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Archive – Q3 week 1 – 19-20

Week of 1/27 – 1/ 31
Reminder….PES!!!!
1/27 -Monday – period 2/3
 
1. Absorption Spectrum – demo vs. Emmission spectrum
 
2.  Weekend Homework Review – Bohr and DeBroglie Form.
 
– Bohr limitations – His model was doomed…..(wave behavior of electrons!)
 
– Louis DeBroglia – Electrons as waves! – Matter has waves properties!!!
 
– De Broglie wavelength ,Compton effect, Davidson  – Germer, William Thompson
 
3. Shrodinger’s equation —> quantum numbers
 
                               – period 4
 
1. Bright Line homework from Thursday – quick review – shows limitation in Bohr model
2.  Weekend Homework Review – Bohr and DeBroglie Form.
 
– Bohr limitations – His model was doomed…..(wave behavior of electrons!)
 
– Louis DeBroglia – Electrons as waves! – Matter has waves properties!!!
 
– De Broglie wavelength ,Compton effect, Davidson  – Germer, William Thompson
 
3. Shrodinger’s equation —> quantum numbers
Review:
atomic structure 2 – Bohr Key p.pdf
View Download
 
atomic structure 3 – de Broglie Key.pdf
Bright Line spectrum Key p.pdf
View Download
The math from the Bohr team – Copenhagen interpretation gave some very hard phenomenon to think about that must be accepted:
 
1) Heisenberg Uncertainty principle
2) Superposition
3) Quantum Entanglement
4) Multiple Worlds Theory – explains the collapse of the superposition (wave function)
 
In the words of Richard Feynman, (Nobel Prize winner, 1965 – quantum electrodynamics)
 
“If you do not like it. Too bad! Go somewhere else! That is how the universe works!” 
                           
A new version of the Newtons’ wheel demo:
 
1/27 -Monday Homework – 
 
1. Quantum Timeline is due tomorrow in the purple crate. Watch the video till the end.
 
2.  Watch the quantum number lesson video and complete the form. This reviews the solutions to the Schrodinger equation that will allow us to determine the energy and relative positions of the electrons around the nucleus based on their quantum energy states.  All this means is that we will get electron configurations from the equation FROM the solutions to the Schrodinger equation.
 
Make sure you have completed your timeline before you begin this lecture.  This lecture about the values that we get from the Schrodinger equation that was “discovered over a weekend” that helped team Einstein (who was pitted against Team Bohr) have more of a visual model of the atom using standing waves of electrons in orbitals).  Remember that Team Bohr described the atom and its behavior entirely with mathematics and equations.  It is the solutions (quantum numbers) to the Schrodinger equation that have lead us to our current electron configuration of atoms that explains all of the chemical nature of atoms.
 
Remember electron configurations this summer?   Na:  1s22s22p63s1
This is the arrangement of electrons in an atom based on electrons existing in standing waves in a region called an orbitals. These orbitals are quantized and their arrangement is also based on its proximity to the nucleus, much like Bohrs first quantum model of the atoms with electrons moving in orbits.  Example:  3s electrons are farther away from the nucleus and have greater energy (unstable) than 1s electrons who are closer to the nucleus (more stable).  The difference to the Bohr model is that electrons are existing as standing waves in a particular geometry.  The basics of the Bohr model is still upheld (that is why he won a Nobel Prize for his first model).
You learned electron configurations this summer and did so without the rest of the atomic structure history and concepts.  We will re-vist electron configurations but will develop the concepts through quantum numbers!
 
The electrons are organized in a atom in 3-d wave function called orbitals! These shapes are defined by the Schrodinger equation and are quantized and thus they are restrictive to each other.
Electrons are not arranged in orbit but do occupy a space outside the nucleus based on proximity and complexity of the shape of the orbital.
 
The electrons are arranged in the atom in the following way:
 
Principle Energy levels = n # (proximity)
Sublevels =  (type of orbital with a distinct 3d shape) = # preferred axis
Orbitals = ml (different versions of the same type of sublevel)= how many orientations in space
Electrons in a single orbital = ms with a specific spin
 
The quantum numbers restrict each other by the following rules:
They are the 4 numbers that define the each electron in the atom. Every electron must have its own unique set of 4 quantum numbers.
N     =  (any integer 1 , 2, 3, to infinity)
      =  (0 to N-1)
ml   = (- l to + l)
ms  = (+1/2, -1/2)
Quantum number lecture:
Making a new one..Did not like the old one..
Quantum number Form: 

Quantum Number form 1920

 

1/28 – Tuesday – 

 
Electrons are arranged from lowest energy to higher energy = Aufbau Principle
 So they fill from lower n to higher n  and from lower energy orbitals (l) to higher energy orbitals. Electrons fill sublevels in order ( s,p,d,f).
 
Electron are arranged in n = principle energy levels
                                                    l  = sublevels
                                                 ms =  individual orbitals (actual 3 -d shape where pairs of electrons can exist)
 
Every Electron has a unique set of 4 quantum numbers (including spin msPauli Exclusion Principle
These quantum numbers describe how the electrons are arranged in the atom based on energy.
 
                                      The basic organization is principle energy level (n = 1, 2, 3, etc.) – “shell”
 
                                                                                     sublevel  (s,  p,  d,  f, ) – type of orbital and all of its orientations
                                                                                                   l = 0,  1,   2,  3
 
                                                                                      orbital (individual orbital of a single orientation of a sublevel)
                                                                                                        ml = -l , 0, +l
This arrangement uses notation that describes the organization of the electrons in principle energy levels and sublevels. The exponent is the number of electrons in the TOTAL sublevel that includes all the orbitals of different orientation of the same sublevel.
 
s = 1 orbital   (l = 0, ml = 0)
p = 2 orbitals (l = 1, ml = -1,  0, +1)
d = 5 orbitals (l = 2, ml = -2, -1, 0, +1, +2)
f  = 7 orbitals (l  = 3, ml = -3, -2, -1, 0, +1, +2, +3)
g =  ?
 
Electron configuration of Al:           Nucleus:  1s22s22p63s23p1
 
Orbital notation for Al:
Each box represents a single orbital.
 
Which electrons are the most stable? Unstable?
*Connections – Our quantum numbers refer to specific energy levels allowable by each quantized element.  The arrangement of electrons in quantized energy orbitals that were solutions to the Erwin Schrodinger equation (quantum numbers) are actually hidden in the periodic table that was already arranged according these energy levels (unbeknownst to Mendeleev and Moseley!!) The periodic table that we use is a condensed version that does not insert the  f block ( l = 3) because it would not fit on most pages!!!!
 

 

 
1. Quantum Number review
2. Electron configurations and orbital notations
 
a) Need to go over writing configurations on worksheet opposite, d orbitals and f orbital shape.
b) Complex electron configurations, shorthand method
c) Extreme overlapping of Uranium
d) valence electrons, core electrons
 
                            period 3/4 –
 
1. Quantum Number review
2. Electron configurations and orbital notations
Atomic Structure 4b -Electron configuration.pdf
View Download
 
Atomic Structure 4b -Electron configuration key.pdf
View Download
Please play this video which was an old flash animation that reviews the energy levels of electron in the orbital notation as electrons fill from lowest energy levels to higher:

Another very good animation that will review electron configuration :
Download file and open in Fire Fox.
 
ElectronConfiguration.html
Download

 

 electron configuration 1   electron configuration 2
1/23  – Tuesday Homework:
1.  Please complete Atomic Structure 4b -Electron configuration.pdf worksheet and review the key.
     If you need I have optional lectures posted above but do not worry I will complete in class.
     Period 2 students -Try this and  I will go over this tomorrow.      
      Period 3 students – We completed in class.
 
 
2: Please complete AP periodicity and Electron Configuration Form 1  based on the presentations below and the blue book. ( I have scanned the pages that you need).
 
IF YOU WANT TO Download the Blue Book for the homework:
 
Blue book Chapter 7 p.pdf
View Download
Powerpoint Presentation 1:
 
 
Blue BooK:
Tonight’s Form: Expand your answers here!  Do not get vague. This form is a place to show off!
                                    Due by 4:30 Wednesday morning.
End of  Tuesday!

 

1/29 – Wednesday – period 2/3
 
1. Review Quantum Number Form 1 worksheet:
 
Quantum Number workheet 1 key p.pdf
View Download
2. Review/hand back homework
*Connections – When we write electron configurations we are really writing the energy levels of each different electron defines by the 4 unique quantum numbers (3 of which are solutions to the Schrodinger equation).
 
Pauli Exclusion Principle – Electrons must have a set of 4 unique quantum numbers.
The implication of this work was to fully understand why the electron “shells” or principle energy levels held even number of  electrons.  The principle also extended to explain why electrons cannot occupy the same quantum state and thus must “stack” in the atom.  This “stacking” repeats and results in chemicals having different chemical properties based on their valence electrons – outermost electrons.
 
               Kernel = Core electrons = Most stable electrons
 
                Valence electrons = Unstable electrons – used in chemical bonding
 
Electrons fill a groundstate (stable lowest energy) atom by following the Aufbau Principle.  If they gain energy as in the case of photon of specific energy they will move away from the nucleus and into an excited state – (unstable high energy) state.  This state is identified by an electron configuration that is not following the Aufbau Principle and thus not filling a lower energy orbitals (closer to the nucleus) first before filling higher energy orbitals (farther away from the nucleus). The excited state configuration represent the electron arrangement of an atom just before it emits the photons in a bright line spectrum
 
Using orbital diagrams (boxes for the shapes of orbitals) we can also identify excited state configurations:
Remember that electrons in the same type of orbital will occupy empty orbitals first before pairing up – HUNDS RULE.
 

*Notice the electrons in the same box (orbital) have the opposite electron spins while the electrons in the same sublevel have parallel spins (going in the same direction).  These unpaired electrons lead to paramagnetism (the attraction to a magnetic field).

 
3. Review of electron configuration – overlapping of energy levels
4. Pauli Exclusion Principle
5.  Wrote the electron configuration of Uranium and Ag / Ag+
 
*Introduced the  term/concept of Zeff, loss of electrons, gain of electrons – will get too
                                period 4
 
1. Review Quantum Number Form 1 worksheet:
 
Quantum Number workheet 1 key p.pdf
View Download
2. Review/hand back homework
3. Review of electron configuration – overlapping of energy levels
4. Pauli Exclusion Principle 
5.  Wrote the electron configuration of Uranium/Shorthand/ and U+6
 
Todays Presentation for overlapping Principle Energy levels – slide 146 – 
PES – slide 49 – 56
Todays Presentation for PES – Photon Emission Spectroscopy: 

Photoelectron Spectroscopy ‎(PES)‎ basics

 
 
                  
1/29 – Wednesday 

 

Homework:  
 

1:  Please complete the form below:

    
Blue book Chapter 7 p.pdf
View Download
Powerpoint Presentation 1:
 
 
Blue BooK:
<
End of Wednesday..

1/30 – Thursday – 
 
Important Vocabulary that you will need in this unit:
 
Z  = #of protons    synonyms   Z = nuclear charge and   Z = Atomic Number (Thanks Moseley!)
 
Zeff = effective nuclear charge – (the nuclear charge that the electron feels as a result of electron – electron interactions ( screening or electron – electron repulsions).
 
n = principle energy level, the larger the n the larger the number of core electrons 
                                                      and larger the orbitals. n defines the proximity of electrons to the  
                                                      nucleus.  The farther that an electron is from the nucleus the lower the  
                                                      coulombic  attractions that the electron feels and thus is less stable           
                                                      than electrons closer to the nucleus.
 
Armed with ZZeff, and n you can explain almost everything in periodicity and electron  
                                                                                                                                           configurations.
 
****Since we are continuously evaluating the energy levels of electrons that are bound in a atom or ion in this unit Ionization energy values are very helpful in determining stability of an electron.  
 
Ionization Energy the energy needed to remove an electron (Einstein’s Binding Energy).
                                             Electrons with higher IE are more stable (takes more energy to remove!)
                                             Electrons with lower IE are less stable (takes less energy to remove!) 
 
Example for Na (sodium):                  IE1          +            Na     —->      Na+        +         e- 
                                                      first Ionization Energy
 
Ionization Energy is often described as the First Ionization Energy (1st IE) or the Second Ionization Energy (2nd IE) and so on…       

     

                                                                   IE2          +           Na+     —->      Na+2        +         e-
                                                              Second Ionization Energy

  

 
So the 2nd IE is the energy needed to remove the second electron.  Would it require the same amount as the 1st IE?  No it would require much more because Na+ the second electron would be removed from a filled principle energy level!! These are core electrons that are more stable. Do not lose site that IE is a measure of electron stability.
 
Stable Electrons =    High coulombic attraction to nucleus =    lower energy orbitals =   High IE
   Lower energy                           Lower n, Higher Zeff                                  closer to the nucleus
 
Successive IE values have verified the existence of Valence electrons!!! Look at the diagram below.
Notice when a successive IE “JUMPS THROUGH THE ROOF”.   

 

                    
                              Na (atom) :   1s22s22p63s                                            Na+1 (atom):  1s22s22p6
 
           Removing valence electron (less stable)                Removing a core electron (more stable)
                          3s electron has higher n                                       2p electron has lower n            
                         3s electron has lower Zeff                                    2p electron has higher Zeff
                                   Z = 11                                                                                       Z = 11
                           
                            IE =   500 kJ/mol          —— 9 x increase——->         IE2 = 4560 kJ/mole
                           removing valence e     “jumps through the roof”           removing core e
 
                                                               Thus Na has 1 valence electron                               
 
 
 – period 2
 
*Excited state of electrons!
 
1.  Classwork worksheet / review /discussed the notes above
 
                             “ALL of  This is about Energy “
 
Classwork:
Electron Config and Periodicity worksheet .pdf
View Download
 
Classwork Key:  I had 2 errors in the class key that displayed today. I fixed it below!
Electron configuration worksheet 1 Key p.pdf
View Download
2. REVIEW and Hand back Form 2 HW/review Form
 
3. Periodic Trends Lab 19 – 
File needed:

Periodic Trends Excel 3 graphs 2012 student file.xlsx

 
 
6.  Periodic Trends Lab 19 – 
File needed:

Periodic Trends Excel 3 graphs 2012 student file.xlsx

Periodic table Trends Lab 19 – activity instructions: (this will be classwork)
 
Please fill out the Ionization energy, Atomic Radii, and Electronegativity graph using table S of the OLD regents reference tables.
 
Please write and discuss every element who bucks the trend on the word document that you pasted the line graph for ionization. Remember that ionization energy is a measure of stability. Use your knowledge of Z, Zeff, and electron – electron interactions to justify why some atoms or groups of atoms are not following the trend (IE going up as you move across a period and IE decreasing as you move down a group). Please number and bullet each point.  There may be a group of elements that may not be following the trend. Please discuss them as well.
Graded as a lab activity. Do not hand this in yet.
 
                                – period 3/4 
 
*Bohrs orbits were really s orbitals
 
1.  Classwork worksheet / review /discussed the notes above
     Key is posted above.
 
Period 4  “ALL of  This is about Energy ” Lecture.
 

 

 
                             “ALL of  This is about Energy “
 
 
Powerpoint Presentation 2:

 

Electron Affinity Animation:

 
1/30 – Thursday Homework:  
1:  Please complete the form below:
 
                
 End of Thursday!                         

1/31 – Friday –  PES!
 
              Ionization Energy          vs.             Electron affinity
 Ionization Energy (IE) Electron Affinity (EA)
 Energy needed to remove an electron  Energy released or absorbed when electron is added
 measures stability of current electrons  measures stability of added electron
            Creates positive ions (cations)                     Creates negative ions (anions)
   ∆H = positive (endothermic)  ∆H = negative (exothermic) or positive (endothermic)
  .50 kJ/mol     +    Na    —–>     Na+     +      e            F          +        e    ——>      F–      +    328 kJ/mol
 1681 kJ/mol    +    F      —–>     F+        +      e-     53 kJ/mole  +   Na        +       e     ——>      Na
 Larger the IE the more stable the e  Larger the negative EA the more stable the added e
 Used for all atoms – Clear Trend  Used primarily for nonmetals but Trend is not clear/ many exceptions

 

 
Given the following EA for the Halogens – group 17

 

Fluorine (F) -328 kJ/mol
 
Chlorine (Cl) -349 kJ/mol
 
Bromine (Br) -324 kj/mol
Iodine (I) -295 kJ/mol
 
 
The EA “generally decreases” down a group because the increased shielding that occurs with more orbitals of electrons (core) between the outermost electrons (valence) when n increases is offset by the larger Z that occurs as you move down a group.  As you move down a group n is the biggest factor why the outermost electron become less stable and held more loosely.  That is why valence electrons are less stable than core electrons.  
 
With Fluorine we would expect it to have the highest EA of the group since its valence electrons are in the smallest n (n= 2) and should release the greatest amount of energy (show more stability) as it grabs one electrons BUT IT DOES NOT.  Because Fluorines electrons exist in a very small space with n= 2 the extra electron will be destabilized a bit by the electron – electron repulsions that will occur in this small space.  The Zeff for this electron that is added will not be as high as we would expect because of the crowded small space for the electron in the second principle energy level.  The rest of the group, follows the expected trend because their valence electrons in exist in larger and larger orbitals as the n increases resulting in lower Zeff due to the increased distance from the nucleus.
 
EA like IE also has EA2 and these values are almost always VERY positive as it will take energy to add an electron to an already negative particle (unless the Z is large enough to offset). This never the case for small values of n.
            Period 2 – 
 
A):  Properties of transitional metals-
      Diamagnetism/Paramagetism  8-3,8-54 to 8-56 slides in powerpoint 1
 
B):
Reviewed the questions on the HW form 3 by 
1.  The 4s fills first due to its shape and its penetration closer to the nucleus.
2. Divergence/crossover of 4s orbitals to explain electron configurations of transitional elements that do not seem to follow the Aufbau Principle:  Examples Cr and Cu.
 
3. Divergence also explains the these elements have multiple oxidation states 
    AND ALL lose their electrons in the s orbital first because it is now a higher energy unstable orbitals compared      to the filling d orbital.
 
 4, Handed back graded Form 3…          
 
         Period 3/4 – 
 
A):  Properties of transitional metals-
      Diamagnetism/Paramagetism  8-3,8-54 to 8-56 slides in powerpoint 1
 
B):
 
Reviewed the questions on the HW form 3 by 
1.  The 4s fills first due to its shape and its penetration closer to the nucleus.
2. Divergence/crossover of 4s orbitals to explain electron configurations of transitional elements that do not seem to follow the Aufbau Principle:  Examples Cr and Cu.
 
3. Divergence also explains the these elements have multiple oxidation states 
    AND ALL lose their electrons in the s orbital first because it is now a higher energy unstable orbitals compared      to the filling d orbital.
 
 4, Handed back graded Form 3…
 
Divergence Lecture if you cannot get enough:
Todays Presentation for PES – Photon Emission Spectroscopy: 

Photoelectron Spectroscopy ‎(PES)‎ basics

 
1/31 – Friday (weekend) Homework: 2 parts
* Next Tuesday Electron Configuration and Periodicity Test.
 
1. Periodic Trends Graphing Lab
2. Lecture on PES/ Complete Form
 
1. Periodic Trends Lab 19 – 
File needed:

Periodic Trends Excel 3 graphs 2012 student file.xlsx

Periodic table Trends Lab 19 – activity instructions: 
 
Please fill out the Ionization energy, Atomic Radii, and Electronegativity graph using table S of the OLD regents reference tables.
 
– Table S of the Regents Reference tables contains the values for First Ionization EnergyElectronegativity, and Atomic Radii. Please make 3 graphs by using my excel file above (it will make a line graph and a 3d graph) of the three trends.  All you have to do is enter the values in the spreadsheet up to atomic # 54 (Xe) for the three Periodic Trends.  Please cut and past the 3 linear graphs into doc that you can print and follow the instructions below:
 
Electronegativity that measures how much an atom attracts electrons IN A BOND.  It really is a value that combines Z, Zeff, and n. You will notice that the Nobel Gases (last column on the periodic table) do not have Electronegativity values because they do not bond so their attraction to electrons in a bond cannot be measured.
Elements with dashes in the table have zero values.
 
For the Electronegativity and Atomic Radii Graphs please Identify the elements that belong to the Alkali family (group 1), Alkaline Earth (Group 2), Halogens (Group 17), and the Noble gases (group 18).
 
For the First IONIZATION ENERGY GRAPH:
 
Please write and discuss every element who bucks the trend on the word document that you pasted the line graph for ionization. Remember that ionization energy is a measure of stability. Use your knowledge of Z, Zeff, n, and electron – electron interactions to justify why some atoms or groups of atoms are not following the trend (IE going up as you move across a period and IE decreasing as you move down a group). Please number and bullet each point.  There may be a group of elements that may not be following the trend. Please discuss them as well.
 
Graded as a lab activity. 
 
2. PES (Photon Emission Spectroscopy) – experimental evidence of electron configurations.
 
Lecture: 

 

Form: 
 

PES – Photon Emission Spectroscopy Form

 

 
End of Week 1 of the 3rd Quarter..