'15 Due
Date
Assignment
NAME: ___________________ per __
(check them off as you complete them)
Wed 9/23
___ Do WS 3.2 (#1-38)
+ Packet 3
Thur 9/24
___ Do WS 3.2 (#38-48)
___ Do WS 3.3 (#1-26)
Nomenclature,
Fri 9/25
___ Do WS 3.3 (#27-56)
___ Do WS 3.4 (#1-9)
Molecular Structure
miniquiz #1
Mon 9/28
___ Do WS 3.4 (all)
Tue 9/29
___ Do WS 3.5
Wed 9/30
___ Do Bonding Lab questions
Thur 10/1
___ Do WS 3.6
Bonding, &
miniquiz #2
Fri 10/2
___ Do WS 3.8 (left column only)
___ let's do WS 3.7 together
Mon 10/5
___ Do WS 3.8
___ start review sheet
* Mustard Day *
Tues 10/6
___
___
___
___
VSEPR
theory
Do WS 3.9
Do WS 3.10
QUIZ TODAY
Come to class with packets ready to be turned in, with the above assignments
in proper order (see below), in your correct folder (1/2 pt), with THIS PAGE as the cover page.
Make sure grade report is stapled in your folder (1/2 pt).
For 1/2 point, be sure not to turn in anything except for what's listed below:
Packet Order:
- assignment sheet
- WS 3.1 ~ 3.10
molybdenum (IV) bicarbonate
pentahydrate!
Mo(HCO3)4 • 5 H2O
IUPAC chemical
nomenclature
+ WS 3.1 - Nomenclature (inorganic)
Case 1
1.
2.
3.
Greek prefixes:
- Nonmetal + Nonmetal
First element retains its name.
Second element gets -ide ending.
Use Greek prefixes to identify the # of atoms
(ignore the 1st one if it's a '1')
Examples:
1-
6-
2-
7-
3-
8-
4-
9-
5-
101/2-
N2O = ______________________
SO3 = ______________________
Now Try These:
carbon dioxide
__________
BrF3 ________________________
diphosphorus pentasulfide __________
C S2 ________________________
Case 2 - ʻsʼ block metal + nonmetal (fixed charged metals)
1. Metal retains its name.
• make sure ions join to form a neutral compound •
2. Non-metal retains its ionic name.
Examples: Na+ + Cl- ---> ___________
Ca+2
+ Cl- ---> ___________
name = __________________________
name = __________________________
+
---> ___________
name = ___________________________
+
---> ___________
name = ___________________________
+
---> ___________
name = ___________________________
Now Try These:
lithium bromide
____________
Al2S 3___________________________
magnesium hydroxide____________
(NH4)2S O3 _______________________
aluminum acetate
Ba(NO3)2 _________________________
____________
WS 3.1, side 2
Case 3 - variable charge metal + nonmetal
1a. Metal gets its charge written w/ Roman # in ( ) after name (w/ new system), -or1b. Metal gets Latin name (w/ old system).
2. Non-metal retains its ionic name.
Examples: Fe+2
Fe+3
+ Cl- ---> ________
+ Cl- ---> ________
name = __________________ (
name = __________________ (
)
)
+
---> ________
name = ____________________
+
---> ________
name = ____________________
Now Try These:
manganese (IV) chloride __________
Cr(OH)3 ______________________
nickel (III) carbonate __________
CuCO3 ______________________
Case 4
- Acids (positive ion = "H+")
1. Drop the ending on the negative ion.
The -ate ending changes to -ic acid.
The -ite ending changes to -ous acid.
The -ide ending uses the prefix hydro- and the suffix -ic acid
Examples:
H2CrO4 = ______________________
HNO2 = ______________________
HCl = ______________________
Now Try These:
hydrobromic acid __________
H3PO3
perchloric acid ____________
_____________________
HI __________________________
“hydrates” when ionic substances have water molecules attached, theyʼre called “hydrates”
Now Try These:
MgSO4 • 3 H2O =
___________________
calcium phosphate tetrahydrate = _________________
NaC2H3O2 • 2 H20 = ________________
+WS 3.2 Nomenclature I
Name (yours) _______________
Covalent Compounds (case #1)
1. oxygen difluoride __________
4. SiF4
_____________________
2. sulfur hexafluoride __________
5. N2O
_____________________
3. silicon dioxide
6. NO2
_____________________
__________
Ionic Compounds (fixed charges - case #2)
7. sodium fluoride
__________
17. KCl
____________________
8. potassium sulfide
__________
18. Na2O
____________________
9. barium cyanide
__________
19. Mg3N2 ____________________
10. magnesium nitrate
__________
20. Na2C O3 ____________________
11. ammonium phosphate __________
21. NH4C 2H3O 2 __________________
12. calcium iodide
__________
22. RaCl2
____________________
13. sodium carbonate
__________
23. K3PO4
____________________
14. calcium chromate
__________
24. Mg3(PO4)2 ___________________
15. barium acetate
__________
25. Cs2 C O3 _____________________
16. lithium iodate
__________
26. Ca(HSO3)2 ___________________
27. copper(I) oxide
Ionic Compounds (variable charges - case #3)
__________
33. Cu2 S _______________________
28. copper(II) oxide
__________
34. FeO _______________________
29. ferric sulfate
__________
35. MoF2 _______________________
30. ferrous sulfate
__________
36. Cr(OH)2 ______________________
31. plumbic hydroxide
__________
37. Fe(HSO3 )2 ____________________
32. tungsten (VI) phosphate __________
38. PbS _______________________
39. phosphoric acid
Acids (case #4)
__________
43. H2 C 2 O 4 _____________________
40. carbonic acid
__________
44. HClO
41. hydrosulfuric acid
__________
45. H2S O4 _____________________
42. hypochlorous acid
__________
46. H2O
_____________________
_____________________
Special Case (hydrates)
47. zinc sulfate hexahydrate
48. Li2CO3 • 5 H20
_____________________
___________________________________
+ WS 3.3
Nomenclature II
Name (yours) ___________________
Hodgepodge (mix of covalent, ionic, & acids)
1. carbon dioxide
_________
27. BrF3
__________________
2. potassium cyanide
_________
28. Li2C 2O 4 __________________
3. selenium disulfide
_________
29. Fe3(PO4)2 __________________
4. potassium chlorate
_________
30. SCl4
5. nitrous acid
_________
31. KHCO3 __________________
6. zinc sulfate
_________
32. SnI2
__________________
7. aluminum acetate
_________
33. HF
hydrofluoric acid
8. copper(II) phosphate
_________
34. PO3
__________________
9. disilicon trioxide
_________
35. PO3 -3
10. chloric acid
_________
36. CaCO3 __________________
11. sodium chloride
_________
37. Fe(IO3)2 __________________
12. aluminum iodide
_________
38. CuCO3 copper (II) carbonate
13. barium cyanide
_________
39. CaF2
__________________
14. carbon disulfide
_________
40. HNO3
__________________
15. strontium nitrate
_________
41. (NH4)2S __________________
16. cuprous phosphate
Cu3PO4
42. SO3
__________________
17. phosphorous acid
_________
43. KNO3
__________________
__________________
__________________
18. potassium hydroxide _________
44. Sn3(PO4)2 __________________
19. bromine heptafluoride _________
45. MgS2O 3 __________________
20. lead (II) sulfide
_________
46. Ca2C
__________________
21. carbon monoxide
_________
47. H2 S
__________________
22. ammonium acetate
_________
48. CCl4
__________________
23. mercuric borate
_________
49. NaHSO3 __________________
24. calcium hydride
_________
50. NH4OH __________________
25. boron trichloride
_________
51. H3B O3 __________________
26. oxalic acid
_________
52. V(BrO3)5 __________________
53. MgSO4 · 7 H2O
______________________________
54. sodium acetate pentahydrate ____________
56. cobalt (III) nitrate tetrahydrate _______________
55. CuCl2 • 6 H2O ________________________________
+ WS 3.4 Bonding
1. What is the octet rule?
- ionic & covalent
2. What are some exceptions to the octet rule? __________________
How many electrons do these exceptions "desire"? ________
3. A _____________ bond is formed when a positive ion and a negative ion come together.
The metal will ( gain / lose ) electrons, giving it a ( positive / negative ) charge.
(circle correct
The nonmetal will ( gain / lose ) electrons, giving it a ( positive / negative ) charge.
word)
Determine how many electrons were transferred to make the following ionic compounds:
4. CaS ___
(Ans for #4-9): 1
5. CaCl2 ___
2
2
2
3
6. Al2O 3 ____
7. AlN ____
8. KBr ____
9. MgO ____
6
10. How many electrons are lost by group IIA elements (alkaline earths) during ionic bonding? ____
11. How many electrons are gained by group VIIA (halogens) elements during ionic bonding? ____
12. A covalent bond is formed when two nonmetals _____________ valence electrons.
13. A single covalent bond has how many electrons? _____
A double covalent bond has how many electrons? _____
A triple covalent bond has how many electrons? _____
did you check
the answer
bank?
Ans (IRO) for #13: 2, 4, 6
In the 1st box, write the Lewis dot structure for the compound.
In the 2nd box, write the molecular structure for the compound (showing the bonds).
F2
H2O
• •
••
F F
•
•
• •
••
• •
•
•
•
•
••
• •
• •
• •
F F
•
•
C H4
C O2
HCl
NH3
See example:
+ WS 3.5
More Lewis Structures
The steps for calculating how many bonds a molecular formula will have are:
1. Add the maximum # of valence e- each element can have (2 for H, 8 for most others)
2. Add the # of valence e- that each element has (adjust for charged molecules as needed)
3. Subtract #2 from #1. This will give you the # of electrons available for bonding.
4. Divide #3 in half. This is the # of bonds in the structure.
Draw the molecular structures for the following:
(we'll do the first 2 together in class)
N 2O
S O42-
HCO31-
CO
C 2H4
BrO3 -1
C O32-
HCN
O3
C H2S
HNO3
PBr3
+ WS 3.6
Polar Bonds & Electronegativity
1. An electronegativity difference between 0 to 0.4 is considered a __________________ bond.
2. An electronegativity difference between 0.5 to 1.9 is considered a __________________ bond.
3. An electronegativity difference between 2.0 and 4 is considered a __________________ bond.
Use the electronegativity values on your periodic table to determine the type of bonding in the
following molecules:
molecule
∆EN
bond type
4.
H-F
5.
H-Cl
6.
H-Br
7.
H-I
8. Which molecule in the table above is the most polar? _______
... the least polar? _______
9. For the following molecules, calculate ΔEN, state the bond type, and draw a dipole (if applicable)
molecule
ΔEN
bond type
ΔEN = (3.98 - 2.20) =
..
H-F
..
..
H-O-H
..
..
1.78
polar bond
..
..
O=C=O
..
..
..
..
..
..
..
Na+ Cl
..
..
..
Cl
.. - Cl
..
H
H
..N
H
10. Who developed the electronegativity scale?
_______________________
11. What are the units for electronegativity anyway?!?! ____________________
+WS 3.7
# of e- pairs
(regions)
VSEPR
e - geometry
bond angle: _______
bond angle: _______
bond angle: _______
bonding
diagram
molecular
geometry
(pull the
ropes)
polarity
example
+WS 3.8
Molecular Geometry
molecule
Lewis
structure
e- regions on
central atom
egeometry
C H2 S
H2 O 2
NF3
C O2
COCl2
O3
S O2
CO
H3 O +
HF
ClO3 -
..
..
..
..
O
.. - O = O
..
3
trigonal
planar
molecular
geometry
P or NP
polarity
+WS 3.9 Exceptions to the Octet Rule
You already know (hopefully) that most elements prefer an _________ of electrons. Boron is one exception, which
only wants ____. There are other exceptions, including elements which can have 10 or even 12 ____________
electrons around them! These extra electrons tend to form bonds in empty, unused “d” orbitals. For example,
______________’s valence shell configuration is 3s23p3.
more bonds.
Ans IRO:
cesium
If its vacant “d” orbitals are used, this element can form
octet
phosphorus
six
valence
Determine the Lewis structure and geometry of the following exceptions to the octet rule:
Lewis structure
PF5
XeF4
XeF2
SF6
I3-
SF4
IF5
ICl4-
electronic
geometry
molecular
geometry
+WS 3.10
Review
The following random compounds need either names or formulas. Help them discover their identity!
1. Cs3P
___________________
3. platinum (IV) iodate _______
6. HBrO3 ________
2. SiCl4
___________________
4. calcium nitrite ________
7. sulfur trioxide _______
9. (NH4)4C ___________________
11. mercury (II) oxalate ________
5. stannous chloride _______
8. Fe(NO3)2 ___________________
10. hydrochloric acid _________
12. CoCO3 _________________
13. CO ___________
14. How many electrons were transferred to make AlCl3? _____
Write the Lewis dot structures, geometry, and polarity for the following:
15. O2
16. HF
17. BCl3
electronic geometry?
electronic geometry?
electronic geometry?
molecular geometry?
molecular geometry?
polar? ___
18. SO3
polar? ___
19. NO3 1 -
electronic geometry?
molecular geometry?
molecular geometry?
molecular geometry?
polar? ___
polar? ___
20. CH4 O
electronic geometry?
electronic geometry?
molecular geometry?
polar? ___
Classify the bond type as non polar, polar, or ionic by calculating ∆EN:
21. KBr
∆EN = ________
bond type = _________
22. CO
∆EN = ________
bond type = __________
23. SiI2
∆EN = ________
bond type = _________
24. Draw all possible resonnance structures for sulphur trioxide:
polar? ___
Bonding Lab
Part 1:
*
Name:
____________________
Start the tutorial
Partner:
_____________________
Ionic Compounds
The ionic bond occurs between a _____________ atom and a ______________ atom. When a
metal atom loses electrons, it becomes a [ positively / negatively ] charged ion. When a nonmetal
atom gains electrons, it becomes a [ positively / negatively ] charged ion.
1. At your lab station, place a pinch of NaCl (table salt) directly on the black lab
counter, and observe it through a magnifying lens, At right, sketch what you see:
Compare these crystals with the large NaCl crystal on display (under the spotlight) What evidence is there that the ions in these crystals are arranged in the specific pattern
discussed in the tutorial?
2. Go watch the demo your instructor has set-up. A test tube filled with NaCl is being heated over a
very hot flame. When the NaCl starts melting, observe whatʼs going on inside the test tube:
3 When the salt is nearly all melted, is the unmelted (solid) NaCl floating at the top or sunk to the
bottom of the liquid (molten) NaCl? ___________
4. Once melted, watch what happens as the molten salt is poured into a beaker. As the molten salt
cools, what do you observe?
5. Predict what would happen if the solidified NaCl were bent:
6. Observe what happens when the solid (frozen) salt is bent.
What occurred? ______________________ Was your prediction correct? _____________
*
Return to your lab station & resume the tutorial
to learn why ionic substances behave this way.
Part 2:
*
Covalent Compounds - Network Covalent
Bonding Lab side 2
7. There are 2 types of covalent substances: network covalent and molecular. The example most
often used for a network covalent solid is diamond. Diamond is made up of carbon atoms covalently
bonded to other carbon atoms which in turn are covalently bonded to more carbon atoms. We do
not have a class set of diamonds for you to try to melt, but we have plenty of silica. Silica is SiO2
(silicon dioxide -- also known as “sand!”). Take a look at the sand under the lens. Aside from having
a slightly different color, it looks very much like the salt. What do you think would happen if you
repeated part 1, but we put sand in the test tube rather than salt? _______________
Why do you think this would happen?
*
Return to the tutorial to learn why network covalent
substances (like diamond and sand) behave the way they do.
Why were we not able to melt the sand in class? _____________________________________
Part 3:
Covalent Compounds - Molecules
8. Take a large test tube with some wax at the bottom, and clamp it to the ring stand. Heat it about 3
cm over a very small (2-3 cm tall) cool flame. Once the wax is mostly melted, turn off the
burner. Does solid wax float or sink in the molten wax? _________
9. Try bending the wax sheet in the plastic bag. What happens? ________________________
What do you think would happen if you struck the wax with a hammer? ____________________
What would happen if you hit a piece of ice with a hammer? _______________________
*
Resume the tutorial to learn more about
molecular solids and why they behave this way.
What does IMF stand for? _____________________________________
How do IMFʼs differ from ionic and covalent bonds?
____________________________________________________________________
Part 4:
Metallic Bonds
10. Take a new penny (minted after 1982) and file away about 1 cm
of the copper along an edge to expose the silvery-gray zinc inside.
11. Light a Bunsen burner, and adjust to a hot flame. Place 400 mL beaker
beneath the flame. Use tongs to hold penny vertically in hottest part of flame.
Be sure to hold penny by the top with the exposed zinc edge at the bottom.
After 15~25 seconds, you should notice a change! Let the molten zinc drip
straight into the beaker below (you may need to give it a little shake). Set the
tongs and whatʼs left of the penny down to cool. What did you observe?
filed area
1¢
12. After 30 seconds, the zinc in the beaker should be cool. Pick it up
burner
and bend it. Observations?
13. Clean-up & throw away penny.
*
Resume the tutorial to learn more about
metallic substances & why they behave this way.
beaker
*
Bonding Lab Follow-Up Questions:
If you need to review the tutorial on bonding, you can find the powerpoint on the class
website. Click on packet #3, look under “optional”.
1. Ionic, metallic, covalent. Which is the strongest? _____________ the weakest? ___________
2. Salt melts at 801 ˚C, sand melts at 1700 ˚C, and zinc melts at 420 ˚C.
Do your answers for #1 above agree with these melting points?
___________
Explain:
3. How do the strengths of bonds between molecules (intermolecular forces, or “IMFʼs”) compare to
the bond strengths in #1 above? ______________________________________________
4. Search online for the melting points (˚C) of the following common covalent molecules:
_____
sugar (sucrose):
butane (C4H 10):
_____
wax (paraffin):
_____
caffeine:
_____
5. In each pair, circle the substance which you think would have the higher melting point.
Explain your choice:
a) K / KF
____________________________________________________________
b) MgBr2 / diamond
c) CH4 / AlCl3
____________________________________________________
__________________________________________________________
6. All metals are solids at room temperature except one. What metal is it? _______ What does
that tell you about this metalʼs melting point? ______________________________________
7. Is NaCl(solid) more or less dense than NaCl(liquid)? _______ How could you tell? ___________
Is wax(s) more or less dense than wax(l)? _______ How could you tell? _____________
Is H2 O (s) more or less dense than H2 O (l)? _______ How do you know? _____________
Which behavior is more common: a solid floating or sinking in its own liquid? _______________
H
8. Consider the diagram at right.
When ice melts, which bonds are
starting to break? A / B / both / neither
And when the water boils?
A / B / both / neither
O
O
H
H
H
O
H
A
H
B
9. Some covalent substances, such as diamond and sand, have incredibly high melting points.
Other covalent substances, such as sugar and butane, have much lower melting points. Explain why.
Hint: your explanation should include words such as: molecules, bonds, IMFʼs, etc...