SCH 4U1 Unit 1 Structure and Properties Day 1 1. Development of the Atomic Theory : ( page 160 - 174 ) 300 B.C. - Democritus: matter made up of tiny indivisible particles, “_____________” 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 ________________________________________. 1
Transcript
SCH 4U1 Unit 1 Structure and PropertiesDay 11. Development of the Atomic Theory : ( page 160 - 174 )
300 B.C. - Democritus: matter made up of tiny indivisible particles, “_____________”
1809 - John Dalton - Billiard ball model - atom was _________, ___________, _____________ sphereLimitations: 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 )
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:
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1 Stove Element-Red Hot 2 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 = mc2.considered light to be __________________________ and explained the photoelectric effect by stating that _____________________________________________________.
Electromagnetic spectrum :
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1. Radiation with shortest wavelength is ______________ and the one with longest is __________2. Microwaves have a _____________ wavelength than x - rays3. Infrared has a _____________ frequency than visible light4. Visible light represents a ______________ spectrum with the colours ROYGBIV arranged in
order of ____________ wavelength and _______________ energy.
- 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 : 3rd to 2nd ________, 4th to 2nd ________, 5th to 2nd ______6th to 2nd _____________
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- when electrons absorb certain wavelengths of light the atoms emission spectrum is called an _________________ spectrum.
- Jumps to the 1st level result in _________________ light and jumps to the 3rd level _________ light
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 determined. Actually, the e- is so small that it has both particle and wave properties so we cannot specify exact orbits. Another problem is when an electron changes energy levels during the emission of atomic spectra.
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.
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
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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. 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 ______2. 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
- apartments will represent the subshells- on the first floor there is a 1s apartment ; on the 2nd floor there is 2s and a 2p apartment- on the the 3rd floor the apartments are _________, ___________, and ______________- on the 4th floor the apartments are _________, _______, __________, and ______________.
- Complete page 182 ( 3,4,5 )
Day 3 3. Magnetic Quantum # (ml ) (pg 182)
- identifies the ________________ of the orbital in space- ml ranges from _________ to ____________- if l = 0 _________ subshell ml = ___________ i.e. _________ orientation
If l=1 __________ subshell ml = __________ i.e. _________ orientations identified as px, py and pz i.e ml =-1 (px) ml =0(py) ml = +1 (pz )
- l=2 identifies the ______ subshell. How many orientations? _______
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If l=3. How many orientations? _______ since m= __________________________________
- general formula for the # of orbitals for each main shell is _________________
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 _______________
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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 2nd one __________
- 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- 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
- individuals represent the electrons ; first person in has a positive attitude 2nd 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 Name Values Meaning
n ___________ 1,2,3 ... n fixes the size of orbital value and energy
l ___________ 0,1,2 ... n-1 fixes the __________ of the orbital l=0 spherical ____ orbital 1 two lobed ____ orbital 2 four lobed ____ orbital 3 more complex ____ orbital
m __________ -l...0...+l defines ___________of orbitals in a magnetic field
s __________ +1/2,-1/2 defines direction of spin or rotation of electron on it axis
Distinguish clearly between an orbital and a orbit ( pg 185 )
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2. Complete the following:1. The first one has been done for you.
Shell
K
L
M
N
PrincipalQuantum(n)
1
___
_____
_____
SecondaryQuantum #(l)
0
___ ____
_____
_____
_____
_____
_____
_____
_____
Subshells
s
____
____
_____
_____
_____
_____
_____
_____
_____
# ofOrbitals
1
_____
_____
_____
_____
_____
_____
_____
_____
_____
Day 4 Quantum Rules
1) The maximum # of electrons in any main shell is _________ ; n=1 ___ , n=2 ___ , n=3 ___ n=4 ____ .
2) The maximum # of subshells in any main shell is________________
3) The # of orbitals for each main shell is _____________e.g. n=1, l=0 --- 1 orbital (m=0) n=2, l =_______ m =_________
l = _______ m =_________
4) Pauli Exclusion Principal ( pg 188 )
- 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=____
-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. __________________________________
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- try N ________________________________________________6) Aufbau Rule
electrons enter subshells of __________________ energy first e.g. 3 _____ is filled before 4 _______
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)
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8. 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.
9. 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.
Assignment Quantum rules Name _______________
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1) Another name for the region in space where the electron may be is a(n):__________
2) 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 _______________
6) The orbital with the highest energy would be the (7s, 6p, 5f, 4p , 4d) ____________ .
7) Wolfgang Pauli noted that an orbital can have 0, 1, 2 but never _________ electrons.
8) A 1s subshell has _________ orbitals and can accommodate a maximum of _________ electrons.
9) A 4f subshell has _________ orbitals and can accommodate _____________ electrons.
10) What subshell is identified by the Quantum #'s n = 3, l = 2 ________________
11) 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) How many electrons in an atom could be identified by the quantum #'s: n = 3, l = 2 ________________
13) 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. n=3, l=3,m=-2 ___________________________b. 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
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Day 5/6 - Electron configurations
1H - 1s1 - 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 )
2He - 1s2
3Li -
4Be -
5B -
6C -
7N -
8O -
9F -
10Ne –
11Na –
15P -
20Ca –
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.
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4) Consider the electron configuration for 26Fe - 1s22s22p63s23p64s23d6 rewritten as: 1s22s22p63s23p63d64s2.
5) The structure for 24Cr is: ____________instead of 3d44s2 because a half-filled sub - shell is especially stable.
6) The structure for 29Cu is ____________ instead of 3d94s2.Give the electron configuration for: (Note: always rewrite at the end so that all principal quantum numbers are together)
1) 35Br -
-
2) 42Mo -
-
3) 82Pb -
-
Next let's try some positive or negative ions: eg. Fe3+
Steps:
1) First give the electron configuration for 26Fe:
1s22s22p63s23p64s23d6
2) Rearrange the order so that subshells with the same principal quantum #'s are together: 1s22s22p63s23p63d64s2
3) Now remove 3 electrons starting from the outermost energy level:
Fe3+ 1s22s22p63s23p63d5
orbital representations :
Type of Transition
4d
5d
4f
5f
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
39Y - 48Cd
57La, 72Hf - 80Hg
58Ce - 71Lu
90Th
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Ex. 2 : try O-2
O - 1s22s22p4 ----------> O-2 - __________________________________________
Try the following:
1) Cr3+
2) Ni2+
3) P-3
4) Se2-
7. Where possible we try to use shorthand notation see page 193
K [ ] ____
Br [ ] _________________
Zn [ ] ___________________
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 Do an orbital representation
Step 3 Circle the 4th 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.
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AssignmentPart I:Determine the element whose outermost electron is being defined by the following quantum numbers.
Element Quantum Number Code1. _____________ n = 1, l = 0, ml = 0, ms = - ½ 2. _____________ n = 2, l = 1, ml = -1, ms = + ½ 3. _____________ n = 2, l = 1, ml = 0, ms = - ½ 4. _____________ n = 2, l = 1, ml = 1, ms = + ½ 5. _____________ n = 3, l = 1, ml = -1, ms = - ½ 6. _____________ n = 3, l = 1, ml = 0, ms = + ½ 7. _____________ n = 3, l = 1, ml = 1, ms = - ½ 8. _____________ n = 4, l = 0, ml = 0, ms = + ½ 9. _____________ n = 3, l = 2, ml = -2, ms = + ½ 10. _____________ n = 3, l = 2, ml = -2, ms = - ½ 11. _____________ n = 3, l = 2, ml = 0, ms = + ½ 12. _____________ n = 3, l = 2, ml = 1, ms = - ½ 13. _____________ n = 3, l = 2, ml = 2, ms = + ½ 14. _____________ n = 4, l = 1, ml = -1, ms = + ½ 15. _____________ n = 4, l = 1, ml = 1, ms = - ½ 16. _____________ n = 4, l = 1, ml = 0, ms = + 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.
√ if not allowed Explanation of error if applicable1. n = 2, l = 1, ml = -1 _____ ________________________________________________________2. n = 1, l = 1, ml = 0 _____ ________________________________________________________3. n = 8, l = 7, ml = 6 _____ ________________________________________________________4. n = 1, l = 0, ml = 2 _____ ________________________________________________________5. n = 3, l = 2, ml = 2 _____ ________________________________________________________6. n = 4, l = 3, ml = 4 _____ ________________________________________________________7. n = 2, l = -1, ml = 1 _____ ________________________________________________________
Part III: What is the maximum number of electrons in an atom that can have these quantum numbers?
1. n = 4 ______2. n = 5, ml = +1 ______3. n = 5, ms = + ½ ______4. n = 3, l = 2 ______5. n = 2, l = 1 ______6. n = 0, l = 0, ml = 0 ______7. n = 2, l = 1, ml = -1, ms = - ½ ______8. n = 3 ______9. n = 2, l = 2 ______10. n = 1, l = 0, ml = 0 ______
WS Orbital Notation Name__________________
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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 4px 4py 4pz
3d 3d 3d 3d 3d
4s
3px 3py 3pz
3s
2px 2py 2pz
Electron configuration _______________________________________ 2s _______________________________________
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.
Zirconium
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Electron configuration____________________________
Reading Assignment ; Make notes on Laser technology page 203, spectrometers pg 205, x-rays pg 206, Cat Scan and MRI ( 206)
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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 ?
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 )
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- Specific subshells represent specific groups in the periodic table- group 1 s1 , group 2 _____________; groups 13 to 18 _________to___________; transition
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)
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 ___ e) 5d ___
b) 4p ___ f) 5f ___
c) 4px ___ g) n = 5 ___
d) 6d ___ 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 ml = -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. 1s22s22p7 b. 1s22s22p63s3 c. 1s22s22p63s23p64s23d12 d. 1s22s22p63s23p6
6. Explain why the following ground-state electron configurations are not possible:a. 1s22s32p3 b. 1s22s22p33s6 c. 1s22s22p73s23p8 d. 1s22s22p63s23p14s23d14
7. Give two examples of:a. an atom with a half-filled subshell
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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 n=3, l=2 and ml = 0 is _________.b. The number of valence electrons in the outermost p subshell of a sulfur atom is
_________.c. The number of unpaired electrons in a Mn2+
ion is _________.d. The subshell with the quantum numbers n=4, l=2 is _________.e. The ml values for a d orbital are ________________________.f. The allowed values of l for the shell with n=2 are _________.g. The allowed values of l for the shell with n=4 are _________.h. The number of unpaired electrons in the cobalt atom is _________.i. The number of orbitals in a shell with n=3 is _________.j. The number of orbitals with n=3 and l=1 is _________.k. The maximum number of electrons with quantum numbers with n=3 and l=2 is
_________. l. When n=2, l can be _________.m. When n=2, the possible values for ml are _________.n. The number of electrons with n=4, l=1 is _________. o. The quantum number that characterizes the angular shape of an atomic orbital is
_________.p. The subshell with n=3 and l=1is designated as the __________ subshell. q. The lowest value of n for which a d subshell can occur is n=_________.
9. Which sets of quantum numbers are unacceptable?a. n=3, l= -2, ml=0, ms= +½ b. n=2, l= 2, ml= -1, ms= -½ c. n=6, l= 2, ml= -2, ms= +½
10. Identify the group of elements on the periodic table which have the following ground state electron configuration:
a. ns2np3 ____ b. ns1 _____ c. ns2np6 _______ d. ns2 ______
Chemical Bonding (224)
1. Distinguish between ionic and covalent bonding
Ionic –
Covalent –
2. Lewis structures - for atoms the # of dots is equal to the last digit in the group number ; complete below for period 3 elements
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Steps To The Writing Of Lewis Structures (pg 225)
From the molecular formula given ,determine how many valence electrons are available from all atoms present in the molecular formula. Calculate the total.
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
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.
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
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
If steps 6 or 7 cannot be accomplished assume that this is a case that is an exception to the octet rule.
9. Draw the Lewis structure showing the covalent bonds with a short line
10. Watch out for oxy-acids the H and O are attached together
Ex.1 SF2
i. S - 6 valence e=sF - 2 x( 7 val. e=s ) Total val. e=s = 20e’s
ii. F- S- F
ii. Complete
Ex.2 CO32-
i. C is bonded to 3 O’sii. C - ___ val. e
3 O - 3 x _______ val. e’s O C O___ outside e’s O val. e’s
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Resonance Structures* For some molecules, there are multiple ways of placing the electrons between the atoms.* Structures that differ only in the arrangement of the electrons are called RESONANCE STRUCTURES. * Resonance structures are indicated using a double headed arrow.
Complete the following:
1) SO42- 2) HClO2
3) SiF4 4) CNO-
5) SO2Cl2 6) SO3
7) HCN 8) NH4+
9) N2O 10) ClF3
11) SO2 12) XeF2
13) ICl2- 14) N2F2
15) XeO4 16) SF4
17) BrF4+ 18) COCl2
Complete page 229 ( 10,12 ) page 230 (4)
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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.
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Determining the Shape of a molecule (pg 243)The best way to determine the architecture of a molecule is to:
1. Determine what the central atom is. 2. Draw the Lewis structure of the molecule. 3. Determine the number of bonding pairs and lone pairs around the central atom. 4. 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 Shape 2 0 AB2 Linear 2 1 AB2E Bent
2 2 AB2E2 Bent 2 3 AB2E3 Linear 3 0 AB3 trigonal
planar 3 1 AB3E pyramidal
3 2 AB3E2 T-shaped 4 0 AB4 Tetrahedral 4 1 AB4E
unsymmetrical tetrahedral ( see - saw )
4 2 AB4E2 squareplanar
5 0 AB5 Trigonal bipyramidal
5 1 AB5E Square pyramidal
6 0 AB6 Octahedral
Refer back to pg 26 and predict the shapes of the various molecules and ions.
3. Molecular Architecture Worksheet - complete the chart below for the following :
1. CCl4 2. ClO2- 3. XeF2 4. XeOF4 5. HCN 6.
ClO3-
7. XeO4
8. OF2 9. CO32- 10. NH3 11. NH4
+ 12. BrF3 13.O3 14. PCl6-
15. H2Se 16. HClO2 17. CH3NH2
Molecule or ion
Lewis Structure E.P.A. Molecular shape
Sketch
34
1. 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) __________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________
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, H2
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 shared electrons in overlapping 1’s’ orbitals
H
Water, H2O
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
1s
1s
37
The orbital diagram for water is:
H
1s
O
2s 2p
H
1s
C. ORBITAL HYBRIDIZATION
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 sp3
The carbon atom is then able to bond with four hydrogen atoms to form the methane molecule, CH4
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, N2, can also be used to show multiple bonds:
N2s 2p
N
2s
2p
bond
bonds
39
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)
3. 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.
For a molecule to be a dipole, two conditions must be met:a) ionic character in at least one of the bonds within the moleculeb) 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
1) all diatomic molecules consisting of unlike atoms have a dipole moment2) triatomic molecules can either be linear (no dipole moment) or bent (dipole moment)3) binary tetratomic molecules can be:
i) planar (no dipole moment) e.g. BF3
ii) pyramidal (dipole moment) e.g. NH3
Dipole Moments of Various Molecules; complete chart
Molecules Electronegativity Difference
Dipole Moment
HCl 1.08
CO 1.91
BeF2 0
CO2 0
OF2 0.30
H2O 1.85
BF3 0
40
Dipoles have a definite direction that can be represented by a vector, showing both magnitude and direction.(pg 254)
1) BeF2 F←Be→F ; addition of vectors 0→ ← non polar (dipoles cancel out)
2) H2O
3) BF3
4) NH3
- 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 Polar or non-polar
linear
bent
trigonal planar
pyramidal
t - shape
see-saw
tetrahedral
trigonal bipyramidal
square planar
Square based pyramidal
Octahedral
Classify the following as polar or nonpolar
a) BeI2 ________ b) CF4 _______c) CH2O ________
d) CO2 _______e) BF2Cl ______ f) CH3Cl ________
g) SF6 _______ h) SCl2 ______
41
Draw 2 structures for each of the following one should be polar the other non-polar :
a. PCl3F2
b. SF4Cl2
Complete page 256 ( 8,10,11 ) ( 1,2,4 )
42
43
Predicting Molecular Geometry 1.
Molecule or Ion (1) OF2 (2) H2CO (3) NO2+ (4) BF3 (5) SbF5
(a) No. of valence e’s
(b) Lewis structure
(c) Approximate bond angle(s)
(d) Electron group arrangement(e) Polar or non-polar molecule?
Ion: Not applicable
(f) Geometry name(g) Hybridization
______________________________________________________________________________2. For each of the molecules below fill in the indicated items in the chart. The central atoms are underlined.
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.
1. For each of the species listed above, do the following and record your answers on the worksheets provided.
2. Determine the total number of valence electrons in the molecule.
3. Determine the Lewis Structure of the molecule.
4. Determine the number of bonding electron pairs and non-bonding electron pairs around the central atoms(s).
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.
1. Classification of Solids: Based on Intra and Intermolecular Forces
Classification 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 stronginter ----> between molecules e.g. london dispersions , dipole-dipole forces and hydrogen bonding
--> weak forces
A. Crystalline Solids – organized pattern of particle arrangement.a. Atomicb. Molecularc. Networkd. Ionice. Metallic
B. Amorphous Solids – indistinct shapes because their particle arrangements lack order.a. Glassb. Polymers
1) Atomic Solids
intra ----> between atoms in a molecule e.g. covalent and polar covalent --> very stronginter ----> between molecules e.g. london dispersions , dipole-dipole forces and hydrogen bonding --> weak forces
- consider Helium (F.P.= -272.2oC) and Argon (F.P.= -189.2oC)- 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
A) Non-Polar Molecules
- consider the table below:
Molecule Total # of Electrons State at Room Temp
F2 18 gas; B.P.= -188oC
Cl2 34 gas; B.P.= -34.6oC
Br2 70 liquid; B.P.= 58.8oC
I2 106 solid; B.P.= 184.4oC
- 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) molecular shape
Account for the increase in B.P. from Fluorine to Iodine ________________________________________________________________________________
- consider the table below:
Molecule # of Electrons B.P.
CH4 10 - 162
C2H6 18 - 87
C3H8 26 - 45
C4H10 34 - 0.5
Questions:
1) Account for the increase in boiling point _____________________________________________________
2) Compare Cl2 and C4H10. Account for the difference in B.P.
______________________________________
3) Two isomers of butane have the boiling points 0oC and - 12oC.
a) C - C - C - C b) C - C - C \ C Which B.P. would you assign to each? Account. ___________________________________
- forces between molecules is dipole-dipole or hydrogen bonding
1) Dipole - Dipole
- 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 electroneg.difference
# of Electrons B.P.
ClF _______ 26 - 101
BrF _______ 44 - 20
CH3F ___and ___ 18 - 78
CH3Cl ___and ___ 26 - 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
Questions :
1. Account for the difference in B.P. between ClF and CH3Cl._______________________________________
2. Account for the difference in B.P. between C2H6 and CH3F
______________________________________________________________________________________ 3. Account for the difference in B.P. between ClF and C3H8
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 )
- 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
- consider the data:
Compound Electronegativity of Halogen
# of Electrons B.P.
HF 4.0 10 + 19.4
HCl 3.0 18 - 83.7
HBr 2.8 36 - 67
HI 2.5 54 - 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 Electronegativity of Group VI Element
# of Electrons B.P.
H2O 100
H2S - 61.6
H2Se -42
H2Te 4.0
- in H2O 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/cm3 at 0oC). When ice melts, this open structure collapses to give a more closely packed pattern
- considering the graph below why would it be predicted that H2O would have the lowest B.P. of the series. ______________________________________________________________________________________Explain waters abnormal behaviour._________________________________________________________ Why does H2O have a higher B.P. than HF? __________________________________________________ What are some other properties of water which can be accounted for by hydrogen bonding ?( page 262 )
Day 13 Characteristics of Molecular Substances (page 270)
1) Majority are at room temperature2) Solids are usually e.g. mothballs, wax3) conductors of electricity and heat since _____________________ 4) melting points5) Intramolecular forces are great i.e. requires a large amount of energy to decompose the
molecule6) Non-polar molecular solids are in water7) Polar molecular substances are in water and if acidic in nature form
_____________ solutions
Worksheet1. Draw the hydrogen bonds that exist in the following diagrams.
a) H b) H H CH3
: O: H H N
H ..O.. . . H3C
H H N: H
. . H N: H ..O
H3C 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) CH3F b) CH3NH2 c) CH3COOH d) CH3CH2CHO
6. 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
I) Structure - 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.
II) Properties
1) Very _________________ but brittle2) _______________ melting points3) ________________ conductors of electricity in the solid state but - __________ conductors in
the molten state4) majority are ___________________ in water5) _______________ 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
- 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 H2O (the H's) ; referred to as ion-dipole attraction (40-600 kJ/mol)
4) Covalent Network Crystals (pg 270)
- (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
Properties1) very _____________ but ________________ ; exception is ______________________2) _______________ conductors of electricity (exception ____________________ )3) ____________________ in most solvents4) 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
- main constituent of quartz- 3 dimensional network solid; M.P.=2600oC- each Si is surrounded by 4 oxygen on the geometry of a tetrahedron
3) Metallic solids ( pg 269 )
A) Structure-
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
- 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)
i) lustre or reflectivity- valence electrons _________and ___________ the energy of all wavelengths of light
ii) electric conductivity - when an electrical force is applied, the delocalized electrons flow readily through the metal
iii) 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
iv) 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
Intermolecular ForcesTypes of Solids* Intermolecular Force(s) Between Particles1. Metallic Crystals (Metals)Examples: Na, Cu, Fe, Mn
Metallic bonding: Valence electrons form mobile _____ of electrons which comprise the metallic bond.
extremely hard compounds with very high melting and boiling points due to their endless 3-dimensional network of covalent bonds.
4. Molecular CrystalsExamples: One or more of the following:
(a) Need H bonded to O, N or F: H2O, HF, NH3.
(a) Hydrogen bonding: Hydrogen bonds are weaker than covalent bonds, but stronger than (b) or (c) below.
(b) C6H6 (benzene), polyethylene, I2, F2, and all the compounds from (a) above.
(b) Dispersion forces (induced dipole – induced dipole or London dispersion forces): universal force of attraction between instantaneous dipoles. These forces are weak for small, low-molar mass molecules, but large for heavy, long, and/or highly polarizable molecules. They usually dominate over (c) below.
(c) CHF3, CH3COCH3 (acetone) (c) Dipole-dipole forces: these forces act between polar molecules. They are much ________than hydrogen bonding.
Note: Van der Waals Forces is a category which includes both categories (b) and (c) above.5. Atomic CrystalsExamples: He, Ne, Ar, Kr, Xe
Dispersion forces: See Section 4(b) above.
*Note: Many of the compounds given as examples are not solids at room temperature. But if you cool them down to a low enough temperature, eventually they will become solids.Physical properties depend on these forces. The stronger the forces between the particles,(a) the ______________ the melting point.(b) the ___________________ the boiling point.(c) the _____________ the vapor pressure (partial pressure of vapor in equilibrium with liquid or solid in a closed container at a fixed temperature).(d) the ______________ the viscosity (resistance to flow).(e) the _______________ the surface tension (resistance to an increase in surface area).
Worksheet1. What type of crystal will each of the following substances form in its solid state? Choices to consider are metallic, ionic, networkcovalent, atomic, or molecular crystals( polar or non-polar)(a) C2H6 __________ (b) Na2O ____________ (c) SiO2 ______________(d) CO2 ______________ (e) N2O5 __________ (f) NaNO3 ______________(g) Al ________________ (h) C(diamond) ______ (I) SO2 ________________(j) Kr ____________
2. Circle all the compounds in the following list which would be expected to form intermolecular hydrogen bonds in the liquid state: If not predict major force.(a) CH3OCH3 (b) CH4 (c) HF (d) CH3CO2H (e) Br2 (f) CH3OH (dimethyl ether) (acetic acid) (methanol)
3. Specify the predominant intermolecular force involved for each substance in the space immediately following the substance. Then in the last column, indicate which member of the pair you would expect to have the higher boiling point.
Substance #1
Predominant Intermolecular
ForceSubstance
#2
Predominant Intermolecular
Force
Substance with Higher Boiling
Point
(a) HCl(g) I2
(b) CH3F CH3OH
(c) H2O H2S
(d) SiO2 SO2
(e) Fe Kr
(f) CH3OH CuO
(g) NH3 CH4
(h) HCl(g) NaCl
(i) SiC Cu
4. Identify the following types of crystals based on properties :
a. extremely hard , very high melting point and does not conduct electricity____________________b. melting points vary; very good conductors ____________________c. L.D. forces only but does not form molecules____________________d. Hard , brittle, conduct in molten state____________________e. poor conductors, low m.p. L.D. forces____________________f. Dipole – dipole forces____________________ g. H- bonds exist
____________________Use the following list and place in the blanks above as examples of the different crystalsCu SiC HCl NaCl Kr H2O SO2 CuO Fe CH4 H2S SiO2 CH3F CH3OH
Properties and Forces in Different Types of Solids
Fill in the various properties on the chart below
Properties of Solids
Ionic Solids
MolecularNonpolar Solids
Molecular Polar Solids
Network Solid
MetallicSolid
Possible Responses
Boiling Points/ Melting Points
Extremely High, HighModerate, Low, Varies
(s)Electrical (l) Conductivity (aq)
________________
__________________
__________________
________________
________________ N/A
Yes or No
If Yes, Reason for Electrical Conductivity
Free flowing electrons,Free flowing ions or Delocalized electrons
Atomic Molecular Crystal Metal Ionic Crystal Network Solid
London forces
Londonforces
Dipole-dipole attractions
Hydrogen Bonds Metallic Bonds IonicBonds
Covalent Bonds
28. He
29. Na
30. CO
31. Ar
32. Ba(OH)2
33. O2
34. H2O
35. NH4Cl
36. Hg
37. P4
38. HCN
39. CaO
40. N2H2
41. H2
42. Pb
43. XeF2
44. SF4
45. SiC
46. Si4H10
47. PH3
48. SiH4
49. H2Se
50. C2H2
51. I2
52. Cu
53. AsH3
54. K2S
Independent Research
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) Sample issue: Nanoparticles have many potentialapplications 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:
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. What impact has the development of synthetic fibres such as nylon had on society?__________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________
8. In what ways has the invention of the silicon chip changed society?__________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________
Watch video www.teachersdomain.org : A Nanotube Space Elevator
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) MnO2
(e) CH3COOH (f) SiC (g) C2H5Cl (h) CH3NH2
(i) CuSO45H2O (j) Extension: UF4 and UF6
Bonded Bond Nature of State at Structure of particles type bonding SATP matter
metal variable strengths, continuous, atoms (a) nondirectional
bonds (b) close-packed
crystals metalloid
very strong continuous,
atoms (c) directional bonds (d) network crystals
ions
(e)
strong nondirectional bonds (f)
continuous, regular crystals
nonmetal
very strong formation of
atoms (g) directional bonds ******* single molecules
Lab: Comparing Different Types of SolidsLab: Comparing Different Types of Solids Names __________________________Names __________________________ Purpose: To determine whether substances are ionic, network covalent, or molecular from
some of their physical propertiesMaterials: Aluminum foil, distilled water, tongs, spot plate, hot plate, sugar, table salt, sand,
naphthalene, sulfurProcedure:Testing for relative melting point1. Use a small strip of aluminum foil, fold the foil to make separate areas for the five solids to
be tested.2. Place a small crystal of each solid on the foil.3. 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. Remove the foil with the tongs and dispose of the foil and contents in the wastebasket.
Turn off the hot plate.Testing for solubility1. Use a spot plate. Add distilled water to 5 wells. Add hexane to 5 other wells2. 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 conductivity1. Test the conductivity of each of the solutions prepared above. Record observations. Use a
+ if the solution conducts and – if it does not.2. When finished, empty the spot plate into the waste jar and clean the spot plate.
Testing for hardness1. 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 Formu
laMelting Point
Order of melting
Relative hardness
Solubility in water
Solubility in hexane
ConductivitySolid Liquid
Sugar C6H12O6
186 C
Salt NaCl 801 C
Sand SiO2 1710 C
Naphthalene
C10H8 80 C
Sulfur S 94 C
Unknown 1
Unknown 2
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
Questions: (Inquiry)1. Compare your order of melting to the melting points given in the table. (1)
3. Once you have identified a substance as either ionic or covalent network, which physical property is best for distinguishing between these two? Explain.(1)_______________________________________________________________________
5. There is a fourth type of solid, metallic. What physical property would you use to identify a substance as a metallic solid? Explain.(1)_______________________________________________________________________
