Chap 6: Covalent
Compound
HRW 6.1 Chemical Bonding
•
•
•
•
•
•
•
•
•
•
•
Define chemical bond.
Explain why most atoms form chemical bonds.
Describe ionic & covalent bonding.
Explain why most bonding is not purely covalent or ionic.
Classify bonding type according to electronegativity differences.
Define molecule & molecular formula.
Explain relationships between potential energy, distance between
approaching atoms, bond length and bond energy.
State the octet rule.
List steps to write Lewis structures.
Determine Lewis structures with single or multiple bonds or both.
Describe how resonance structures are used.
6-1: Types of Chemical Bonding
• Ionic - results from the _______________
between large numbers of _________ and
__________.
• Covalent - results from the __________of
electron pairs between two ________.
• Metallic - results from chemical attraction
between metal atoms and the sea of
electrons surrounding them.
What about covalent
compounds?
• The electrons in each atom are attracted to
the ___________of the other.
• The electrons ___________ each other,
• The nuclei ____________each other.
• They reach a distance with the
__________________ energy.
• The distance between is the
______________.
How does H2 form?
• The nuclei repel
+
+
How does H2 form?
• The nuclei repel
• But they are attracted to electrons
• They share the electrons
+
+
Figure 6-2
• _____________ = space that shared
electrons move within
Energy and Stability
• Energy is __________
when atoms from a
covalent bond
• Potential Energy = stored
energy
• Potential energy
determines ___________
• Bond length is the
distance between two
atoms at their
__________ potential
energy
• Figure 6-4
Energy
Potential
0
Internuclear Distance
Tr32 p 165 Bond Length & Stability
What’s H2 molecule’s bond length?
What does the dashed line represent?
How much energy must be added to the bond to break it?
Why doesn’t He form a diatomic molecule?
Bond Energy
• Bond forming = __________________
– _____________ , Negative
• Bond breaking = ______________
– Endothermic, __________
• As bond length____________, more
energy is _______. Therefore, more
energy is needed to break the bond.
• As bond length _________, _________
the bond energy
Chapter 6
Visual Concepts
Bond Energy
Tr 32G p. 173
What is the trend between number of bonds and bond
energy (bond strength) ?
What is the trend between bond number and bond length?
Check Your Understanding!
1. Why is the H2 molecule more stable than the
individual atoms that bond to form it?
2. What is created when two atomic orbitals
overlap?
3. What happens to the potential energy of
two atoms as they approach each other
to form a covalent bond?
4. What name is given to the distance
between two atoms in a covalent bond at
which the potential energy is minimum?
5. List an example of substances that have
covalent bonds
Covalent Bonding
• Covalent bond: __________________
• Nonpolar covalent bond- electrons are
______________________
• Polar covalent bond- shared
____________
E. Chart of electronegativities:
F. Element with highest electronegativity: ___
Lowest electronegativity: _______
Chapter 6
Visual Concepts
Electronegativity
Dipole
• A molecule that has one partial positive
end and one partial negative end
• The symbol  is used to mean ____________.
• _____ is used to show a partial positive charge
• _____ is used to show a partial negative charge charge
• example: __________
• Because the F atom has a partial negative charge, the
electron pair is more likely to be found nearer to the
fluorine atom
VI. Bond Polarity
A. Non-Polar Covalent (_______ Covalent) Bond
1. Electrons are shared _______.
2. The bond has 100% covalent character and 0% _____
character.
3. The electronegativity difference between atoms in the bond
is __.
4. Example: H2
H
H
B. Polar Covalent Bond
1. Electrons are shared _________.
2. The electronegativity difference between atoms in the bond
is < ___.
H
Cl
3. Example: HCl
C. Ionic Bond
1. Electrons are transferred from one atom to the other.
2. The electronegativity difference between atoms in the
(Rule of 2.0)
bond is ≥ ___.
3. Example: LiF
The electronegativity difference is so great that the electron is
transferred to the non-metal.
D. Examples: What type of bonds are each of the following:
(a) KF
(b) HBr
(c) Br2
(d) H2S
(e) CsCl
Tr31 Predicting Bond Type from Electronegativity Differences
EN for K = 0.8 & Cl = 3.0.
What type bond for KCl?
H = 2.1; S = 2.5. For H2S?
P = 2.1; H = 2.1. For PH3?
Electronegativity Problems
• Determine the electronegativity difference
bond type and the more-negative atom for
the following:
C
and H
C and S
O and H
Na and Cl
Cs and S
Tr30 Fig. 6.1 p. 162 Ionic vs. Covalent Bonding
Why does the ionic bonding occur between atoms of different
sizes?
Why are similar-sized atoms covalently bonded?
Comparison of electron density of
nonpolar H-H bond with polar H-Cl bond
Why is the electron density greater around chlorine?
Chapter 6
Lecture
Notes Part 2
Bonding in Metals
• Metallic Bonds:
– Metallic bonds- consists of the attraction of
the __________________________ for the
___________charged _______________
– Properties of metals– __________: Conducts heat & electricity
– ___________: Can be drawn into wires
: Can be hammered or forced into
– _____________
shapes
6.2 Drawing and Naming
Molecules
Electron Dot diagrams
• A way of keeping track of
_____________ electrons.
– Write the symbol.
– Put one dot for each valence
electron.
– Don’t pair up until they have
to.
X
Tr26A Fig 6.10 p 170 Electron Dot Notation
Why are we interested
in the outermost e-s?
How does electron-dot
notation help us
determine bond
formation?
The Electron Dot diagram for
Nitrogen
Nitrogen has 5 valence
electrons.
 First we write the symbol.
Then add 1 electron at a
time to each side.
Until they are forced to pair up.

Electron Dot Structures
Symbols of atoms with dots to represent the
valence-shell electrons
1
2
13
14
15
16
17
H
Li
18
He:

Be

B 


C



N




O :F



:Ne :

Learning Check

A.
X would be the electron dot formula for
1) Na
B.
2) K
3) Al

 X  would be the electron dot formula

1) B
2) N
3) P
The Octet Rule:
The Diatomic Fluorine Molecule
F
F
1s
2s
2p
1s
2s
2p
Each has
seven valence
electrons
The Octet Rule:
The Diatomic Oxygen Molecule
O
O
1s
2s
2p
1s
2s
2p
The Octet Rule:
The Diatomic Nitrogen Molecule
N
N
1s
2s
2p
1s
2s
2p
Lewis Structures
Lewis structures show how
__________________ are arranged among
atoms in a molecule.
Lewis structures reflect the central idea that
stability of a compound relates to
_______________ electron configuration.
Shared electrons pairs are covalent bonds and
can be represented by two dots (:) or by a single
line ( - )
The HONC Rule
Hydrogen (and Halogens) form
__________________ bond
Oxygen (and sulfur) form __________________
bonds
One double bond, or two single bonds
Nitrogen (and phosphorus) form _______________
bonds
One triple bond, or three single bonds, or one double bond
and a single bond
Carbon (and silicon) form ________________
bonds.
Two double bonds, or four single bonds, or a triple and a
single, or a double and two singles
Covalent Bonding
• Single Covalent Bonds: a bond in
which ____atoms share a pair of
______________
– Unshared pairs- _____________ or
_________________ pairs;
• pairs of e- not shared between atoms
– Water molecule
X. Lewis Structures of Covalent Compounds: A representation of the
covalent bonding in a molecule.
A. Covalent bonds are shown as ______. Example: H2 H H
B. Lone pairs of electrons are shown as _____. Example: O2
C. ONLY ________ electrons are shown.
••
••
O O
D. General steps for drawing Lewis structures:
••
••
1. Sum the valence electrons in the compound.
2. Add __ for each negative charge. Subtract 1 for each __ charge.
3. Generally place the element that makes the _________ number of
bonds in the center.
(least electronegativity)
4. Draw ______ bonds to the other atoms off of the central atom.
5. Place electrons around the ______ atoms until an _____ is
reached.
6. If you run out of electrons, start forming ______ or _____ bonds.
7. If you have EXTRA electrons after all have octets, place them on
the ______ atom.
8. In the end, all atoms should have an octet that need an octet
(____ is an exception), and the total number of electrons should be
placed on the molecule.
Try These!
1. HBr
2. NH3
3. CH3NH2
Deficient Electrons
Deficient 2 e-:
Make a ____________ into a ___________.
Deficient 4 e-:
Make one single bond into a ____________
or make 2 single bonds into 2
_____________
1. Count the total number of valence electrons
2. Place the atom that makes most bonds in the middle. (least electronegative
other than hydrogen)
3. Draw single bonds to the other atoms off of the central atom.
4. Place electrons around peripheral atoms to fill octet, then to central atom
5. if electrons run out, start making double, triple bonds
4. CH4
5. O3
5. CO2
1. Count the total number of valence electrons
Add 1 electron for each – charge , subtract 1 from each + charge
2. Place the atom that makes most bonds in the middle.
3. Draw single bonds to the other atoms off of the central atom.
4. Place electrons around peripheral atoms to fill octet, then to central atom
5. if electrons run out, start making double, triple bonds
1. SO42-
2. NH4+1
3. IO-1
Resonance
Occurs when more than one valid Lewis
structure can be written for a particular
molecule.
O
O
O
These are resonance structures.
the double side arrow is used to show that
the actual molecule is the AVERAGE of the two.
Resonance in Benzene, C6H6
Binary Covalent Compounds
 Compounds between two _____________
 First element in the formula is named first.
 Keeps its element name
 Gets a prefix if there is a subscript on it
 Second element is named second
 Use the root of the element name plus
the _________ suffix
 Always use a prefix on the second
element
 1 = ______
 2 = _______
 3 = _______
 4 = _______
 5 = ______
 6 = hexa
 7 = hepta
 8 = octa
 9 = nona
 10 = deka
List of
Prefixes
Naming Covalent Compounds
P2O5 =
CO2 =
CO =
N2O =
Practice – Write the Formula
Compound Name
Carbon dioxide
Carbon monoxide
Diphosphorus pentoxide
Dinitrogen monoxide
Silicon dioxide
Carbon tetrabromide
Sulfur dioxide
Phosphorus pentabromide
Iodine trichloride
Nitrogen triiodide
Dinitrogen trioxide
Compound Formula
Practice – Name the Compounds
Compound
Formula
N2O4
SO3
NO
NO2
As2O5
PCl3
CCl4
H2O
Compound Name
Don’t Confuse with Ionic
Compounds!
• Ionic Compound = Metal + NonMetal
• Just add –ide on nonmetal! ONLY!!!
– Na+ + Cl- = NaCl
• _______________________ ( -ide at the
nonmental)
– Li2O = 2 Li + + O• ____________________
Covalent (Molecular)
Compound
Ionic Compound
VSEPR
Model/Theory
(Valence Shell Electron Pair Repulsion)
Valence Shell Electron Pair Repulsion
Theory
_________________________ of a
molecule based on Lewis structures
The structure around a given atom is
determined principally by minimizing
electron pair repulsions.
VSEPR
• For __________, the two double bonds around
the central carbon atom _________ each other
and remain far apart.
• For BF3, the three single bonds around the
central fluorine atom will be at a maximum
distance apart.
Predicting a VSEPR
Structure
Draw _________________.
Count the number of bonds. (count double and
triple bonds as 1 bond)
Count the number of lone pair of electrons on
the ______________ atom
Determine the name of molecular structure
from positions of the atoms from the chart.
Determine the Molecular Shape of
SCl2
Determine the shape of these molecular
compounds
1. PF3
2. CO2
3. NH4+
Polarity
A molecule, such as HF, that has a
center of positive charge and a center
of negative charge is said to be
______________, or to have a
_____________________
Determine if the molecule is polar
1. Indicate partial charges
2. Draw the arrow going towards the - .
3. Add the arrows .
4. If arrows do not cancel out, then it is polar! .
O
O
H
C
O
H
F
P
F
F