Download unit 5 hw packet - District 196 e

Honors Chemistry
Name ____________________
Unit #5B Plan – Bonding Concepts and Carbon Chemistry
I)
II)
Assignments / Handouts
 Lewis structures and naming H-carbons
 Electronegativity chart
 Molecular shape guide
 Chemical bonding and Lewis structures
 Bonding concepts ws
 Text problems / resonance structure asst
 Naming / Drawing H-carbons ws
 Polymers info and ws
 Polymers (uses, recycling codes, refineries)
 Text – chapter 7-8 study guide
 Polymer lab
 Building isomers activity
 More TBA
Text Support
 Chap 7
 Chap 8
 Chap 22
 Chap 23
pg 192-213
pg 220-255
pg 760-788
pg 796-829
(info – pg 4)
(info – pg 3)
(notes / info – pg 2)
(notes – pg 5-6)
(required – pg 7-8)
(required – pg 8 --- attach to back of packet)
(required – pg 9-10)
(required – pg 11-14)
(info – pg 15-16)
(required – pg 17-18)
ionic and metallic bonding
covalent bonding
hydrocarbon compounds
functional groups
III)
Learning Targets (I can…)
A)
Use an electronegativity chart to predict whether individual bonds in molecules are ionic,
polar covalent or nonpolar covalent.
B)
Predict shapes of molecules based upon their location in the periodic table.
C)
Use molecular shape and electronegativity differences between atoms to determine
whether an overall molecule is polar or nonpolar.
D)
Draw electron dot diagrams (Lewis structures) to show the bonding of atoms.
E)
Identify molecules in which hydrogen bonding occurs.
F)
Build, draw, and name hydrocarbons.
G)
Build, draw, and identify isomers of hydrocarbons.
H)
Use my knowledge of carbon chemistry to solve polymerization and biochemistry
problems
I)
Identify and analyze current issues involving carbon chemistry.
IV)
Tentative Schedule
Mon
11/23 Polymer lab
Tue
11/24 Polymer lab wrap-up
Ionic, polar covalent, and nonpolar covalent bonding
Wed
11/25 Lewis (electron dot) structures
Mon
11/30 Molecular shapes
Distinguishing between polar and nonpolar
Tue
12/1
Intro to H-carbons
Wed
12/2
Building isomers and naming H-carbons
Thu
12/3
Functional groups
Mon
12/7
Review
Tue
12/8
Exam – bonding concepts and H-carbons
Hints for Drawing Lewis (Electron Dot) Structures
1. Add the number of valence electrons in each atom to determine the total number of
valence electrons. (For polyatomic anions, add one electron for each unit of negative
charge. For polyatomic cations, subtract one electron for each unit of positive charge.)
2. Put electrons around each atom. Start with the 4 base electrons before pairing electrons
up. (It is usually best to start with the atom that has the fewest valence electrons --excluding hydrogen.)
3. Atoms bond when electrons need to be shared to complete an octet around each atom.
4. The total number of electrons in Lewis structure should equal the total number of
valence electrons in the molecule (or ion).
 Hydrogen is “full” with 2 electrons. All other atoms need an octet of electrons to be
stable.
 A “shortage” of electrons usually means double or triple bonds are needed.
 Oh yes, there are exceptions.
Rules for Naming Hydrocarbons
1.
2.
3.
4.
Count carbons in longest chain to get “root” name.
Check for single, double and triple bonds (or OH group). Use appropriate suffix.
Number carbons starting at the end nearest the double/triple bond or OH group.
Check for “branches”. (If there are no double/triple bonds or OH groups, number
carbons starting at end nearest the outermost “branch”.) Put in alphabetical order.
5. Put commas between consecutive numbers; dashes between numbers and words.
Alkane Roots (single bonds)
Methane
CH4
Ethane
C2H6
Propane
C3H8
Butane
C4H10
Pentane
C5H12
Hexane
C6H14
Heptane
C7H16
Octane
C8H18
Nonane
C9H20
Decane
C10H22
Branches
methyl
ethyl
propyl
butyl
(etc.)
fluoro
chloro
bromo
iodo
Functional Groups
Alkene --- double bond
Alkyne --- triple bond
Alcohol --- OH group
# of Branches
2 = di
3 = tri
4 = tetra
More Functional Groups (R or R’refers to carbon chain)
Ether
R-O-R’
Aldehyde
R-(C=O)-H
Ketone
R-(C=O)-R’
Carboxylic acid R-(C=O)-OH
Ester
R-(C=O)-O-R’
Amine
R-N(H2) or R-N(R’2) or R-N(R’H)
Aromatics
R-(C6H6)
-CH3
-C2H5
-C3H7
-C4H9
-F
-Cl
-Br
-I
Chemical Bonding Notes
Conditions for Chemical Bonding
 Atoms must be physically close to each other

Atoms must attract each other
o Attractive forces > repulsive forces

Energy removal
o Energy is released (given off) as bonds formed
o Energy needs to be absorbed (added) to break bonds
Bond Length –distance between nuclei of 2 atoms when their attractive forces exceed their
repulsive forces by a maximum amount
Chemical Bond – electrons of 2 atoms are simultaneously attracted to the nuclei of both atoms
 Ionic Bond – unequal sharing of electron pairs between 2 nuclei

Nonpolar Covalent Bond – equal sharing of electron pairs between 2 nuclei

Polar Covalent Bond – somewhat equal (or unequal) sharing of electron pairs
between 2 nuclei
Lewis (Electron Dot) Structures Intro
Draw Lewis structures for each of the following covalent compounds. Your Hints for Drawing
Lewis Structures will be helpful.
H2O
CF4
CH2F2
Cl2
HF
NH3
CH4
NH4+1
OH-1
Unit #5B – Bonding Concepts
Name_____________________________
1. Use your electronegativity chart to predict if the following molecules have ionic, polar
covalent or nonpolar covalent bonds. (Look at the individual bonds --- not the overall
symmetry.)
A. NaCl
G. CF4
B. BaCl2
H. CO2
C. OCl2
I. O2
D. H2O
J. NH3
E. CH4
K. H2S
F. MgO
L. Cl2
2. Draw Lewis (electron dot) structures for the following molecules or ions.
A. NH3
H. H2S
B. Br2
I. NH4+
C. ClF
J. BH4-
D. PH3
K. CH3OH
E. SiH4
L. O2
F. C2H6
M. N2
G. C2H4
N. C2H2
3. Predict the shapes of the following molecules.
A. NH3
D. OF2
B. Br2
E. N2
C. BCl3
F. CH2Cl2
4. Are the following molecules polar or nonpolar? (Consider overall symmetry.)
A. NH3
D. OF2
B. Br2
E. N2
C. BCl3
F. CH2Cl2
5. For each of the following molecules:
 Draw a Lewis (electron dot) structure
 Determine the shape of the molecule
 Determine whether the molecule is polar or nonpolar
A. H2O
D. NF3
B. CH2F2
E. F2
C. CF4
F. CH3Cl
Text pg 256-258 --- Do problems 51, 52, 53, 68, 79, 82.
Text pg 234 --- Identify 3 chemical formulas (on pg 234) that are best represented by
resonance structures. Draw all the resonance structures for each of these formulas.
Attach to back of this homework packet.
Unit 5B Book Assignment (Chapters 7-8)
1. What is the connection between valence electrons and the octet rule?
2. Cations are formed when metals ___________ electrons.
3. Anions are formed when nonmetals ___________electrons.
4. What is a formula unit?
5. The coordination number for both Na+ and Cl- in NaCl is 6. What does this mean?
6. Metals are good conductors of electricity. Why?
7. What is a metallic bond?
8. Define alloy.
9. Give 3 examples of a molecular (covalent) compound.
10. The “representative unit” of an ionic compound is called a formula unit. Why?
11. The “representative unit” of a molecular (covalent) compound is called a
molecule. Why?
12. Draw the electron dot structure (Lewis structure) for each of the following
molecular compounds. (Include unshared pairs of electrons.)
CH4
NH3
H 2O
13. Which molecule contains: (choose from C2H6, C2H4, C2H2)
a. Only single bonds
b. One double bond
c. One triple bond
14. What is a coordinate covalent bond? Give an example.
15. List 3 exceptions to the octet rule. Give an example of each.
16. How is bond dissociation energy and covalent bond strength related?
17. There are 2 resonance structures for ozone? Does ozone oscillate back and forth
between these 2 structures? Explain.
18. A single bond has ________ sigma bonds and _________ pi bonds.
19. A double bond has ________ sigma bonds and _________ pi bonds.
20. A triple bond has ________ sigma bonds and _________ pi bonds.
21. What does VSEPR stand for?
22. What is orbital hybridization?
23. What is the difference between a polar covalent and nonpolar covalent bond?
24. List 3 types of intermolecular forces.
25. What is a network solid. Give an example.