Tend to be _______________solids, liquids, or gases at room temperature
 __________ melting and boiling points
 __________ conductors of heat and electricity
 Non-electrolytes – do not conduct electricity in water
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Naming Binary Molecular Compounds
3 Rules to Name
1. Name the first element using the _______________name
2. Second element in the formula- use the root word and __________ in –ide
ex: Oxygen  Oxide
Sulfur  Sulfide
Hydrogen _______________
3. Add a prefix to ___________ words to indicate the _______________of atoms
1. mono-
6. hexa-
2. di-
7. hepta-
3. tri-
8. octa-
4. tetra-
9. nona-
5. penta-
10. deca-
Exceptions to the Rules
1. When the formula contains one atom of the _______________element, omit (leave out) mono
ex: CO2  ______________ ________________, not Monocarbon Dioxide
2. Drop the final letter in the prefix if the element begins with a _______________
- For prefixes 1 & 4-9
ex: CO  ______________ __________________, not Carbon monooxide
Example 1
S4N2  _______________ _______________
Example 2
SO3  _______________ _______________
Example 3
P4S5  _______________ _______________
Example 4
CO  _______________ _______________
Example 5
NH3  _______________ _______________ common name: _______________
Example 6
CH4  _______________ _______________
Example 7
As2O3  _______________ _______________
Example 8
N2O5  _______________ _______________
Example 9
P4O10 _______________ _______________
common name: _______________
Lewis Structures
5 rules for Lewis Structures
1. Find the total # of _______________electrons for the molecule
2. Find the center atom (the element with the _______________ # of atoms)
3. Draw bonds. Connect the other atoms to the center atom. Then subtract ____ electrons from the total # of
valence electrons for each bond drawn.
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4. Distribute electrons around each atom to give a total of ____ electrons except H, Al, B, & Be
5. If there are not enough _______________to give 8 around each atom, create _______________& triple
bonds.
Example 1: CF4
Example 2: NH3
Example 3: H2S
Ions: With Ions we add or take away _______________ . Put ions in _________________ and the charge on the
outside
Example 4: NH4+
Double Bonds
Example 5: CO2
Triple Bond
Example 6: CO
Incomplete Octet
Example 7: BCl3
Boron does not have 8e- around it. It is _______________but it is okay; it is one of our exceptions.
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NT: VSEPR- V_________________ S__________ E__________________ P_________ R____________________
 Lewis Structures are _________________, whereas VSEPR Molecules are _________________
 VSEPR predicts the _________________ or “_________________” of the molecule
 ________ pairs of electrons influence the shape by pushing other atoms as far apart from each other as possible
Molecular Shape
Molecule
Lewis Structure
Number of
(shared)
bonding pairs
of electrons
Number of lone
pairs of
electrons
Total Number
of electron
pairs
Molecular
Shapes
(look at lone &
bond pairs)
H2
Linear
CO2
Linear
BH3
Trigonal
planar
SiF4
Tetrahedral
PH3
Trigonal
Pyrimidal
SCl2
Bent
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Ball and Stick Model
Bond
Angle
Electronegativity and Polarity
Electronegativity

Relative ability of an atom to _______________electrons in a chemical bond

_______________has the highest e.n. value.

Trend: ______________ across (left to right) a period and _______________ down a group
Types of Covalent bonds

Non-polar covalent— _______________sharing of electrons

Polar covalent—unequal _______________of electrons
Non-polar covalent

Sharing of electrons _______________

Usually occurs when two _______________atoms are bonded together.

Examples: H2, O2, N2, Cl2, Br2, I2, F2
Polar covalent

_______________ sharing of electrons

unequal sharing caused by 2 elements with different ____________________ (different abilities to
attract electrons)

The bond is called a dipole (two poles)

Creates a molecule with _____________ charges

Partial charges symbolized by (delta) + and -

The _______________electronegative atom is located at the partially negative end

Example:
+ ⟼
H—Cl or H—Cl
Polar molecule or not?



The _______________ of a molecule usually tells if a molecule is polar or not
If the VSEPR shape is _______________it is usually non-polar
If the molecule is asymmetric it is _______________
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Name
Polar example
VSEPR Model
Non-polar example
VSEPR Model
Linear
HCl
CO2
Trigonal planar
CH2O
AlH3
Tetrahedral
CH3OH
CH4
Trigonal
pyramidal
NH3
N/A
N/A
H2O
N/A
N/A
(Always polar)
Bent
(Always polar)
Intermolecular Forces
 The force that exists between ________________ molecules
 This force attracts molecules to each other
 3 Types
— Dispersion force or induced dipole moment between molecules; only force in ___________
molecules; _____________ force (caused by the motion of electrons) (Ex: CH4)
— Dipole-dipole the force between two _____________ molecules; ______________ force
(Ex: HCl)
— Hydrogen bond forms between the hydrogen end of one dipole and fluorine, oxygen, or
nitrogen (that have at least one lone pair) end of another dipole; _______________force
(Ex: H2O)
 _______ molecules have dispersion forces
 All ____________ molecules have dipole-dipole forces and dispersion forces
 Molecules that hydrogen bond have all 3
Solubility of polar molecules
 _______________ are due to intermolecular forces
 Like dissolves like
— Polar substances will dissolve _______________molecules (and ionic compounds)
— ___________________ substances will dissolve non-polar molecules
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