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## Q3 : Week 2 – 2/5 – 2/9

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#### Refresh this page every time you arrive!  I update this daily!!!

If you have not already please join the REMIND for this class.

– Regarding Homework assignments – IF you cannot complete the assigned homework before school the next day you must email me.

Please Make sure Your read all instructions for all assignments.

Make Sure Powerschool grades are accurate!!

___________________________________________________                                                     Jump to: Monday Homework / top 2/5 – Monday A Day 2, 3a Lab/4

Main focus –

a) To use the solutions to the Schrodinger Equation to write electron configurations.

b) To define the quantum numbers to standing waves of orbitals (wave potentials).

c) To build an orbital diagrams.

Period 2,3a,4:

1. Shrodinger’s equation —> quantum numbers

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!”

2.  Review the quantum number form.

Quantum Number form 1920 Key p.pdf

3.  Complete the Quantum numbers worksheet with the class to cements the concepts of quantum numbers that are solutions to the Shroedinger’s Equations.

Quantum Numbers – worksheet 1.pdf

Quantum Numbers – worksheet 2.pdf

Atomic Structure 4b -Electron configuration.pdf

4.  Wrote the configuration of the Uranium to demonstrate how the periodic table is a guide to the electron configurations (especially the overlapping of energy levels).

Need to show students the shapes of orbitals and where the g begins!!

TODAY’s NOTES:

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)
m= -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 = 3 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?

*ConnectionsOur 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!!!!

Need to show students the shapes of orbitals and where the g begins!!

Light into the Bohr Model:

___________________________________________________                                                                                                               jump to: top               2/5 – Monday’s Homework: –

1.  Please complete the Quantum Numbers – worksheet 1.pdf if you did not do so in class and review with the key below: (period 2, 3a needs to finish the last 2 and review with the key below)

(period 4 finished in class!

Quantum Numbers – worksheet 1.pdf

Quantum Numbers – worksheet 1 Key.pdf

2.  Please view the lecture of me reviewing the first 4 elements in the Quantum Numbers – worksheet 2.pdf and then complete the worksheet UP TO Ca (CALCIUM) ONLY!!. Review with the key below.

Quantum Numbers – worksheet 2.pdf

Quantum Numbers – worksheet 2 Key.pdf

3. Please write all the configuration from the worksheet below UP TO Ca (CALCIUM ONLY!) and review with key or with the 2 lectures posted at the bottom of this page if you need help.

Atomic Structure 4b -Electron configuration.pdf

Atomic Structure 4b -Electron configuration key.pdf

Review of quantum number worksheet 2 (the 4 atoms) : Start at 17:30 – If you need more review of what these quantum numbers mean you can listen to me reviewing that before 17:30.

 electron configuration 1 electron configuration 2

____________________________________________________________________                                Jump toTuesday Homework / top  2/6 – Tuesday – B Day – 2, 3b/4 Lab

Main focus –

a) To write electron configurations for atoms and ions using the quantum numbers

b) To identify filled, occupied, and partially filled principle energy levels, sublevels,                and orbitals.

Period 2: –  Shape of orbitals! (presentations and flash)

1.  Review the quantum numbers worksheet 1.

Quantum Numbers – worksheet 1 Key.pdf

2. Review the quantum numbers worksheet 2.

a) Overlapping of principle energy levels (orbitals) – Scandalous!!!

Quantum Numbers – worksheet 2 Key.pdf

3.  Writing Electron Configuration of Uranium – writing configurations lesson – shorthand!

– Wrote the electron configuration of Uranium and Ag / Ag/Shorthand/ and U+6

Atomic Structure 4b -Electron configuration.pdf

Atomic Structure 4b -Electron configuration key.pdf

Period 3b, 4: – Shape of orbitals! (presentations and flash)

1.  Quantum numbers Review with presentation

2. Review the Quantum Numbers worksheet 2.

3. Review of electron configurations

TODAY’s NOTES:

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)
m= -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 = 3 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 Nitrogen:           Nucleus:  1s22s22p3

Orbital notation for N:
Each box represents a single orbital.

The orbital diagram of Nitrogen above reveals a new situation where electrons will occupy empty orbitals OF THE SAME SUBLEVEL before they pair up in a single orbital.
This did not occur in the s sublevel because there is only 1 orbital in the s sublevel.

IF l = 0 (s sublevel) → ml = 0 (only 1 solution thus only 1 orbital for the s sublevlel)  thus a second electron had no other choice but to pair up in an s orbital due to the Aufbau Principle. That is this second electron was more stable (lower in energy ) in a paired s orbital than a p orbital of higher energy.

If the electron in an sublevel with multiple degenerate orbitals (different orientations in space) will occupy an empty orbital before pairing up = Hund’s Rule.

*ConnectionsOur 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!!!!

Need to show students the shapes of orbitals and where the g begins!!

Aufbau Principle Flash Energy inversions :

Quantum Numbers to Electron Configurations:

___________________________________________________                                                                                                                    jump to:  top         2/6 – Tuesday’s Homework: –

1.  Please complete Atomic Structure 4b -Electron configuration.pdf worksheet and review the key if you have not done already.

Atomic Structure 4b -Electron configuration.pdf

Atomic Structure 4b -Electron configuration key.pdf

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 snipit for the homework:

Blue book Chapter 7 p.pdf

Powerpoint Presentation 1:

Blue BooK:

. \

2 : Tonight’s Form:  AP Periodicity and Electron Configurations Form 1 23 – 24

I will have this homework graded and returned back to you tomorrow morning.  YOU WILL ONLY HAVE ONE RESPONSE! If you do not complete this it I CANNOT ACCEPT IT LATE AS I WILL BE GOING OVER IT as part of my lesson tomorrow.

End of Wednesday!

__________________________________________________________________________                           Jump to: Wednesday Homework / top  2/7 – Wednesday A Day 2, 3a Lab/4

Main focus –

a) To Review Groundstate vs. Excited electron configurations.

b) To define Z, Zeff, and Ionization energy.

c) To review electron screening and electron – electron repulsions (Zeff)

Period 2,3a/4:

1. A review of the Atomic Structure 4b -Electron configuration.pdf worksheet (electron configurations).

Atomic Structure 4b -Electron configuration key.pdf

a) Go through the complete electron configuration of entire periodic table with                                                                                            slide 154 – g orbitals –> groundstate vs. Excited state

2. A Review of the Quantum number worksheet 2. (reviews overlapping of principle energy levels)

Quantum Numbers – worksheet 2 Key.pdf

Electron Config and Periodicity worksheet 2.pdf                                                                                  View Download

3. Review of last nights Form.

AP Periodicity and Electron Configurations Form 1 KEY p.pdf                                                                                     View Download

4.   Define, Z, Zeff, electron screening and electron – electron repulsions.

TODAY’s NOTES:

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 you can explain almost everything in periodicity and electron
configurations.

****Since we are continuously evaluating the energy levels of electrons that are bound in an 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…

– The Second Ionization Energy is the Energy needed to remove a second electron (after a single electron was removed)

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

– Wrote the electron configuration of Uranium and Ag / Ag/Shorthand/ and U+6

“ALL of  This is about Energy “

ALL of this is about Energy Lecture!

Need to show students the shapes of orbitals and where the g begins!!

Light into the Bohr Model and beyond!

_____________________________________________________.                                                                                     jump to:  top          2/7 – Wednesday Homework: –

1.  Please complete the Electron Configuration and Periodicity worksheet 2 and review the key.

Electron Config and Periodicity worksheet 2.pdf                                                                                                      View Download

Electron Config and Periodicity worksheet 2 key.pdf                                                                                                      View Download

2.  Please complete the Form below based on the resources posted above the form.  Please answer the questions completely as if you were answering a free response question. Do not be vague but get to your point quickly!    A longer response is not always the correct reponse.

Powerpoint Presentation 1:

Blue BooK:

2 : Tonight’s Form:  AP Periodicity and Electron Configurations Form 2 23-24 –

Please use complete sentences and DO NOT BE VAGUE!!  Be careful to answer each question in each numbered question. There is often more than one question per numbered question.  Writing more does not guarantee the answer will be correct.  The real skill is to write the least amount of words that will answer the question completely.  Get in and get out!

________________________________________________________________________                           Jump to: Thursday Homework / top  2/8 – Thursday – B Day – 2, 3b/4 Lab

Main focus –

a) To explain the trends in the periodic Table in terms of Z, n, and Zeff.

b) To compare and contrast Ionization Energy with Electron Affinity.                                    c) To identify paramagnetic and diamagnetic properties of atoms from electron                       configurations,

Period 2/3a:

1. Complete the Review the AP Periodicity and Electron Configuration Form 1 with key.

AP Periodicity and Electron Configurations Form 1 KEY p.pdf                                                                                     View Download

2. Review the AP periodicity and Electron Configuration Form 2 with key.

AP Periodicity and Electron Configurations Form 2 – KEY.pdf

a) Paramagnetism
b) Electron affinity

3.  Valance electrons AND “Bam!!!” IE energy values = the proof of valance electrons!

4.  Begin classwork worksheet:

Classwork:
Electron Config and Periodicity worksheet .pdf

Classwork Key:
Electron configuration worksheet 1 Key p.pdf
Period 4:

1. Review the AP Periodicity and Electron Configuration Form 1 with key.

AP Periodicity and Electron Configurations Form 1 KEY p.pdf                                                                                     View Download

2. Review the AP periodicity and Electron Configuration Form 2 with key.

TODAY’s NOTES:

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 (screening) 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.

Paramagnetism – weak attractions to a magnetic field.
Diamagnetism –  No attractions to a magnetic field.

Electron Affinity Animation  :

Last Slides for orbital shapes!

__________________________________________________                                                                                            jump to:  top            2/8 – Thursday Homework: –

1. Please complete the AP Periodicity and Electron Configuration Form 3 Below using both Power Point 1 and Powerpoint 2 below:

Powerpoint Presentation 1:

Powerpoint Presentation 2:

Blue BooK:

1 : Tonight’s Form:  AP Periodicity and Electron Configurations Form 3 22-23 –

Please use complete sentences and DO NOT BE VAGUE!!  Be careful to answer each question in each numbered question. There is often more than one question per numbered question.  Writing more does not guarantee the answer will be correct.  The real skill is to write the least amount of words that will answer the question completely.  Get in and get out!

End of Thursday!

_________________________________________________________________                                         Jump to: Friday Homework / top 2/9 – Friday A Day 2, 3a Lab/4

Main focus –
a) To explain orbital diagrams of transitional metals through diverging orbitals

b) To review the electron shielding vs electron – electron repulsions

Period 2,3a:

1.  Review the AP Periodicity and Electron Form 3 with the key.
2.  Divergence with s and d to explain +2 cations and exception orbital diagrams.
3.  properties of Transitional metals

4.  Complete the following worksheet:

5.  Lab 21 – Periodic Graphs excel – If time permits..

Classwork:
Electron Config and Periodicity worksheet .pdf

Electron configuration worksheet 1 Key p.pdf

Electron Config and Periodicity worksheet 3.pdf

Electron Configuration worksheet 3 Key p.pdf

Period 4:

1 – 3. : Same as above

4. Classwork 1:

Electron Config and Periodicity worksheet 3.pdf

Electron Configuration worksheet 3 Key p.pdf

TODAY’s NOTES:

Properties of transitional metals –

We learned that divergence of the 3d orbital is responsible for the properties of the transitional metals which are the elements in the d – block.  The d 0rbital unlike the s and the p holds a maximum of 10 electrons and combined with the outermost s orbital that is very close in energy with the outermost d orbital provides a “super sublevel” where there 12 electrons reside in what becomes sort of a valence shell for these elements.

1. High Conductivity of Electricity – High number of mobile electrons (low IE) in metallic bonding

2. The Largest Paramagnetism  Largest number of degenerate orbitals that could contain the
largest number of unpaired paralleled spin electrons.

3. Multiple Oxidation States –  Many choices for stability of electrons based on minimizing
electron – electron repulsions given the 6 orbitals (s and d) that
electrons can move to and from. Transitional elements cannot
achieve noble gas configurations because they cannot lose or gain
the high number electrons that this would require.
Fe would have to LOSE 8 electrons OR Gain 10 electrons to achieve
Kr or Xe electron configurations. Fe has too high of a Z to lose 8
electrons and its Z is not high enough to gain 8 electrons.

4.  Valence Electrons from Multiple Principle Energy Levels (n) –
Electrons are lost by metals because of relatively low IE but
electrons  lost by d – block metals are from the “super sublevel”
of (n) s and (n- 1) d electrons that are very similar in energy.

5.  Form Colored Solutions –   Crystal Field Splitting!!!!!   Remember!!!!

Because they can have high oxidation states due the large number of electrons in their “super sublevel” they can draw electrons pairs from other molecules (ligands) to form stable complexes that cause the degenerate d orbital to split into 2 energy levels.  This splitting of the d orbital based on the electrons being drawn into the d orbitals of the d – block metal because of the large coulombic attraction of d block elements that have high Z and large oxidation states creates an opportunity for these elements to absorb photons of visible light (negative theory of light) resulting in the complexes that transmit a photons of light that is not absorbed.

Mn+7  in  KMnO4 = purple (in the oxidation titration lab) ——>  Mn+2  (Clear)

Cu oxidized by nitric acid (in the % by mass of copper in Brass lab) = blue green solution and we used a spectrophotometer to measure how much light is missing (absorbed):
*Remember Cu solutions are blue-green because they make complexes in water:

Cu+2   +   6H2O     —–>     [Cu(6H2O)6]+2

 The d orbitals that will interact directly with the incoming ligand (electrons from oxygen in water) will Destabilize that d orbital because of electron – electron repulsions and thus that orbital will contain electrons with lower Zeff.  The d orbitals that do not directly interact with the incoming ligand are not as destabilized and thus a GAP is created and the degenerate d orbitals are split into 2 levels by a gap small enough in energy that photons of visible light can match!

FROM WEEK 6 of QUARTER 2!
Connections:
In Lab 13 – Volumetric REDOX Titrations – We used Fe+2 ions to Standardize the KMnOsolution.
The MnO4  solution  oxidized the Fe+2  into Fe+3 .  The Fe+2  reduced the The MnO4  solution to turn from purple (Mn+7 form) to colorless (Mn+2).

IF the the Fe+2  solution sits in volumetric flask it slowly changes into Fe+3 as oxygen in the flask and in the solution oxidizes the the Fe+2 solution. Notice the development of color in the 24 hour period.

I have to make the  Fe+2  solution fresh each day.  Why is there a color developing?

As the Fe+2  becomes  Fe+3  the solution has a greater ability to absorb wavelengths of visible light because Fe+3 makes a complex ion with water that attracts water with such high Coulombic attractions that the water’s lone pair of electrons are pushed into the iron’s (Fe) d – orbitals directly interfering with some d – orbitals and not others based on their orientation in space.  The orbitals that “feel” the oxygens electrons become destabilized (increase in energy) and there is a spit in the energy level between these orbitals that were initially the same energy.

This splitting of d orbital energy levels provides a pathway for electrons to transition to higher energy levels when low energy visible light energy (photons) are absorbed.  Not all of the wavelengths of light are absorbed thus the ones that are not absorbed are what we see.
The spectrophotometer picks up the transmittance (which is the what is absorbed or missing from our eye.)

A photon of light of a certain wavelength is absorbed by electrons in lower energy d orbitals that can now transition into higher energy d orbitals (that split due to electrons being pulled directly into the orbitals). The absorbed photon is now missing from the entire spectrum of light that is illuminating the complex and the color shown is what is left.

Aufbau Principle Flask Energy inversions :

Light into the Bohr Model:

______________________

Lab 21 – Periodic Trends Graphing

File needed:

Periodic Trends Excel 3 graphs 2012 student file.xlsx

Periodic table Trends Lab 20 – 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.

Divergence (energy inversion) Lecture of 3d orbitals :

______________________________________________                                                                                                     jump to:  top        2/9 – Friday Homework: –

1.. PES (Photon Emission Spectroscopy) experimental evidence of electron configurations.

All that we have been learning regarding the arrangement of electrons in the Current Model of the atom – Wave Mechanical or Electron Cloud would be pure nonsense UNLESS there was no experimental evidence to support it.

The Three most important experiments include:

Ionization Energies 1st, 2nd, etc. – The WHOA amount that proves Core electrons and Valence e.
Also proves that electrons are filling energy shells as we move across the
periodic table.
Magnetic Behavior –  measures a whole integer of magnetic force due to electrons with unpaired
electrons having the same spin that create a magnetic moment – proves
that orbitals of the same sublevel are degenerate.

PES (Photon Emission Spectroscopy) An application of the photoelectric effect that will measure  the energy and the relative amount of electrons each energy level in an atom or ion. It proves the electron configurations are correct as we know it OR the Aufbau Principle.

Please watch my lecture on the PES  and compete the form below that will be on auto-reply.

Photon Emission Spectroscopy Lecture :

1 : PES – Photon Emission Spectroscopy –

PES DATA From Lecture:

End of Week 2!