Name______________________ #_____
Period___
Honors Chemistry
Ms. K
Pages 174-215
(excluding Section 4 and Intermolecular Forces, pages 203-207)
Task Checklist:
Look at Online Textbook Materials
Look at Section 1 Visual Concepts
Look at Section 2 Visual Concepts
Look at Section 3 Visual Concepts
Look at Section 5 Visual Concepts
(the first four only)
Take Section 1 Self-Check Quiz
Take Section 2 Self-Check Quiz
Take Section 3 Self-Check Quiz
Take Section 5 Self-Check Quiz
(Questions 1-4 only)
Chemical Bonding Concept Map
Other Tasks
Read Section 1
Read Section 2
Read Section 3
Read Section 5 (excluding
pages 203-207
Review Chapter Packet
Examine Class Website
caffeine
“In K nature, most atoms
Ms.
Chemistry
are joined to other
atoms by chemical
bonds.”
Name___________________________ #_______
Date__________________ Period______
Chapter 6 ? Chemical Bonding
All of nature favors arrangements in which
the lowest potential energy is achieved (more stable).
chemical bond ? a mutual electrical attraction between the nuclei and valence electrons of two atoms that
binds the atoms together
ionic bond ? a chemical bond that results from the electrical attraction between cations and anions. How
are the cations and anions formed? Electrons are taken from one atom by another atom.
ionic compound ? a compound composed of ions, bound together by ionic bonds
• solid at room temperature
• high melting points (thus are usually solid at room temperature)
• formula unit represents the lowest ratio of ions that combine to form a neutral compound
• most are crystalline solids
• when dissolved in water, the ionic compounds will break up into ions (dissociate)
• the solutions of ionic compounds will conduct electricity (electrolytes)
ions – can be cations or anions; can be monatomic or polyatomic
monatomic ions ? ions formed from one atom
Ex: Na + or O-2
polyatomic ions ? ions formed from two or more atoms bonded together
Ex: NH4+ or SO4-2
dissociation ? when an ionic compound dissolves in water to break apart into hydrated ions
NaCl(s)
H 2 O(l)
→ Na + (aq) + Cl- (aq)
electrolyte – An ionic compound that dissolves in water, producing ions that conduct electricity.
nonelectrolyte – A compound that does not produce ions when dissolved, not conducting electricity.
covalent compounds ? a compound composed of all nonmetals
• Also called molecular compounds
• solid, liquid, or gas at room temperature
• low melting points
• molecular formula represents the actual ratio of atoms that combine to form a neutral compound
• when dissolved in water, the molecular compounds DO NOT break up into ions
• the solutions of molecular compounds DO NOT conduct electricity (nonelectrolytes)
Figure 5: Page 179
Ionic, Polar Covalent, and Nonpolar Covalent:
•
•
•
0.0 to 0.3
0.3 to 1.7
1.7 and up
nonpolar covalent
polar covalent
ionic
What kind of bond would these atoms form?
1. N and S
2. S and C
3. Mg and Cl
4. C and F
5. Ba and O
Which one of these bonds
has the least ionic character?
Ms. K
Chemistry
Name_____________________________ #_____
Date______________________ Period_________
Electron Dot Structures
The valence electrons are the ones involved in bonding.
Electron-dot notation – the dots represent the valence electrons and the symbol
represents the nucleus and all the other electrons
Na
Mg
Al
Si
P
S
Cl
Ar
Ionic Compounds:
Na
+
Cl
Mg
+
Cl
+
Cl
Covalent Molecules:
H
+
H
Cl
+
Cl
Lewis Structures ? formulas for covalent compounds using the electron dot notation,
except, the bonds (electron pairs) are shown by dashes
H-H
Cl-Cl
O=O
single bond – 1 pair or 2 electrons are shared
double bond – 2 pairs or 4 electrons are shared
triple bond – 3 pairs or 6 electrons are shared
NO such thing as a quadruple bond.
How do you know how to draw these?
CH2Cl2
C
H
H
Cl
Cl
Ms. K
Chemistry
Name_____________________________ #_____
Date______________________ Period_________
VSEPR Theory Notes
octet rule ? most atoms will gain or lose elect rons to have 8 valence electrons (e in the highest energy
level)
Exceptions: H, He, Li, Be, B, and some atoms P and higher (by atomic number) on the periodic table
VSEPR theory ? Valence Shell Electron Pair Repulsion Theory
Repulsion between the sets of valence electrons surrounding an atom causes these sets to be oriented as
far apart as possible.
These sets are called ?regions of electron density”. These regions can be any of the following:
• Single bond (2e connecting 2 atoms)
• Double bond(4e connecting 2 atoms)
• Triple bond(6e connecting 2 atoms)
• Lone pair (unbonded pair) (2e alone on an atom)
When these regions spread out around the central atom in the molecule, they create certain ?shapes? to
the molecule. The shape of the molecule will depend on whether the regions are bonded pairs of
electrons or nonbonded pairs (lone pairs) of electrons and how many regions there are.
Shape
Bond Angles
linear
180o
trigonal planar
bent
tetrahedral
trigonal pyramidal
bent
trigonal bipyramidal
see-saw*
T-shaped*
linear*
octahedral
square pyramidal*
square planar
*You are not required to know these.
# Regions of
Electron Density
2
# Bonded Pairs
of Electrons
2
# Lone Pairs of
Electrons
0
Octet Rule
? octet rule ? most atoms will gain or lose
electrons to have 8 valence electrons (e- in
the highest energy level)
? Exceptions: H, He, Li, Be, B, and some atoms
P and higher on the periodic table
Why is an atom like Ca not stable
with 2 valence electrons?
VSEPR Theory
Valence
Shell
Electron
Pair
Repulsion
Repulsion between the sets of
valence-level electrons
surrounding an atom causes
these sets to be oriented as far
apart as possible.
Theory
How many would it have to lose to
have 8?
Regions of Electron Density
What is a Region of electron density?
? Single bond (2e- connecting 2 atoms)
? Double bond(4e- connecting 2 atoms)
? Triple bond(6e- connecting 2 atoms)
? Lone pair (unbonded pair) (2e- alone on an
atom)
TRIGONAL PLANAR
120o
3 Regions of Electron Density
3 Bonds
3 bonded pairs
of electrons
LINEAR
180o
2 Regions of Electron Density
2 Bonds
bonded pair
of electrons
bonded pair
of electrons
BENT
119o
3 Regions of Electron Density
2 Bonds & 1 Lone Pair
2 bonded pairs
of electrons
1 lone pair
of electrons
1
TETRAHEDRAL
109.5o
4 Regions of Electron Density
4 Bonds
TRIGONAL PYRAMIDAL
107o
4 Regions of Electron Density
3 Bonds & 1 Lone Pair
1 lone pair
of electrons
4 bonded pairs
of electrons
3 bonded pairs
of electrons
BENT
105o
4 Regions of Electron Density
2 Bonds & 2 Lone Pairs
TRIGONAL BIPYRAMIDAL
120o & 90o
5 Regions of Electron Density
5 Bonds
OCTAHEDRAL
90o
6 Regions of Electron Density
6 Bonds
SQUARE PLANAR
90o
6 Regions of Electron Density
4 Bonds & 2 Lone Pairs
SF6
ICl4-
2
Ms. K
Chemistry
Name______________________________ #______
Date____________________ Period_______
Lewis Structures Worksheet #1
1.) HI
5.) CH4
9.) BeF2
2.) H2S
6.) NH3
10.) BH3
3.) CH2Cl2
7.) H2O
11.) SO2
4.) O2
8.) Br2
12.) CO
13.) C6H5OH (ring structure)
14.) C2H5Cl
Rules for Drawing Small Molecules
Remember, each line represents a pair of electrons.
1.
2.
3.
4.
5.
6.
This is NOT what
happens on the
molecular level.
This is just our
method to draw the
correct Lewis
structure for any
small molecule.
Choose a central atom (If C is a choice, choose that.)
Bond all other atoms to it. (Unless it is a carbon chain.)
“Octet rule” everything except the small exceptions (H, He, Li, Be, and B)
Count the valence electrons to be used. (From the periodic table)
Count the valence electrons that have been drawn.
If the valence electrons to be used = the valence electrons you drew, then YOU ARE ALMOST DONE. Go
to Step 10.
7. If you have drawn too many electrons, then erase a pair of electrons from the central atom and erase a pair of
electrons from an attached atom. Create a double (triple) bond between them.
8. If you have not drawn enough electrons, then add a lone pair of electrons to the central atom. It must be an
exception to the octet rule and must be able to hold more than 8 electrons because of its size. (P and higher
by atomic number on the periodic table)
9. Go to Step 5.
10. Adjust for VSEPR Theory. Spread the regions of electron density out so that it reflects the appropriate
shape.
Shapes:
• linear – 2 regions of e- density
• trigonal planar – 3 regions of e- density
• tetrahedral – 4 regions of e- density
• trigonal pyramidal ? 4 regions of e- density (1 lone pair)
• bent or angular – 4 regions of e- density (2 lone pairs)
•
•
•
trigonal bipyramidal ? 5 regions of e- density
octahedral – 6 regions of e- density
square planar – 6 regions of e- density (2 lone pairs)
• others
Ms. K
Chemistry
Name______________________________ #______
Date____________________ Period_______
Lewis Structures Worksheet #1
1.) HI
5.) CH4
9.) BeF2
2.) H2S
6.) NH3
10.) BH3
3.) CH2Cl2
7.) H2O
11.) SO2
4.) O2
8.) Br2
12.) CO
13.) C6H5OH (ring structure)
14.) C2H5Cl
Ms. K
Chemistry
Name_____________________________ #_____
Date______________________ Period_________
Bonds, Resonance, and Ions
Bond Length, Bond Strength, and Bond Energy:
bond length – single bonds > double bonds > triple bonds
bond strength – triple bonds > double bonds > single bonds
bond energy – energy required to break a bond
triple bonds > double bonds > single bonds
Questions:
1. What bond would be the strongest: a single bond between C and C, a double bond between C and
C, or a triple bond between C and C?
2. Therefore, which one would take the most energy to break?
3. Which one would be the shortest?
Resonance:
Some molecules and ions cannot be represented adequately by a single Lewis Structure.
SO3
According to our structure, there are shorter bonds and stronger bonds. BUT…in the lab, they have
measured these bonds?.THEY ARE ALL THE SAME. WHY?
Resonance –
Polyatomic Ions:
Polyatomic ions Lewis Structures (Polyatomic ions are really small molecules that have gained or lost
electrons)
OHH3O+
NO3-
CN-
Ms. K
Chemistry
Name______________________________ #______
Date____________________ Period_______
Lewis Structures Worksheet #2
(be sure to include the resonance structures if there are any)
a. PCl3
b. SF6
c. NH3
d. BeH2
e. XeF4
f. PCl5
g. NH4+
h. H3O+
i. NO+
j. SO3
k. CO3-2