1.
Development of the Atomic Theory
300 B.C. -
: ( page 160 - 174 )
Democritus: matter made up of tiny indivisible particles, “_____________”
1
1809 -
John Dalton - Billiard ball model - atom was _________, ___________, _____________ sphere
Limitations: Could not explain why atoms combined in certain ratios.The atom is not the smallest particle
1897 - J.J. Thomsom - Plum Pudding model - atom was _____________ charged with ___________ charged
_______________ embedded in the sphere. If Thomsom’s atom was bombarded with alpha particles ( which are _________ charged ) it would be expected that the particles would
________________________________________.
1911 -
Ernest Rutherford - gold foil experiment - most alpha particles go __________ ____________; a few are ________________ at large angles - Nuclear atom model : atom mostly _________ ____________ nucleus is ______________ charged ; extremely ___________ but takes up most of the _______ of the atom.
Major flaw in model is that the electron in motion should give off _____________ that is lose energy and eventually spiral into the nucleus.( Page 174 )

Complete pg 166 (3)
2
_____________________________________________________________________________________
_____________________________________________________________________
1932 - James Chadwick: Alpha particle bombardment, explains why mass of nucleus does not equal mass of protons present, “neutrons”
1900 - Max Plank(pg 169) - atoms absorb or release energy in discrete packages called ______________and in terms of light are referred to as __________________; evidence ____________ radiation and the ____________________ effect.
- Blackbody – perfectly black object that does not reflect any light and emits various forms of light as a result of its high temperature.
Examples of Blackbodies:
-
-
-
1
2
Stove Element-Red Hot
Light Bulb Filament-White Hot
In 1900, Max Planck developed the formula, E=hf (where E is Energy and h is a constant (Planck’s constant) and f is
- the frequency of light). Thus he was stating that light emitted from a blackbody was not continuous but in multiples of a small quantity of energy.
Planck stated that the energy from a blackbody is _______________; restricted to whole # multiples of certain energy
What does the different sizes of the circles represent ? _________________________________________
In the above diagram what does the energy threshold represent ? __________________________
Is the energy of the electrons produced by the photoelectric effect dependent of the light intensity ?
______________________________________________________________________________________
The photoelectric effect was discovered by German Heinrich Hertz in 1888. When light, particularly light of high energy shines on the surface of a metal, ___________ are emitted from the surface. Max Planck (1858-1947) discovered different metals had different __________ when they released electrons. This threshold was based on
_____________________________of the light not the _____________________________.
1905 - Einstein - unknown Swiss clerk named Albert Einstein published three papers. The first paper gave his explanation of the photoelectric effect. He proposed light was made up of small bundles of light energy, called ___________ or ________________ and that the E = hf (h is known as ____________ constant). The second paper explained Brownian motion (gas behaviour). The third paper was about the now famous link between matter and energy, E = mc 2 . considered light to be __________________________ and explained the photoelectric effect by stating that _____________________________________________________.

Electromagnetic spectrum :
1.
2.
3.
4.
Radiation with shortest wavelength is ______________ and the one with longest is __________
Microwaves have a _____________ wavelength than x - rays
Infrared has a _____________ frequency than visible light
Visible light represents a ______________ spectrum with the colours ROYGBIV arranged in order of
____________ wavelength and _______________ energy.
Complete page 173 (4)
_________________________________________________________________________
_________________________________________________________________________
Line Spectrum(pg 175)
- distinguish between bright-line emission spectra and adsorption or dark-line spectra
410 nm 434 nm
1 nm = 1 x 10 -9 m = “a billionth of a meter”
486 nm 656 nm
____________________________________________________________________________________________
____________________________________________________________________________________________
_____________________________________________________________________________________________
Line spectrum of Hydrogen consists of 4 lines _________,__________,________and _________
3

4
-
Day 2 1913 - Niels Bohr - Solar System Model
Bohrs theory had 2 major postulates (pg 176) :
1.__________________________________________________________________________________________
2.__________________________________________________________________________________________
-
Bohrs atomic Model :
If energy is added to the H atom in the form of electrical current e’s are ______________ to higher energy levels as they fall to lower levels energy is emitted in the form of light called
__________________ spectra ; for the H atom specifics jumps result in certain colours : 3
2 nd ________, 4 th to 2 nd ________, 5 th to 2 nd ______6 th to 2 nd _____________ rd to when electrons absorb certain wavelengths of light the atoms emission spectrum is called an
_________________ spectrum.
- Jumps to the 1 st level result in _________________ light and jumps to the 3 rd level _________ light determined. Actually, the e- is so small that it has both exact orbits. Another problem is when an electron changes energy levels during the emission of atomic spectra.
Weakness of Bohrs Model (pg 177) works for only 1 electron systems such as H and does not explain the emission spectra of other elements e.g mercury which has many additional lines of yellow and neon which has many lines of red. These additional lines suggest that there are _________________- associated with the main shells
Bohr visualized the e as a _____________by which its exact location and momentum could be
and
properties so we cannot specify
An electron can never be “in-between” energy levels so it is clear that this model has limitations and a new model needs to be generated.

5
-
-
-
-
-
-
-
1924 - Louis de Broglie
- page 199 - suggested a dual nature, the electron has both ________ like and __________ like characteristics ; extremely __________ objects have a significant wavelength ---------> streams of moving electrons produced diffraction patterns support this idea
The idea that matter may be represented by a wave was a new idea and gave rise to the equation:
 = h
mf
Complete page 180 (4,5,8,9 )
1926 -
-
Erwin Schrodinger(pg 199)
In 1926, after Austrian Erwin Schrodinger heard of de Broglie’s “electron wave” it occurred to him that this idea could be used to solve the electron location problem in an atom. Schrodinger proposed using wave
(quantum) mechanics – a new branch of mathematics - and applying it to the behaviour of the electron. The result was our present orbital theory of the electron. Orbitals, electron clouds, take on certain shapes or sublevels
(s-spherical, p-hourglass, d-cloverleaf). The energy level, n, must equal the number of sublevels. Within a sublevel the number of orbitals will change. For the s – sublevel there is 1 orientation, for the p-sublevel, 3; for the d-sublevel, 5.
- Quantum mechanical Model - electron has _______-like characteristics.
1927 - Werner Heisenberg (pg 200)
-
-
-
Uncertainty principle - impossible to know both the __________ and _______________ of an electron at any specific time since measuring one will affect the other. probability regions where e’s may be found are called ______________ ( 90% of the time ) watch video on www.teachersdomain.org
: Light Particles Acting Like Waves: The
Uncertainty Principle
Quantum Numbers (pg 181 )
1.
-
2.
-
Principal Quantum # ( n ) represents the ____________ energy levels i.e. n=1,2,3 etc ; the maximum e’s in any main shell is represented by the formula ______
Secondary Quantum # ( l ) identifies the _____________ of the orbital thus identifying ___________ in each main shell l =0,1,2, n-1 subshells are identified by letters : ____, ______, _______ and _______ ( some people don’t forget ) l=0 ______ subshell l=1 ________ subshell l=2 _______ l=3 _____________
If n=1 , l= _____ thus _______ subshell called the _________ subshell
If n= 2, l= _____, _______ thus a _______ subshell and a ________ subshell and the pattern continues; the d subshell is not found until the _________ energy level and the f not until the ____ energy level see page 201 ----> the shape of the s orbital is _______________ and the p orbital is ___________
General formula for # of subshells for each main shell is ________________

-
Apartment Analogy
floors will represent the energy levels
apartments will represent the subshells
-
on the first floor there is a 1s apartment ; on the 2 nd
on the the 3 rd on the 4 th
floor there is 2s and a 2p apartment
floor the apartments are _________, ___________, and ______________
floor the apartments are _________, _______, __________, and ______________.
Complete page 182 ( 3,4,5 )
-
-
-
Day 3 3. Magnetic Quantum # (m l
) (pg 182) identifies the ________________ of the orbital in space m l
ranges from _________ to ____________ if l = 0 _________ subshell m l
= ___________ i.e. _________ orientation
6
If l=1 __________ subshell m l = _ i.e m l
=-1 (px) m l =
0(py) m l
_________
= +1 (pz ) i.e. _________ orientations identified as px, py and pz
- l =2 identifies the ______ subshell. How many orientations? _______

-
If l=3. How many orientations? _______ since m= __________________________________ general formula for the # of orbitals for each main shell is _________________
3d x 2 - y 2
2p y
3d yz
2p z
3d xz
+
1s
2p x
3p x
2s
3s
3d xy
3p z
4s
3d
3p y z 2
Questions
1) What would be the subshell designation (eg. 1s) for the following: a) n=2, l=1 ______ b) n=3, l=2 _________ c) n=4, l=0 ________ d) n=5, l=3 ________
Apartment analogy - rooms represent the orbital ; s apartment has ______ room , p ______, d ________, f _______________
7

-
-
-
3.
-
Spin quantum # (s ) ( Pg 183) as e’s orbit they spin on their own axis clockwise spin is taken as +1/2 and counterclockwise as -1/2. An orbital can hold a maximum of ______ e’s the first will have a _____ value and the 2 if an electron is spinning in one direction about an arbitrary axis, it will generate a magnetic field with a certain polarity if it is spinning in the opposite direction, the magnetic field will be reversed nd one __________ opposite spins and attracting magnetic fields might account for the tendency of 2 electrons to co-exist in a limited region of space
the two electrons in a filled orbital will take a position where the attracting force if the magnetic field is equal to the repulsive force owing to the similar negative charge
8
-
individuals represent the electrons ; first person in has a positive attitude 2 nd has a negative attitude
in apartments rooms are filled 1 person at a time until each room is half full then they pair up.
Complete page 184 ( 1,3 )
Summary :
Symbol n l m s
Name
___________
___________
__________
__________
Values
1,2,3 ... n
0,1,2 ... n-1
-
+ 1
0...+
/
2
,1 /
2
Meaning fixes the size of orbital value and energy fixes the __________ of the orbital
l=0 spherical ____ orbital
1 two lobed ____ orbital
2 four lobed ____ orbital
3 more complex ____ orbital defines ___________of orbitals in a magnetic field defines direction of spin or rotation of electron on it axis
Distinguish clearly between an orbital and a orbit ( pg 185 )

9
2 . Complete the following:
1. The first one has been done for you.
Shell Principal
Quantum(n)
K
L
M
N
1
___
_____
_____
Secondary
Quantum #(l)
0
___
____
_____
_____
_____
_____
_____
_____
_____
Subshells s
____
____
_____
_____
_____
_____
_____
_____
_____
# of
Orbitals
1
_____
_____
_____
_____
_____
_____
_____
_____
_____
Day 4 Quantum Rules
1)
2)
The maximum # of electrons in any main shell is _________ ; n=1 ___ , n=2 ___ , n=3 ___ n=4 ____ .
The maximum # of subshells in any main shell is________________ eg. n=1 --- s subshell n=2 --- ________, ____________ n=3 ________,__________,__________ n=4 ________,__________,__________,___________
The # of orbitals for each main shell is _____________ 3) e.g. n=1, l=0 --- 1 orbital (m=0) n=2, l =_______ m =_________
l = _______ m =_________
Pauli Exclusion Principal ( pg 188 ) 4)
- no 2 e, s in the same atom can have the same set of _________ quantum numbers. e.g. H - 1 electron can be identified by the quantum #' s :
n=____ , l=____ , m=____ , s=____
-
6)
-in He, the last electron can be identified by:
n=____ , l=____ , m=____ , s=____
5)
-
Hund's rule successive last electrons enter _________ orbitals until all possible orbitals contain _________ then they pair up e.g. carbon 6. __________________________________ try N ________________________________________________
Aufbau Rule electrons enter subshells of __________________ energy first e.g. 3 _____ is filled before 4 _______

10
Below are energy level diagrams . Complete for the elements Barium, Iodine and Platinum. The first e is represented by an arrow up and the second by an arrow down. (pg 187)

11

8.
9.
12
The diagram above represents an energy level diagram . It represents the order in which electrons enter the atom. The circles with the electrons shown with arrows also are known as schematic diagrams.
The order of filling can also be shown as illustrated in diagram page 188.
Complete a energy level diagram for Zn and for Br. See above.

13
Assignment Quantum rules Name
1)
2)
Another name for the region in space where the electron may be is a(n):__________
Suppose the electron was in the 1s orbital. The principal quantum number is ________ ; the orbital has a shape which is _________________ .
3) Consider an electron in the 3px orbital. This is one of the _______________orientations for a p orbital and it can contain a maximum of ____________ electrons.
4) If scientists discovered an atom with orbital designation of "g", it would possibly have ______ orientations.
5) An e may be in the following orbitals ( 7s, 6p, 4d, 5s,4f ) . The one of lowest energy is _____________ . The one with the greatest number of possible orientations is the _______________
The orbital with the highest energy would be the (7s, 6p, 5f, 4p , 4d) ____________ .
8)
9)
6)
7)
10)
11)
Wolfgang Pauli noted that an orbital can have 0, 1, 2 but never _________ electrons.
A 1s subshell has _________ orbitals and can accommodate a maximum of _________ electrons.
A 4f subshell has _________ orbitals and can accommodate _____________ electrons.
What subshell is identified by the Quantum #'s n = 3, l = 2 ________________
For the Neon atom, which electron i.e. # 1 to 10 would be identified by the #'s n = 2, l = 1, m = +1, s = + 1/2 ____________________
12)
13)
How many electrons in an atom could be identified by the quantum #'s: n = 3, l = 2 ________________
How many electrons can be identified by the quantum #'s : n=2, l=1, m=-1, __________________
14) How many electrons in an atom could be identified by the quantum #'s: n=3, m=0, s=-1/2
15) Why are the following quantum #’s not possible for an atom ? a.
b.
n=3, l=3,m=-2 ___________________________ n=3, l=2,m=-3 ___________________________
16) In the space below draw an energy level diagram for Al and Ni and show how the electrons are placed in the orbitals

Day 5/6 -
1
2
H -
He - 1s
Electron configurations
1s
2
1 - the first one represents the _________ shell; the s represents the _____ shell and the exponent 1 represents ___electron
Complete the following and also include orbital representations: ( pg. 192 )
14
3
Li -
4
Be -
5
B -
6
C -
7
N -
8
O -
9
F -
10
Ne –
11
Na –
15
P -
20
Ca –
3) The filling on the 3d subshell, with slightly higher 4s already occupied, represents an unusual situation.
Between Sc and Zn, the 3d sublevel is filling.

Type of Transition Situation
- 4d is filling with 5s already occupied
- 5d is filling with the 6s already occupied
- 4f is filling with 5s, 5p and 6s occupied
- 5f is filling with the 6s, 6p, and 7s full
Elements Involved
4d
5d
4f
5f
57
39
La,
58
Y -
72
90
48
Cd
Hf -
Ce -
71
Th
80
Lu
Hg
4)
5)
Consider the electron configuration for
26
Fe - 1s 2 2s 2 2p
The structure for
24 stable.
Cr is: ____________instead of 3d 4
6 3s
4s 2
2 3p 6 4s 2 3d 6 rewritten as: 1s 2 2s 2 2p
Rearrange the order so that subshells with the same principal quantum #'s are together:
1s 2 2s 2 2p
2 2p 6
6 3s
3s 2
2 3p
3p
6 3d
6 3d 5
6 4s 2
3)
Fe
Now remove 3 electrons starting from the outermost energy level:
3+ 1s 2 2s
orbital representations :
6 3s 2 3p 6 3d
6)
2)
1)
35
42
Mo -
3)
82
-
Pb -
-
Next let's try some positive or negative ions: eg. Fe 3+
Steps:
1)
2)
The structure for
29
Cu is ____________ instead of 3d numbers are together)
9 4s 2 .
Give the electron configuration for: (Note: always rewrite at the end so that all principal quantum
Br -
-
First give the electron configuration for
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6
26
Fe:
6 4s
because a half-filled sub - shell is especially
2
15
.

K
Br
Zn
7.
4) Se 2-
Ex. 2 : try O -2
O - 1s 2 2s 2 2p 4 ----------> O
Try the following:
-2 - __________________________________________
1) Cr 3+
2) Ni 2+
3) P -3
Where possible we try to use shorthand notation see page 193
[ ] ____
[ ] _________________
[ ] ___________________
16
8.
Electrons in atoms can be identified by using the 4 quantum #’s. Consider the quantum #’s n=4,l=2,m=2,s=+1/2
Identify an electron in the element Cd . Which electron is it ?
Step 1 Write the e configuration
Step 2
Step 3
Do an orbital representation
Circle the 4 th energy level; the d subshell ; the fifth orbital ; and point to the 1st electron
Step 4 Count the electrons until you can to the identified one.
It is ___________ electron.

17
Assignment
Part I:
Determine the element whose outermost electron is being defined by the following quantum numbers.
Element
1. _____________
2. _____________
3. _____________
4. _____________
5. _____________
6. _____________
7. _____________
8. _____________
9. _____________
10. _____________
11. _____________
12. _____________
13. _____________
14. _____________
15. _____________
16. _____________
Quantum Number Code n = 1, l = 0, m n = 2, l = 1, m l n = 2, l = 1, m l n = 2, l = 1, m l n = 3, l = 1, m l l
= 0, m s
= - ½
= -1, m
= 0, m
= 1, m
= -1, m s
= + ½ s
= - ½ s
= + ½ n = 3, l = 1, m n = 3, l = 1, m l l n = 4, l = 0, m l
= 0, m
= 1, m s
= - ½ s
= + ½ s
= - ½ s
= + ½ n = 3, l = 2, m l n = 3, l = 2, m l n = 3, l = 2, m l
= 0, m
= -2, m
= -2, m
= 0, m s s
= + ½
= - ½ s
= + ½ n = 3, l = 2, m n = 3, l = 2, m n = 4, l = 1, m l l l
= 1, m
= 2, m n = 4, l = 1, m l n = 4, l = 1, m l
= 0, m
1. n = 2, l = 1, m
2. n = 1, l = 1, m
3. n = 8, l = 7, m
4. n = 1, l = 0, m
5. n = 3, l = 2, m
6. n = 4, l = 3, m s
= - ½ s
= + ½
= -1, m
= 1, m s
= + ½ s
= - ½ s
= + ½
Part II:
Which of the following sets of quantum numbers are not allowed in the hydrogen atom? For the sets of quantum numbers that are incorrect, state what is wrong.
7. n = 2, l = -1, m l
Part III: l l l
√ if not allowed Explanation of error if applicable
= -1 _____ ________________________________________________________ l
= 0 _____ ________________________________________________________
= 6
= 2
_____ ________________________________________________________
_____ ________________________________________________________ l
= 2 _____ ________________________________________________________ l
= 4 _____ ________________________________________________________
= 1 _____ ________________________________________________________
What is the maximum number of electrons in an atom that can have these quantum numbers?
1.
2.
5.
6.
n = 4 n = 5, m
3.
n = 5, m s l
= +1
= + ½
4.
n = 3, l = 2 n = 2, l = 1 n = 0, l = 0, m
7.
n = 2, l = 1, m l
10.
n = 1, l = 0, m l l
8.
n = 3
9.
n = 2, l = 2
= 0
= -1, m
= 0 s
= - ½
______
______
______
______
______
______
______
______
______
______

18
WS Orbital Notation Name__________________
Let arrows reflect electrons and fill in the orbitals for atoms listed to the left. Then write the electron configuration and core notation for that element.
1. Calcium
3d 3d 3d 3d 3d
4p x
4p y
4p z
4s
3p x
3p y
3p z
3s
2p
2s
1s
2. Nitrogen x
2p y
2p z
2p x
2p y
2p z
Electron configuration _______________________________________
2s
1s
Electron configuration _______________________________________
_______________________________________
Core notation _____________________________________________
3. Silicon
4s
3d 3d 3d 3d 3d
4p x
4p y
4p z
3p x
3p y
3p z
3s
_______________________________________
Core notation _____________________________________________
4s
3p
3s x
3p y
3p z
3d 3d 3d 3d 3d
4p x
4p y
4p z
2p
2s
1s x
2p y
2p z
Electron configuration _______________________________________
_______________________________________
Core notation _____________________________________________

19
4. Manganese
4s
3p x
3p y
3p z
3s
3d 3d 3d 3d 3d
4p x
4p y
4p z
2p
2s
1s
5. Vanadium x
2p y
2p z
2p x
2p y
2p z
Electron configuration _______________________________________
_______________________________________
Core notation _____________________________________________
4s
3p
3s x
3p y
3p z
3d 3d 3d 3d 3d
4p x
4p y
4p z
2s
1s
Electron configuration _______________________________________
_______________________________________
Core notation _____________________________________________
6. Selenium
4s
3d 3d 3d 3d 3d
4p x
4p y
4p z
3p x
3p y
3p z
3s
2p x
2p y
2p z
2s
1s
Electron configuration _______________________________________
_______________________________________
Core notation _____________________________________________
Fill in the names of the orbitals on the lines below the circles. Then fill in the orbital notation, electron configuration and core notation using this type of chart.

8. Iodine
20
Electron configuration____________________________
_______________________________________________________
Core Notation _________________________________________
Electron configuration______________________
__________________________________
__________________________________
Core Notation _________________________________________

9. Chromium *** Be Careful
21
Electron configuration____________________________________
______________________________________
_________________________________________
10. Copper ***
Core Notation _________________________________________
Complete pg 191 (3)
Electron configuration_______________________________
__________________________________
__________________________________
Core Notation __________________________________

Reading Assignment ; Make notes on Laser technology page 203, spectrometers pg 205, x-rays pg
206, Cat Scan and MRI ( 206)
22

23
Independent Research Name ____________________________________
1.
Assess the benefits to society of technologies that are based on the principles of atomic and molecular structure( e.g. magnetic resonance imaging (MRI), X-ray crystallography, nuclear energy, medical applications of spectroscopy and mass spectrometry). Consider the issue : In medicine radioisotopes are bonded with chemical compounds to form radioactive tracers, which are then injected into the patients bloodstream. The radiation emitted by the tracers allows doctors to obtain images of organ systems, facilitating the early and accurate diagnosis of disease. However to avoid radioactive contaminants, care must be taken in the storage, use and disposal of this material. Answer the questions that follow :
1. How does infrared spectroscopy aid in criminal investigation ?
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
2.
How has the use of X-ray crystallography and mass spectrometry advanced our understanding of atomic and
molecular structure ?
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
3. What social benefits are associated with the above advances?
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
4. Electron spin resonance(ESR) is an analytical technique that is based on the spin of the electron. State some examples of the uses of ESR in at least 2 different areas?
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
5. How is the MRI technique similar to and different from ESR and provide examples of the usefulness of
MRI results?
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
_______________________________________________________________________

-
-
-
-
-
-
-
Day 7 : Characteristics of s,p,d and f block elements ( pg 188 )
- Specific subshells represent specific groups in the periodic table
- group 1 s 1 , group 2 _____________; groups 13 to 18 _________to___________; transition
-
-
-
-
24 metals ____to______; lathanides _______to_______actinides ______to_____. for main group elements the last number of their group # represents the number of ___________ electrons group 1 the alkali metals have a configuration that ends in ________ thus very ___________; must be stored under ________________________ group 2 the alkaline-earth metals end in ______________; still quite reactive but not as reactive as group
1 main group elements in group 1,2 and 13 tend to ___________ electrons main group elements from 14-18 tend to ______________ electrons ; e.g. group 16 tend to gain _____ e’s group 18 elements the noble gases are extremely _____________ since ________ shell is full; configuration ends in _____________ groups 3 to 12 represent the ________________ elements e.g. Ni ends in _____s ____ _____d ____ since size _________ down the periodic table and ___________ across the table the most reactive metal would be _____________________; most reactive non-metal is ________________ reactivity for metals _____________ across the table small atoms have very________ ionization energies and very _________ electron affinities tend to
________ electrons; the most reactive halogen is _________________ the most stable noble gas is ___________________
Complete page 194 ( 6,9 ) Page 197 ( 1-6 ) 9, 13, Page 202 (2,5)
Complete page 219 ( 1-19 ) pg 220 ( 2,3,6,9,12,13,15-18 )
Make up a short note on superconductivity (pg201)
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________

Q U A N T U M N U M B E R P R A C T I C E
25
1. Rank the following orbitals in the H atom in order of increasing energy: 3s, 2s, 2p, 4s, 3p, 1s, and 3d.
__________________________________________________________________
2. How many orbitals in an atom can have the following quantum number or designation? a) 3p ___ b) 4p ___ e) 5d ___ f) 5f ___ c) 4p x
___ d) 6d ___ g) n = 5 ___ h) 7s ___
3. Answer the following questions as a summary quiz on the chapter. a) The quantum number n describes the _______ of an atomic orbital. b) The shape of an atomic orbital is given by the quantum number ____. c) A photon of orange light has _____ (less or more) energy than a photon of yellow light. d) The maximum number of orbitals that may be associated with the set of quantum numbers n=4 and l=3 is ____. e) The maximum number of orbitals that may be associated with the quantum number set n=3, l=2, and m l
= -2 is ___. f) When n=5, the possible values of l are ______. g) The maximum number of orbitals that can be assigned to the n=4 shell is ____.
WS Quantum Numbers and Orbitals
1.
Indicate which of the following orbital destinations are possible. a. 7s b. 1p c. 5d d. 2d e. 4f f. 5g g. 7i
2.
Without referring to a text, periodic table or handout, deduce the maximum number of electrons that can occupy an: a. s orbital _____ b. the subshell of p orbitals _______ c. the subshell of d orbitals ______ d. the subshell of f orbitals_______ e. the subshell of g orbitals_______
3.
Explain why there are 10 members of each d transition metal series. ____________________________
___________________________________________________________________________________
4.
Explain why there are 14 members of each f inner-transition metal series. ________________________
___________________________________________________________________________________
5.
Indicate which of the following electron configurations is ruled out by the Pauli exclusion principle. a. 1s 2 2s 2 2p 2p 6 3s 3 c. 1s 2 2s 2 2p 6 3p 6 4s 2 3d 2p 6 3s 2 3p 6 7 2s 2 3s 2 12 2s 2
6.
Explain why the following ground-state electron configurations are not possible: b. 1s 2 b. 1s 2 2p 7 3s 2 3p 8 d. 1s 2 d. 1s 2 2s 2 a. 1s 2 2s 3 2p 3
7.
Give two examples of:
2s 2 3 3s 6 c. 1s a.
an atom with a half-filled subshell
2 2s 2 2p 2p 6 3s 2 3p 1 4s 2 3d 14 b.
an atom with a completely filled outer shell c.
an atom with its outer electrons occupying a half-filled subshell and a filled subshell.

8.
Fill in the blanks with the correct response: a.
The number of orbitals with the quantum numbers b.
The number of valence electrons in the outermost p subshell of a sulfur atom is _________. c.
The number of unpaired electrons in a Mn d.
The subshell with the quantum numbers e.
The m
=3,
=2,
=6,
l
values for a d orbital are ________________________.
= -2, m
= 2, m
= 2, m
for the shell with
for the shell with j.
The number of orbitals with
n.
The number of electrons with
=3 and
=2 are _________.
=4 are _________.
=3 is _________.
=1 is _________. k.
The maximum number of electrons with quantum numbers with l.
When m.
When
=2,
can be _________.
=2, the possible values for m
=4,
l
are _________.
=1 is _________. o.
The quantum number that characterizes the angular shape of an atomic orbital is _________. p.
The subshell with q.
The lowest value of n for which a d subshell can occur is n=_________.
9.
Which sets of quantum numbers are unacceptable? l l l
=3 and
=0, m s
= -1, m
= -2, m s
=3,
=2 and
=2 is _________.
l
= 0 is _________.
=1is designated as the __________ subshell.
= +½
= -½ s
= +½
2+
=4,
ion is _________. f.
The allowed values of g.
The allowed values of h.
The number of unpaired electrons in the cobalt atom is _________. i.
The number of orbitals in a shell with
=3 and
=2 is _________. a.
b.
c.
10.
Identify the group of elements on the periodic table which have the following ground state electron configuration: np 3 ____ b. ns 1 _____ c. ns 2 np 6 _______ d. ns 2 ______ a. ns 2
Chemical Bonding (224)
26
1.
2.
Distinguish between ionic and covalent bonding
Ionic –
Covalent –
Lewis structures - for atoms the # of dots is equal to the last digit in the group number ; complete below for period 3 elements

27 ii. ii.
4.
5.
Steps To The Writing Of Lewis Structures (pg 225)
1. From the molecular formula given ,determine how many valence electrons are available from all atoms present in the molecular formula. Calculate the total.
2. Add one valence electron to the total for each negative charge on the species if any, and subtract one valence electron for every positive charge on the species if any
3. Decide which atom in the molecular formula will be the central atom and arrange all other atoms around it in as symmetrical a way as possible. The atom with the highest valency is considered the central atom.
If there are more than one atom possible, the atom with the lowest electronegativity is usually considered to be the central atom.
Attach all atoms to the central atom with a pair of valence electrons from the pool (total from step 1.)
Distribute the rest of the valence electrons from the pool so that each atom obeys the octet rule. Start with the outside atoms. Additional electrons go around the central atom.
6. If there are electrons left over in the pool look for multiple bonding by shifting atoms so that the left over valence electrons can be incorporated
7. If there are not enough electrons left over in the pool look for multiple bonding by converting pairs of non-bonding electrons into bonding electrons
8.
9.
10.
If steps 6 or 7 cannot be accomplished assume that this is a case that is an exception to the octet rule.
Draw the Lewis structure showing the covalent bonds with a short line
Watch out for oxy-acids the H and O are attached together
Ex.1 SF
2 i.
S - 6 valence e = s
F - 2 x( 7 val. e = s )
Total val. e = s = 20e’s
F- S- F
Complete
Ex.2 CO
3
2- i. C is bonded to 3 O’s ii. C - ___ val. e
3 O - 3 x _______ val. e’s
O C O
___ outside e’s
O
val. e’s
•
•
•

Complete the following:
1)
3)
SO
SiF
4
2-
4
5) SO
2
Cl
2
7) HCN
9) N
2
O
11) SO
2
13) ICl
2
-
15) XeO
4
17) BrF
4
+
Complete page 229 ( 10,12 ) page 230 (4)
2) HClO
2
4) CNO -
6) SO
3
8) NH
4
+
10) ClF
3
12) XeF
16) SF
4
2
14) N
2
F
2
18) COCl
2
28

29
Day 8 A. VSEPR Theory (pg242)
1.
Molecular shapes are using the VSEPR ( valence shell electron pair repulsion) theory which states that because electrons repel, molecules adjust their shapes so that the valence - electron pairs are as far apart as possible to minimize electron repulsion Note: unshared electrons take up more room than the bonding electrons. In terms of decreasing repulsion : LP – LP > LP – BP > BP - BP
Electrostatic Repulsion
- Data have shown that bond angles for atoms in molecules with p orbitals in the outer energy level do not conform to the expected 90  separation of an x, y, z axis orientation. This variation can be expressed by electrostatic repulsion between valence electron charge clouds or by the concept of hybridization. This valence energy level electron pair repulsion model is sometimes called the VSEPR model.
Electrostatic repulsion uses as its basis the fact that like charges will orient themselves in such a way as to diminish the repulsion between them.

Determining the Shape of a molecule (pg 243)
The best way to determine the architecture of a molecule is to:
1.
2.
3.
4.
Determine what the central atom is.
Draw the Lewis structure of the molecule.
Determine the number of bonding pairs and lone pairs around the central atom.
Refer to the following chart and determine the shape of the molecule.
VSEPR CHART A = central atom B = bonding pair E = lone pair
# of bonded electron pairs # of lone pairs VSEPR Formula
2
2
2
2
3
3
3
4
4
4
5
5
6
30
Shape
0 AB
2
Linear
1
2
3
0
1
2
0
1
2
0
1
AB
AB
2
0
AB
2
AB
3
E
3
AB
AB
AB
AB
2
E
E
2
3
3
4
4
E
E
2
E
Bent
Bent
Linear trigonal planar pyramidal
T-shaped
Tetrahedral unsymmetrical
AB
4
AB
5
AB
5
AB
6
E
2
E tetrahedral ( see - saw )
squareplanar
Trigonal bipyramidal
Square pyramidal
Octahedral
Refer back to pg 26 and predict the shapes of the various molecules and ions.
Watch tutorial on www.teachersdomain.org
: molecular shape
Complete pg 246(2,3) pg 247(9) pg 249(10,11) pg 250 ( 1-3 )

Molecular Geometry – Summary
Valence
Electron
Pairs
2
Electron pair
Geometry
Terminal
Atoms
2
3
3
2
Lone Pairs Molecular Geometry
Bond
Angle(s)
0
0
1
31
4 0
4
3 1
2 2

Valence
Electron
Pairs
Electron pair
Geometry
Terminal
Atoms
5
Lone Pairs
0
4 1
5
3 2
2 3
6 0
6
5 1
Molecular Geometry
32
Bond
Angle(s)
4 2

3.
-
1. CCl
8. OF
2
4
Molecular Architecture Worksheet complete the chart below for the following :
15. H
2
Se
Molecule or ion
2. ClO
9. CO
3
2
-
2-
16. HClO
2
3. XeF
10. NH
17. CH
2
3
3
Lewis Structure
NH
2
4. XeOF
11. NH
4
4
+
E.P.A.
5. HCN
12. BrF
3
Molecular shape
6. ClO
3
13.O
3
7. XeO
14. PCl
4
6
-
Sketch
33

4.
34
-
What bond angles are associated with the following shapes : linear__________, trigonal planar______________, bent or angular________________, pyramidal______________, tetrahedral_________,trigonal bypyramidal______/_________, octahedral_______________.
Outline the contributions of Dr. Bader (pg 249)
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________
Outline the contributions of Dr. Ronald Gillespie (pg 242)
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
_________________________________________________________________________
_________________________________________________________________________

Summary Table :
Total # of
Groups of e-
Electron Pair
Geometry
2 linear
3 ______________
4
5
6 tetrahedral octahedral
Approximate
Bond Angle
_______
_____________
120o (in plane) &
90o (above & below)
90o
# of
Bonding
Directions
(# of X)
2
# of
Lone Pairs
(# of E)
Geometry Name
(VSEPR class)
0
3
2
4
3
2
5
4
3
2
6
5
4
3
2
0
1
0
1
2
0
1
2
3
0
1
2
3
4 linear
(AB2)
___________
(AB3)
T-shaped
(AB3E2) linear
(AB2E3) octahedral
(AB6) square pyramidal
(AB5E)
________
(AB4E2)
_________
(AB2E) tetrahedral
(AB4)
_____________
(AB3E) bent
(____) trigonal bipyramidal
(AB5)
___________
(AB4E)
T-shaped
(AB3E3) linear
(AB2E4)
Shape
35
Examples
BeH2,
CO2
BF3, NO3
–
SO2
CH4
NH3
H2O
PCl5
SF4
ClF3
XeF2
SF6
BrF5
XeF4

 A covalent bond is described as two __________________orbitals so that they can be occupied by a shared pair of electrons of ____________________spin
 The bond results in a _______________________in the energy of the atoms forming the bond
Hydrogen gas, H
2
 A simpler method of showing the orbitals that are involved in bonding is the orbital diagram
 Outer ‘s’ and ‘p’ orbitals are represented by squares and their electron populations by arrows
 A rectangle is drawn around the orbitals involved in bonding to indicate the bond (shared pair of electrons
H

1s

1s shared electrons in overlapping 1’s’ orbitals

H
Water, H
2
O
 In the water molecule the 1’s’ orbital for each of the hydrogen atoms overlaps two of the 2’p’ orbitals of the oxygen atom
36

 The orbital diagram for water is:
O

H
H


2s

2p


1s


1s
 In order for carbon to have half-filled orbitals (and form 4 bonds), it is suggested that one of the 2s electrons is promoted to the vacant 2p orbital
 When the 2s orbital and the three 2p orbitals of a carbon atom are mathematically averaged or hybridized, the result is four equivalent orbitals which point towards the corners of a tetrahedron
 The four equivalent orbitals are called sp 3
 The carbon atom is then able to bond with four hydrogen atoms to form the methane molecule, CH
4
37

38
MULTIPLE BONDS
 Multiple bonds can be explained by the presence of two types of bonds: o sigma bond  , - the end-to-end overlap of s orbitals, p orbitals, hybrid orbitals or any combination of these o pi bond  , - two orbitals overlap side by side that are represented by the second and/or third lines in the structural diagrams for double and triple bonds
 an orbital diagram for nitrogen gas, N
2
, can also be used to show multiple bonds:
N

N


2s





2p
bond
 bonds



39
-
3.
Day 9 Types of bonds Polarity of Molecules ( page 251 )(pg 254)
1.
Distinguish between ionic, polar and non-polar covalent bonds
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
2.
In classifying bonds use the following rule : if electronegativity difference is> 1.7 = ionic; between 0.5 and 1.7 = polar covalent; between 0 and 0.5 = mostly covalent
Complete page 253 (1,2) molecules are said to be polar if the distribution of electric charge on an object is not uniform, such that one region has an excess of negative charge while another has an excess of positive charge. The polarity of a molecule is measured by calculating the dipole moment of a molecule.
1)
2)
3)
For a molecule to be a dipole, two conditions must be met: a) b) ionic character in at least one of the bonds within the molecule molecular shape must permit a net charge displacement
Symmetrical molecules are non-polar in most cases as long as the atoms attached to the central atom are identical.
Rules all diatomic molecules consisting of unlike atoms have a dipole moment triatomic molecules can either be linear (no dipole moment) or bent (dipole moment) binary tetratomic molecules can be:
i) planar (no dipole moment) e.g. BF ii) pyramidal (dipole moment) e.g. NH
3
3
Dipole Moments of Various Molecules; complete chart
Molecules Electronegativity
Difference
Dipole Moment
HCl
CO
BeF
CO
OF
H
2
BF
2
2
2
O
3
1.08
1.91
0
0
0.30
1.85
0

Dipoles have a definite direction that can be represented by a vector, showing both magnitude and direction.(pg 254)
1) BeF
2
F← Be → F ; addition of vectors 0 →
←
non polar (dipoles cancel out)
2) H
2
O
3) BF
3
4) NH
3
40
- a) d) g)
Referring to your summary page for molecular shapes, you should be able to determine the shapes that would be non-polar. Complete the following but assume that the atoms attached to the central atom are identical.
Shape linear
Polar or non-polar bent trigonal planar pyramidal t - shape see-saw tetrahedral trigonal bipyramidal square planar
Square based pyramidal
Octahedral
Classify the following as polar or nonpolar
BeI
CO
SF
6
2
2
________ b)
_______ e)
_______ h)
CF
4
_______
BF
2
Cl ______
SCl
2
______ c) f)
CH
CH
2
3
O ________
Cl ________

a.
Draw 2 structures for each of the following one should be polar the other non-polar :
PCl
3
F
2 b.
SF
4
Cl
2
Complete page 256 ( 8,10,11 ) ( 1,2,4 )
41

42

1.
Molecule or Ion
(a) No. of valence e’s
(b) Lewis structure
(c) Approximate bond angle(s)
(d) Electron group arrangement
(e) Polar or non-polar molecule?
(f) Geometry name
(g) Hybridization
(1) OF
2
(2) H
2
CO (3) NO
2
+
Ion: Not applicable
(4) BF
3
(5) SbF
5
______________________________________________________________________________
2. For each of the molecules below fill in the indicated items in the chart. The central atoms are underlined.
Molecule
(a) No. of valence e - > s
(b) Lewis structure
(c) Approximate bond angle(s)
(d) Electron group arrangement
(e) Polar or non-polar molecule?
(f) Geometry name
(g) Hybridization
(1) SO
2
(2) HBF
2
(3) XeF
4
(4) CH
2
Cl
2
(5) NF
3
43

44
Day 10 Experiment 1 Molecular Models
1. Purpose Names _________________________
The purpose of this lab exercise is to use Lewis Structures and Valence-Shell Electron Pair Repulsion (VSEPR, commonly pronounced “vesper”) theory to predict shapes of small molecules and polyatomic ions and to construct models of compounds.
2. Procedure
Molecular Species List:
BrF
PF
N
2
5
5
3
2-
(Br – 6 holes) (F – 1 hole)
(P – 5 holes) (F- 1 hole)
O (N – 4 holes) ( O – Red ( 6 holes )
CO
OCl
2
(C – black (4holes) (O – Red ( 6 holes )
(O – 6 holes) (Cl -1 hole)
NCl
ClF
SiCl
XeF
SF
4
SF
ICl
3
3
4
4
(N – 4 holes) (Cl -1 hole)
(Cl – 5 holes) (F – 1 hole)
(Si – 4holes) (Cl -1 hole)
( Xe – 6 holes ) (F – 1 hole)
6
2
1-
( S – 5 holes) (F – 1 hole)
( S -6 holes ) (F – 1 hole)
( I – 5 holes ) (F – 1 hole)
1. For each of the species listed above, do the following and record your answers on the worksheets provided.
2.
3.
4.
Determine the total number of valence electrons in the molecule.
Determine the Lewis Structure of the molecule.
Determine the number of bonding electron pairs and non-bonding electron pairs
5. Determine the electronic and molecular geometry around each centralized atom in the molecule from the
Lewis Structure and VSEPR theory. Predict if polar or non-polar . For ions state NA.
6. Construct an exact geometric model of the molecule. (You may work with a partner to construct models.)
7.
Sketch an exact 3-D representation of the molecule from the model. (See VSEPR handout).
3. Report
Simply finish recording all of the requested information on the worksheets and turn them in before leaving lab.
There is no formal lab report required for this lab.

Experiment 1 Lab Sheet
1. Molecule =
Valence electrons =
Bonding pairs =
Non-bonding pairs =
Electronic geometry =
Molecular geometry =
Polarity =
3-D drawing:
2. Molecule =
Valence electrons =
Bonding pairs =
Non-bonding pairs =
Electronic geometry =
Molecular geometry =
Polarity =
3-D drawing:
3. Molecule =
Valence electrons =
Bonding pairs =
Non-bonding pairs =
Electronic geometry =
Molecular geometry =
Polarity =
3-D drawing:
Names __________________________________
Lewis Structure:
Lewis Structure:
Lewis Structure:
45

4.
5.
Molecule =
Valence electrons =
Bonding pairs =
Non-bonding pairs =
Electronic geometry =
Molecular geometry =
Polarity =
3-D drawing:
6.
Molecule =
Valence electrons =
Bonding pairs =
Non-bonding pairs =
Electronic geometry =
Molecular geometry =
Polarity =
3-D drawing:
Molecule =
Valence electrons =
Bonding pairs =
Non-bonding pairs =
Electronic geometry =
Molecular geometry =
Polarity =
3-D drawing:
Lewis Structure:
Lewis Structure:
Lewis Structure:
46

7.
8.
9.
Molecule =
Valence electrons =
Bonding pairs =
Non-bonding pairs =
Electronic geometry =
Molecular geometry =
Polarity =
3-D drawing:
Molecule =
Valence electrons =
Bonding pairs =
Non-bonding pairs =
Electronic geometry =
Molecular geometry =
Polarity =
3-D drawing:
Molecule =
Valence electrons =
Bonding pairs =
Non-bonding pairs =
Electronic geometry =
Molecular geometry =
Polarity =
3-D drawing:
Lewis Structure:
Lewis Structure:
Lewis Structure:
47

10. Molecule =
Valence electrons =
Bonding pairs =
Non-bonding pairs =
Electronic geometry =
Molecular geometry =
Polarity =
3-D drawing:
11. Molecule =
Valence electrons =
Bonding pairs =
Non-bonding pairs =
Electronic geometry =
Molecular geometry =
Polarity =
3-D drawing:
12. Molecule =
Valence electrons =
Bonding pairs =
Non-bonding pairs =
Electronic geometry =
Molecular geometry =
Polarity =
3-D drawing:
Lewis Structure:
Lewis Structure:
Lewis Structure:
48

-
Day 11 Intermolecular forces ( pg 257 )
49
1.
Classification of Solids: Based on Intra and Intermolecular Forces
based on macroscopic properties that arise from the arrangement of their component particles.
-
-
- intra ----> between atoms in a molecule e.g. covalent and polar covalent --> very strong inter ----> between molecules e.g. london dispersions , dipole-dipole forces and hydrogen bonding --> weak forces
A.
Crystalline Solids – organized pattern of particle arrangement. a.
Atomic b.
Molecular c.
Network
B.
d.
Ionic e.
Metallic
Amorphous Solids – indistinct shapes because their particle arrangements lack order. a.
b.
Glass
Polymers
1) Atomic Solids intra ----> between atoms in a molecule e.g. covalent and polar covalent --> very strong inter ----> between molecules e.g. london dispersions , dipole-dipole forces and hydrogen bonding --> weak forces consider Helium (F.P.= -272.2
o C) and Argon (F.P.= -189.2
o C) force between atoms is called "London Dispersion Forces" ( pg 258 ) consider diagram below ; At a particular instant, the electron distribution in atom 1 becomes unsymmetrical, with more electrons on the _______ side of atom. At that instant the left side will have a slight charge with respect to the right side( _____________________ Dipole); The charge on the right side of 1 the electrons in atom 2 giving the two atoms the same charge distribution (___________________ Dipole). The positive end of atom 1 and the negative end of atom 2 attract each other. An instant later, as electrons shift again, the left side of atom 1 may develop a slight negative charge, causing the charge distribution in the neighbouring molecule to shift in a similar manner.
Considered to be flucuatating dipoles an easy way to view this force is that the atomic nucleus of an atom attracts the e = s of neighboring atoms( see page 258)
Note: This is a very weak force ( 1-10 kJ/mol ) and is directly related to the ______________present. Since
Argon has 18e and He only 2e; Ar has a higher M.P. > since London Dispersion Forces are stronger.

2) Molecular Solids
50
A) Non-Polar Molecules
-
- consider the table below:
Molecule
F
Cl
Br
2
2
2
Total # of Electrons
18
34
70
106
State at Room Temp gas; B.P.= -188 gas; B.P.= -34.6
o C liquid; B.P.= 58.8
o C
solid; B.P.= 184.4
o C o C I
2 in non-polar molecules London Dispersion Forces also operate
Magnitude of L.D. forces between molecules is determined by: i) # of electrons ii) molecular size(# of atoms per molecule) iii)
Account for the increase in B.P. from Fluorine to Iodine
molecular shape
________________________________________________________________________________
- consider the table below:
Molecule
C
C
C
CH
4
2
3
H
H
H
4
6
8
10
# of Electrons
10
18
26
34
B.P.
- 162
- 87
- 45
- 0.5
-
-
Questions:
1) Account for the increase in boiling point _____________________________________________________
2) Compare Cl
2
and C
4
H
10
. Account for the difference in B.P. ______________________________________
3) Two isomers of butane have the boiling points 0 o C and - 12 a) C - C - C - C b) C - C - C
\ o C.
C
Which B.P. would you assign to each? Account. ___________________________________
___________________________________________________________________________
B.
Polar Molecules e.g. HCl, HI, ICl, CH
3
Cl forces between molecules is dipole-dipole or hydrogen bonding

1) Dipole - Dipole
51
- in polar molecules such as HCl the positive side of one molecule attracts the negative side of a neighbouring molecule ( force generally weak 3-4 kJ/mol )
consider the following:
Molecule
ClF
BrF
CH
CH
3
3
F
Cl
electroneg.
difference
_______
_______
___and ___
___and ___
# of Electrons
26
44
18
26
B.P.
- 101
- 20
- 78
- 24
- in comparing ClF and BrF you can see that ___________ has a greater dipole moment and a greater # of electrons thus dipole-dipole forces are greater as well as London Dispersion Forces thus the boiling point is greater
Note: - it is difficult to make generalizations about the relative strengths of the intermolecular forces unless we restrict ourselves to comparing molecules of either similar size and shape or similar polarity and shape
1.
2.
Questions :
Account for the difference in B.P. between ClF and CH
3
Cl._______________________________________
Account for the difference in B.P. between C
2
H
6 and CH
3
F
______________________________________________________________________________________
3. Account for the difference in B.P. between ClF and C
3
H
8
______________________________________________________________________________________
Note: the term Van der Waals Forces is subdivided into the forces:
1) London Dispersion 2) Dipole - Dipole

Complete page 260 (1-5)
Day 12
2) Hydrogen Bonding ( pg 261 )
52
-
-
- when hydrogen is bonded to a highly electronegative element such as F, N, or O it can serve as a bonding bridge to an atom in another molecule which is relatively negative and contains lone electron pairs
10 x the energy of V.W. forces (10-40 kJ/mol)
1/10 the energy of covalent bonds e.g. HF

-
53
# of Electrons
10
18
36
54
B.P.
+ 19.4
- 83.7
- 67
- 35
According to the trend HF should have a B.P. less than - 83.7. Why is its B.P. the highest?
________________________________________________________________________________
-consider the data: Complete chart
Compound
H
H
H
H
2
2
2
2
O
S
Se
Te
Electronegativity of Group
VI Element
# of Electrons B.P.
100
- 61.6
-42
4.0
- consider the data:
Compound Electronegativity of
HF
HCl
HBr
HI
Halogen
4.0
3.0
2.8
2.5
- in H
2
O there are ________________ lone pairs of electrons, which can attract the positive hydrogen of another water molecule as shown in the diagram on the next page. When the motion of the water molecules is restricted as in ice, hydrogen bonding between molecules is directed by the tetrahedral shape of the water molecules, leaving significant _______________________ holes as in the diagram page 263. It is hydrogen bonding that explains the unusually lower density of ice than liquid water
(.917g/cm 3 at 0 o C). When ice melts, this open structure collapses to give a more closely packed pattern considering the graph below why would it be predicted that H
2
O would have the lowest B.P. of the series.
______________________________________________________________________________________
Explain waters abnormal behaviour._________________________________________________________
Why does H
2
O have a higher B.P. than HF? __________________________________________________
What are some other properties of water which can be accounted for by hydrogen bonding ?( page 262 )
______________________________________________________________________________________
_______________________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________
__________________________________________________________________________

54
11.2
- make a brief note on the work of Dr LeRoy pg 265
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________

55
Question
1) Dimethyl ether (CH
3
)
2
O has B.P. of -24.9 whereas its isomer ethanol CH
Account for the difference.
3
CH
2
OH has B.P. of 78 o C.
- C - O - C - ------> ether and - C - C - O - H -------> alcohol
______________________________________________________________________________________
Complete page 264 ( 9, 11 ) page 266 ( 1,3,4,5,6,8,15)

Day 13 Characteristics of Molecular Substances (page 270)
56
1)
2)
3)
4)
5)
6)
7)
Majority are at room temperature
Solids are usually e.g. mothballs, wax
conductors of electricity and heat since _____________________
melting points
Intramolecular forces are great i.e. requires a large amount of energy to decompose the molecule
Non-polar molecular solids are in water
Polar molecular substances are in water and if acidic in nature form _____________ solutions
Worksheet
1.
Draw the hydrogen bonds that exist in the following diagrams. a) b) H H CH
3
H
:
O
:
H
.
.
O
.
.
H
H
.
.
.
.
O
H
H
N
. .
H3C
H
3
H
H N:
C
H N:
H
2.
Is a hydrogen bond really a bond? ________. If not, then what is it? ____________________________
________________________________________________________________.
3.
Why do hydrogen bonds only exist between molecules that contain H and molecules that have fluorine, oxygen and/or nitrogen? ________________________________________________________________
4. Rank the following types of IMFA from weakest force of attraction to strongest force of attraction.
(ion-ion attractions, Ldf, hydrogen bonding, dipole-dipole)
_____________________________________________________________________________________
5.
How do the boiling points and melting points of molecules with hydrogen bonds compare to the boiling points and melting points of molecules without hydrogen bond?_________________________________
_____________________________________________________________________________________
6. In which of the following substances are molecules capable of hydrogen bonding? Circle all correct choices. If not indicate major force. a) CH
3
F b) CH
3
NH
2 c) CH
3
COOH d) CH
3
CH
2
CHO

7.
Classify the following as non-polar molecular or polar molecular and outline the major force to overcome to melt the substance once it is in a solid state
Substance Classification Intermolecular
Attractions (major)
CF
4
C
3
H
8
HCNO
HF
CS
2
PCl
3
SF
6
SO
2
57
I)
4)
5)
-
-
II)
1)
2)
3)
-
- ionic compounds are built up from ions packed together in a compact orderly array in such a way that charged ions are close together and ions of the charge are separated. The type of packing depends on the relative sizes and the charges on the ions. bonding: is non-directional - ions arrange themselves in an ionic solid to maximize attractions between opposite ions and to minimize repulsions between like charged ions.
Properties
Very _________________ but brittle
_______________ melting points
________________ conductors of electricity in the solid state but - __________ conductors in the molten state majority are ___________________ in water
_______________ vapour pressures the hardness and high melting points are due to the strong __________________ attractions between the closely packed ions (400 – 4000 kJ/mol) the brittleness of ionic crystals can be explained as follows: when stress is applied a layer of ions is forced out of place in such a way that ions of ______________________ charges are next to each other; this leads to strong _____________________ forces and a fracture in the crystal

→
☺ ☻ ☺ ☻ ☺ ☻ ☺ ☻
☻
☺ ☻ ☺ ☻ ☺ ☻ ☺
☺ ☻ ☺ ☻ ☺ ☻ ☺ ☻
→
☺ ☻ ☺ ☻ ☺ ☻ ☺ ☻
☻
☺ ☻ ☺ ☻ ☺ ☻ ☺
☺ ☻ ☺ ☻ ☺ ☻ ☺ ☻
→
☺ ☻ ☺ ☻ ☺ ☻ ☺ ☻
58
- - - - - - - - - - - - - - - - - - - -
☻
☺ ☻ ☺ ☻ ☺ ☻ ☺
☺ ☻ ☺ ☻ ☺ ☻ ☺ ☻
- for ionic crystals to conduct electricity they must be in liquid or molten state where ions are
_________________ to move ( ________________ ions)
- the solubility in water is explained since both consists of charged particles, thus there is an attraction between the positive ions and the ______________ end of the water molecule and the negative ions are attracted to the _____________ end of H
2 kJ/mol)
O (the H's) ; referred to as ion-dipole attraction (40-600
4) Covalent Network Crystals (pg 270)
-
-
1)
2)
3)
4)
(network solids) e.g. quartz, diamond, graphite a vast # of atoms are rigidly held in space and held by a network of covalent bonds in all directions
Properties very _____________ but ________________ ; exception is ______________________
_______________ conductors of electricity (exception ____________________ )
____________________ in most solvents very ____________________ melting points (e.g. M.P. of diamond is _______________ )
Allotropic Forms of Carbon
allotropy is defined as the existence of an element in two or more forms in the same physical state (e.g. diamond and graphite
watch videos www.howstuffworks.com
: Pure Carbon: The Chemistry of Diamonds and Graphite and Pure
Carbon: The Science of Nanotubes

Structure
Diamond
-____________dimensional; each C surrounded by ______ others
-
-
-
Graphite
____________ dimensional each C surrounded by ________others
_________ plates are arranged in
_____________; force between layers is ____________________
- _________ bonds exist
Hardness
M.P.
Density
Conduction of
Electricity
Solubility
Uses hardest substance ( pg 271 )
3700
3.5 g/cm 3 non conductor
_____________ in ordinary solvents gem stones, cutting, drilling and grinding soft and greasy
3600
2.26 g/cm 3 good conductor due to free mobile electrons
same dry lubricant for machines working at extremely high temp., lead in pencils, electrodes, tennis rackets, fishing rods
59

Research the structure of buckyballs (pg 272) and potential uses.
60
-
-
- c.
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
Silicon Dioxide(pg 271) main constituent of quartz
3 dimensional network solid; M.P.=2600 o C each Si is surrounded by 4 oxygen on the geometry of a tetrahedron
3) Metallic solids ( pg 269 )
- consider the outer electron configuration of the following metals
-
-
-
Na - ----- ----- ----- -----
3s 3p
Mg - ----- ----- ----- -----
3s 3p
Al - ----- ----- ----- -----
3s 3p note in the above there are but a few valence electrons (1,2 or 3) and there are many
_________________ orbitals which are very close in energy to those containing the _______________ electrons metallic atoms are very close to each other, the empty orbitals ____________ to form a continuum of available space through which electrons can travel since the electrons are no longer restricted to just one atom they can spread throughout they are known as ____________________ electrons

-
-
61 thus metals are composed of ___________________ charged metallic ions surrounded by a _________ of valence electrons metallic bond is due to ___________________ attraction between the positively charged ions and the sea of mobile electrons (75-1000 kJ/mol) conditions necessary for metallic bonding: -
1) ________ valence orbitals 2) ______________ ionization energies
B) Properties of Metals(pg 269) i) ii) iii) iv) lustre or reflectivity- valence electrons _________and ___________ the energy of all wavelengths of light electric conductivity - when an electrical force is applied, the delocalized electrons flow readily through the metal heat conductivity - heat applied to one section of a metal increases the motion of electrons at this point.
This motion is transmitted to nearby electrons and the motion (kinetic energy) of the rapidly travelling electrons results in the heating of other parts of the crystal workability - metals are malleable i.e. they can be hammered into thin sheets and ductile i.e. can be drawn into thin wires ; metallic bonds are _________________ thus one plane of atoms may slide over another when a stress is applied
1. Metallic Crystals (Metals)
Examples: Na, Cu, Fe, Mn
2. Ionic Crystals (Ionic Solids)
Examples: NaCl, MgCl
2
, MgO
3. Covalent Crystals (Network
Solids)
Examples (small class!): C(diamond), SiC(s),
SiO
2
(quartz)
: Valence electrons form mobile _____ of electrons which comprise the metallic bond.
: Attraction of charged ions for one another.
Lattice energy is a measure of ionic bond strength.
Network solids are extremely hard compounds with very high melting and boiling points due to their endless 3-dimensional network of covalent bonds.
4. Molecular Crystals
Examples:
(a) Need H bonded to O, N or F: H
2
O, HF,
NH
3
.
(b) C
6
H
6 (benzene), polyethylene, I 2
, F and all the compounds from (a) above.
2
,
One or more of the following:
(a)
: Hydrogen bonds are weaker than covalent bonds, but stronger than (b) or (c) below.
(b)
(induced dipole – induced dipole or
London dispersion forces): universal force of attraction between instantaneous dipoles. These forces are weak for small, lowmolar mass molecules, but large for heavy, long, and/or highly
(c) CHF
3
, CH
3
COCH
3 (acetone)
molecules. They usually dominate over (c) below.
(c)
: these forces act between
molecules. They are much ________than hydrogen bonding.
Note:
is a category which includes
categories (b) and (c) above.
5. Atomic Crystals
Examples: He, Ne, Ar, Kr, Xe
: See Section 4(b) above.
*Note: Many of the compounds given as examples are
solids at room temperature. But if you cool them down to a low enough temperature, eventually they will become solids.

62
depend on these forces. The
the forces between the particles,
(a) the ______________ the
.
(b) the ___________________ the
.
(c) the _____________ the
(partial pressure of vapor in equilibrium with liquid or solid in a closed container at a fixed temperature).
(d) the ______________ the
(resistance to flow).
(e) the _______________ the
(resistance to an increase in surface area).
Worksheet
1. What
will each of the following substances form in its solid state? Choices to consider are
,
, network
, atomic, or
crystals( polar or non-polar)
(a) C
2
H
6
__________ (b) Na
2
O ____________ (c) SiO
2
______________
(d) CO
2
______________
(g) Al ________________
(e) N
2
O
5
__________
(h) C(diamond) ______
(f) NaNO
3 ______________
(I) SO
2
________________(j) Kr ____________
2. Circle
the compounds in the following list
in the liquid state: If not predict major force.
(a) CH
3
OCH
3
(b) CH
4
(c) HF
(dimethyl ether)
(d) CH
3
CO
2
H
(acetic acid)
(e) Br
2
(f) CH
3
OH
(methanol)
3. Specify the
involved for each substance in the space immediately following the substance.
Substance
#1
Predominant
Intermolecular
Force
Substance #2
Predominant
Intermolecular
Force
Substance with
Higher Boiling Point
(a) HCl(g) I
2
(b) CH
3
F
(c) H
2
O
(d) SiO
2
(e) Fe
(f) CH
3
OH
(g) NH
3
(h) HCl(g)
CH
3
OH
H
2
S
SO
2
Kr
CuO
CH
4
NaCl
(i) SiC Cu
4.
Cu
Identify the following types of crystals based on properties : a. b. c. d. e. f. g. extremely hard , very high melting point and does not conduct electricity ____________________ melting points vary; very good conductors
L.D. forces only but does not form molecules
Hard , brittle, conduct in molten state poor conductors, low m.p. L.D. forces
Dipole – dipole forces
H- bonds exist
Use the following list and place in the blanks above as examples of the different crystals
SiC HCl NaCl Kr H
2
O SO
2
CuO Fe CH
4
H
2
____________________
____________________
____________________
____________________
____________________
____________________
S SiO
2
CH
3
F CH
3
OH

63
Properties and Forces in Different Types of Solids
Fill in the various properties on the chart below
Properties of
Solids
Ionic
Solids
Molecular
Nonpolar
Solids
Molecular
Polar
Solids
Network
Solid
Metallic
Solid
Possible Responses
Boiling Points/
Melting Points
(s)
Electrical
(l)
Conductivity
(aq)
If Yes, Reason for Electrical
Conductivity
Thermal
Conductivity
Brittle
________
________
_________
_________
_________
_________
________
________
________
________
N/A
Extremely High, High
Moderate, Low, Varies
Yes or No
Free flowing electrons,
Free flowing ions or
Delocalized electrons
Yes, conductor or
No, insulator
Yes or No
Malleable
Hardness 3-D _ _ _
_ _
2-D
Yes or No
Extremely Hard, Hard
Soft, or Varies
Unit at
Lattice Sites
Possible forces of attraction between units ranking them from weakest to strongest if more than one.
1) cations in a sea of electrons
2) cation/anion attract
3) polar molecules
4) nonpolar molecules
5) atoms covalently
bonded
1)
2)
3)
4)
5)
6) metallic bond
Ldf
H-bonding
Dipole-dipole
Ion-ion attraction
One Gigantic Molecule
thus no IMFA
Examples dip-dip ex.
--------------
H-bonding ex.
3-D ex.
-------------
2-D ex.
List at least three examples for each type of solid
Research BN ( Boron nitride ). How would you classify it ? What are some of it’s applications ?
____________________________________________________________________________

____________________________________________________________________________
64

IMF's, Liquids, & Solids
17.
18.
19.
20.
21.
22.
11.
12.
13.
14.
15.
16.
23.
24.
25.
26.
27.
4.
5.
6.
7.
1.
2.
3.
8.
9.
10.
Indicate the strongest IMF holding together crystals of the following:
Atomic Molecular Crystal
London forces
Dipoledipole attractions
Hydrogen
Bonds
NH
3
Kr
HCl
F
2
KMnO
4
NaCl
SO
2
CO
2
C
3
H
8
CH
4
CH
3
Cl
HF
C
6
H
6
NO
H
2
SO
4
WC
Si
SiO
2
C
(graphite)
N
2
CH
3
OH
Ag
(C
2
H
5
)
2
NH
NaOH
Al
PCl
3
XeF
4
London forces
Metal
Metallic
Bonds
Network
Solid
Covalent
Bonds
Ionic
Crystal
Ionic
Bonds
65

Atomic
London forces
28.
He
29.
Na
30.
CO
31.
Ar
32.
Ba(OH)
2
33.
O
2
34.
H
2
O
35.
NH
4
Cl
36.
Hg
37.
P
4
38.
HCN
39.
CaO
40.
N
2
H
2
41.
H
2
42.
Pb
43.
XeF
2
44.
SF
4
45.
SiC
46.
Si
4
H
10
47.
PH
3
48.
SiH
4
49.
H
2
Se
50.
C
2
H
2
51.
I
2
52.
Cu
53.
AsH
3
54.
K
2
S
London forces
Molecular Crystal
Dipole-dipole attractions
Hydrogen Bonds
Metal
Metallic Bonds
66
Network
Solid
Covalent
Bonds
Ionic Crystal
Ionic
Bonds

Independent Research
67
Evaluate the benefits to society, and the impact on the environment, of specialized materials that have been created on the basis of scientific research into the structure of matter and chemical bonding (e.g., bulletproof fabric, nanotechnologies, superconductors, instant adhesives)
Nanoparticles have many potential applications in medicine, including the improvement of drug delivery systems, the enhancement of diagnostic images, and use in surgical robotics, all of which could improve the effectiveness of our health care system.
However, nanoparticle contamination can have a negative effect on the environment.Research the following evaluating the benefits and the impact on the environment:
1.
Kevlar ( bulletproof fabric )
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
2.
Nanotechnolgies
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
3.
Instant adhesives
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
4.
What precautions are taken to protect the health and safety of people working with nanoparticles?
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
5.
What properties of disposable diapers enable them to hold so much liquid?
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
6.
What impact has the widespread use of disposable diapers had on the environment?
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
7.
8.
______________________________________________________________________________________
What impact has the development of synthetic fibres such as nylon had on society?
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
In what ways has the invention of the silicon chip changed society?
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________

68
Watch video www.teachersdomain.org
:
A Nanotube Space Elevator
1.
What is nanotechnology?
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
2.
Describe several different arrangements of carbon molecules. What properties of carbon atoms explain their use in nanotechnology?
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
3.
In what other ways do you see nanotubes being used?
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
4.
Science fiction stories of the past have predicted technologies that we have today. Can you identify any stories where this has happened? Do you think Arthur C. Clarke's idea will come true?
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________
______________________________________________________________________________________

Bonding and the State of Matter of Pure Substances : 1. Complete the following blanks by writing in the type of bond and likely state(s) of matter at SATP, for each example.
2. For each of the following substances, list all the types of chemical bonds believed to be present according to our current theories and predict the most likely state of matter at SATP.
(a) Ge
(b) Co
(c) CO
(d) MnO
2
COOH (e) CH
3
(f) SiC
(g) C
2
H
5
Cl
(h) CH
3
NH
2
(i) CuSO
4
5H
2
O
(j) Extension: UF
4 and UF
6
Bonded Bond particles type
Nature of bonding
State at
SATP
Structure of matter metal atoms (a) variable strengths, nondirectional bonds (b) continuous, close-packed crystals metalloid atoms (c) very strong directional bonds (d) continuous, network crystals ions strong nondirectional bonds continuous, regular crystals
(e) nonmetal atoms (g) molecules (h)
(i) (j) very strong directional bonds
Relatively weak bonds: nondirectional bonds; somewhat directional bonds; directional bonds
(f)
*******
(k) formation of single molecules continuous, often irregular arrangements
Day 14 -
Day 15 -
Lab Properties of substances pg 202
Review pg 273 ( 2,4,7,9 ) pg 276 ( 2,3,4,6,8,9 ) pg 281 (1-20 omit 9) pg 282(2,4,12,14-26), 28 pg 286(1-40) pg 288 (1,4,7,8,9a,c,e, 13,14,15,16,21,22,23,(25-30)
Day 16 Test
69

70
L a b : : C o m
Purpose: p a r r i n g D i f f f f e r e n t T y p e s o f f S o l l i d s N a m e s _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _
To determine whether substances are ionic, network covalent, or molecular from some of their physical properties
Materials: Aluminum foil, distilled water, tongs, spot plate, hot plate, sugar, table salt, sand, naphthalene, sulfur
Procedure:
Testing for relative melting point
1.
2.
3.
Use a small strip of aluminum foil, fold the foil to make separate areas for the five solids to be tested.
Place a small crystal of each solid on the foil.
Carefully place the foil on the hot plate. Turn on the hot plate – no higher than 5. Observe the order of melting. Record the order of melting.
4.
1.
2.
Remove the foil with the tongs and dispose of the foil and contents in the wastebasket. Turn off the hot plate.
Testing for solubility
Use a spot plate. Add distilled water to 5 wells. Add hexane to 5 other wells
Add a small crystal of each solid to each well and stir with a toothpick. Record your observations. Note if the solubility is high, low, or none. Repeat the above using hexane as the solvent.
Testing for conductivity
1. Test the conductivity of each of the solutions prepared above. Record observations. Use a + if the solution conducts and – if it does not.
When finished, empty the spot plate into the waste jar and clean the spot plate. 2.
Testing for hardness
1.
Rub one solid onto the surface of another. Rank the relative hardness of each solid on a scale of 1 to 5. A solid that receives no scratch marks when rubbed by the other four is the hardest (5). A solid that can be scratched by the other four is the softest (1).
Data: (Complete) ( 21 marks ) ( Communication )
Substance Formula Melting
Point
Order of melting
Relative hardness
Solubility in water
Solubility in hexane
Conductivity
Solid
Liquid
Sugar C
6
H
12
O
6
186  C
Salt
Sand
NaCl
SiO
2
Naphthalene C
10
H
8
Sulfur
Unknown 1
S
Unknown 2
801  C
1710
 C
80  C
94  C

71
Conclusions : Knowledge ( 20 marks )
Substance Type of Solid Major Force Reason for classification
#1
Reason for classification
#2
Sugar
Salt
Sand
Naphthalene
Sulfur
Unknown 1
Unknown 2
5.
4.
2.
Questions: (Inquiry)
1. Compare your order of melting to the melting points given in the table. (1)
_______________________________________________________________________
3.
________________________________________________________________________
________________________________________________________________________
Which physical property is best to identify ionic substances? Explain. (1)
_______________________________________________________________________
________________________________________________________________________
Once you have identified a substance as either ionic or covalent network, which physical property is best for distinguishing between these two? Explain.(1)
_______________________________________________________________________
________________________________________________________________________
Which physical property is best used to identify the molecular substances ( those made up of molecules)?
Explain.(1)
_______________________________________________________________________
________________________________________________________________________
There is a fourth type of solid, metallic. What physical property would you use to identify a substance as a metallic solid? Explain.(1)
_______________________________________________________________________
________________________________________________________________________

Lab Discussion : Inquiry
1. For each type of compound, list the types of intermolecular forces that exist.(3) a) polar molecule _______________________ b) nonpolar molecule _______________________ c) molecule with OH, NH, HF bonds _____________________________
2. i) For each compound below, write down the types of intermolecular forces that exist between its molecules. ii) Circle the substance that will have the a) LiCl or HCl
___________________ _______________________ b)
____________________ c)
Xe
CH
3
CH
2 or I
2 or Cu
_________________________ __________________
OH or CH
3
CH
2
CH
3
___________________ _________________________ d) NO or N
2
___________________ __________________________ e) H
2
Se or H
2
Te
___________________ ___________________________ f) OF
2 or CO
2
___________________ ____________________________ g) NH
3 or PH
3
___________________ _____________________________ h) C
2
H
6 or C
4
H
10
___________________ _____________________________ i) CH
3
OH or CH
3
F
___________________ ______________________________
boiling point. (13) k) l)
NH
3
____________________
CH
3
CH
2
____________________ or PH
3
______________________________
OH or CH
3
CH
2
CH
2
OH
______________________________ m) CH
3
COOH
____________________ or (CH
3
)
2
C=O
______________________________ n) SiC or
_____________________
K /20
KCl
______________________________
I /21 C / 21
72

73

74