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Q3: Week 4  – 2/28 – 3/4

                                                                                                                                                                                 Jump toTuesday,   Wednesday,  Thursday,  Friday                                                                                                              ______________________________________________________________

2/28 – Monday B Day – 2, 3b/4 Lab 

Main focus –                                                                                                                                                         

    a) To Review the properties of transitional metals

    b) To Review the isoelectric series.


Period 2:  –  

1.  Complete the backside of Electron Config and Periodicity worksheet.

Electron Config and Periodicity worksheet 3.pdf
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Electron Configuration worksheet 3 Key p.pdf
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2. Introduce the Isoelectric Series and complete classwork.      

Isoelectric series – 
Electron Config and Periodicity worksheet 4.pdf
Electron Configuration worksheet 4 Key p.pdf
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Period 3/4:  – 

1.  Same as 1 and 2 above.

2. Lab 21 – Periodic trends Graphs



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!



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.   
PES DATA From Lecture: 

Aufbau Principle Flask Energy inversions :

Quantum Numbers to Electron Configurations:


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).
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.
NOT a Formal Lab!



2/28 – Monday Homework: –                                                                                                                                                                                                                           

1:  Please complete the following form that is on auto-reply as a review for tomorrows test on atomic structure.


2: Please Review the following atomic structure worksheets.   You could also review with the Atomic Structure        Review page posted in quarter 3.
Atomic structure 2 – bohrs.pdf

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atomic structure 2 – Bohr Key p.pdf

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Photoelectric Effect – Form Key.pdf
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atomic structure 3 – de Broglie Key.pdf
Atomic structure 3 – de Broglie.pdf

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Quantum Number form 1920 Key p.pdf
Atomic Structure 4b -Electron configuration.pdf
Atomic Structure 4b -Electron configuration key.pdf
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Electron Config and Periodicity worksheet .pdf
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Electron configuration worksheet 1 Key p.pdf
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Electron Config and Periodicity worksheet 3.pdf
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Electron Configuration worksheet 3 Key p.pdf
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Electron Config and Periodicity worksheet 4.pdf
Electron Configuration worksheet 4 Key p.pdf
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PES – Photon Emission Spectroscopy Form – key p.pdf
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2 : AP Periodicity and Electron Configurations Form 2:

End of Tuesday!


3/1 – Tuesday – A Day – 2/3a Lab, 4 

Main focus –                                                                                                                                                         

    a) To take an Exam on Atomic structure and Periodicity

 Period 2/3: 

1. Lab 21 – Periodic trends Graphs

2. Atomic Structure Test

 Period 4: 

1. Atomic Structure Test                                         



3/1 – Tuesday Homework: –                                                                                                                                                                                                                           

1. Please complete the take-home portion of the test (page 3). It is due tomorrow.


3/2 – Wednesday – B Day – 2, 3b/4 Lab 

Main focus –                                                                                                                                                         

    a) To Define Metals, Non-metals, and metalloids in terms of atomic radii, First                           Ionization Energy, and Electronegativity.

    b) To introduce the basics of bonding principles.

    c) Introduce the Lewis Dot Diagrams  for Ionic Compounds and covalent compounds.

Period 2:  

1. Metals vs. Nonmetals using the Periodic trends 3-D graph.
– marked a periodic table with metals, nonmetals, and metalloids     
– Defined metals and nonmetals based atomic radii, electronegativity and Ionization Energy
– started ionic bonding
Blank periodic table with Reference Periodic Table .pdf
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2.  Born- Haber Lattice Energy
3. Ionic Bonding – naming /honeymooners
4.  Covalent Bonding setup. 
Period 3/4:  
1: Periodic Table development 
            Tom Lehr
            Alkali reactivity
            atom shack
2. Metals vs. Nonmetals using the Periodic trends 3-D graph.
– marked a periodic table with metals, nonmetals, and metalloids     
– Defined metals and nonmetals based atomic radii, electronegativity and Ionization Energy
– started ionic bonding
3.  Born- Haber Lattice Energy
4. Ionic Bonding – naming /honeymooners
5.  Covalent Bonding setup. 
– completed the alkali reactivity presentation with the reaction between Cs and F.


Periodic Trends from our Lab.
 Where’s Waldo?  Where is Hydrogen? 
 Where are the small atoms (non-metals) and where are the large atoms (metals)?


Family of elements – same valence electrons
– alkali metals – group 1 – 1 valence electrons —————> all become +1 ions
– alkaline earth metals – group 2  – 2 valence electrons—–> all become +2 ions
– halogens – group 17 – 7 valence electrons ——————->all become -1 ions
– noble gas – group 18 – 8 valence electrons ——————> most do not become ions
– Periodic trends:  atomic radii, and Ionization Energy,  
    AND Electronegativity = attractions for other elements electrons.
    Electronegativity is really a term that combines (Z, Zeff, and n).  It is easier to describe an elements attraction 
    for other atoms electrons IN A BOND!  The electronegativity scale for all elements (0 – 4) has been                         determined for all elements.  Fluorine has the highest (4.0) and Cesium has the lowest (.8).  
    *The Noble Gases do not have a value for electronegativity ( 0 ) because they do not form bonds and it cannot be  measured.  You will notice in your periodic table that Kr and Xe do have small electronegativity values and well as possible oxidation states due to their small reactivity due to the larger n!

Reactivity of Metals (BIGS!):

Bonding Presentation:


3/2 – Wednesday Homework: – 

1.  Please watch the lecture below that reviews both types of bonding and how to complete the lewis diagrams for each.
2. Complete the back side of the ionic bonding worksheet. Review with the key.
Ionic ditto 1 -Electron Dot Diagrams .pdf
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Ionic ditto 1 -Electron Dot Diagrams KEY.pdf
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3. Please complete the 1st column of electron dot diagrams for covalent compounds on the backside of the covalent worksheet and review with the key.
Covalent ditto 1 -Electron Dot Diagrams .pdf
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Covalent ditto 1 – Electron Dot Diagrams key .pdf
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Ionic and Covalent Bonding Basics:



3/3 – Thursday – A Day – 2/3a Lab, 4 

Main focus –                                                                                                                                                         

    a) To Review Ionic Bonding and their crystal lattice structures that their empirical                  formula represents.

    b) To Identify Molecular formulas and they are exclusive for Molecular compounds.

     c) To Identify Dimetry Mendeleev as the author of the period table (Period Law)

Period 2/3:  – 

1: Periodic Table development 
            Tom Lehr
            Alkali reactivity
            atom shack
2.   Metallic vs. Nonmetallic reactivity
Alkali Reactivity Presentation ——> Ionic Bonding 
3.  Ionic Lewis Dot Diagrams
 Period 4:  – 
1: Periodic Table development 
            Tom Lehr
            Alkali reactivity
            atom shack
2.  Ionic Lewis Dot Diagrams




Bonding Presentation:

Lattice Energy presentation:


3/3 – Thursday Homework: – 

1. Please view the Hybridization Lecture up to 8:05 only.
2. Watch the sp hybridization Lecture in its entirety and follow along with me using the worksheet below.
We will be learning the following.
       a) sigma bonding
       b) pi bonding
       c)  sp hybridization
       d) lewis dot diagrams
       e) polarity
SP hybridization.pdf
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SP hybridization KEY p.pdf
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3. Then read Grodski Notes below!
4. Complete Form below. It will be on auto-reply with a total of 3 submissions.

2 :Hybridization Lecture :

2 :sp hybridization worksheet Lecture :

3 :TODAY’s Homework NOTES:
We have been learning a valence bond theory that explains almost all the experimental data that we obtain in terms of bond angle (x-ray crystallography), bond energies, and polarity for molecules (covalently bonded particles).  We can predict very successfully the above listed properties of molecules by just putting dots on a piece of paper and writing Lewis structures, using VSEPR theory (how electrons in orbitals repel each other to find a stable shape), and utilizing the concepts of hybridization theory.  Remember Hybridization it is just a theory that explains how molecules achieve the known experimental data, using our known understanding of electrons acting in waves in atomic orbitals.  Think that the atomic orbitals that we have learned BECOME new orbitals when they bond.  This should make sense as electrons act as waves and should constructively interfere with other orbitals of atoms they bond with.
Remember that we use the hybridization theory for central atoms BECAUSE we need to!!! The bond angles and bond energy that we predict for using the different atomic orbitals of the central atom cannot be explained unless we believe that these valence orbital mix or hybridize!!!
Lets take a look at the BeF2 orbital sketch below.  BeF needs to be sp hybridized to explain its linear structure that we know experimentally.  In order for it to bond on BOTH SIDES YOU NEED 2 separate orbitals and the 2 central sp hybridized orbitals in the central Be allow for this.  Notice the terminal F atoms do not need hybridization to explain their terminal bonding. All they are doing is filling their outermost 2p orbital.  Remember unhybridized p orbitals cannot bond on both side since EACH P orbital hold 2 electrons and if it bonding on one side it is now full. This is why each F atom on the outside does not have any electrons on the outer porbital BECAUSE it full due to the covalent bonding with the central Be sp hybrid orbitals. Figure 2 below highlights the px orbital from the F atom illustrating that the right and left side ( of each dumbell ) is the same orbital orbital which is now full due to the bonding with central atom (Be).
figure 1

*the 2s orbital is drawn much smaller than the 2p to help view the overlap of orbitals in covalent bonds.

Notice that the central atom has 2 Be – F bonds and they measured in strength called bond enthalpy. These bond enthalpy’s (the energy needed to break a bond) have been determined experimentally to be the same. If there was no hybridization then they would not be the same as one electron in a 2s orbital of Be would bind with one F while another electron in the 2p would bond with the other F.  Because an s orbital is more stable than a p orbital, a bond in the s orbital would have a higher Bond Enthalpy (much like a more stable orbital has a higher IE!).
Do not lose site that these orbital diagrams are just visualizations of what we think the new bonding arrangement of of atoms joined together by using each other to fill orbitals (bonding) looks like.  You can see that the solutions to the Schrodinger equation gives a way to explain the bonding between atoms visually. Remember that these orbitals are wave functions that represent electrons in standing waves or in electron density probabilities.
figure 2
Now you will notice the direct overlap between the F atoms px orbital and the Be sp hybrid orbital so that 2 electrons are shared between each orbital.  This type of bonding of sharing valence electrons to fill valence orbitals is called covalent bonding.  The type of covalent bonding that has a direct overlap of orbitals is called sigma (𝞼) bonds. We also learned about the second type of bonds called pi (π) bonds that occur with UNHYBRIDIZED p orbitals of the central atom with unhybridized p orbitals of the non central orbitals. Notice that Be ONLY had a 2s and 1 2p orbital that hybridized and there we no other electrons in the 2p so there ARE NO CENTRAL unhybridized p orbitals for the central atom this molecule does not have any π bonding. π bonding will only occur as the second or the third bond but never as a bond on its own.
4 : Hybridization Form 1 – 2021


End of Thursday!


3/4 – Friday B Day – 2, 3b/4 Lab 

Main focus –                                                                                                                                                         

    a) To identify sp and sp2 hybridization. molecular and electron domain geometries,                bond angles, by writing Lewis Dot Diagrams for Molecular compounds.

    b) To review sigma bonding and pi bonding

    c) To identify polar and nonpolar covalent bonds.


Period 2, 3/4– 

1. Why the theory of hybridization (FOR THE CENTRAL ATOMS) is needed by writing the Lewis Diagram for BeF2.

     a) Review of last nights last nights last question.

       Lewis Diagrams —-> identification of hybridization, sigma bonds, pi bonds, polar bonds, dipole moments, polarity

SP hybridization KEY p.pdf
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2.  Sp2 Hybridization introduction:

sp2 hybridization.pdf
sp2 hybridization KEY p.pdf
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In the sp2 hybridized structures that we will learn about tonight, we will see that electron domain geometry will not always be the same as the molecular geometry.  
Electron domain geometry is the shape of the (bonded and non-bonded) electrons that surround the central atom 
while the 
Molecular geometry is the shape of where the atoms are bonded around the central atom independent of the electrons.  This means that we do not consider the lone electrons but only the atoms. Be careful here.  The electrons surrounding the central atoms (bonded or non- bonded) determine the hybridization and the general shape of the hybridized family.  Electrons repel themselves in the central atom to find the most stable structure, (VSEPR theory) but we only look at where the atoms are positioned in bonds around the central atom to determine the shape of the molecule. 
In the sp hybridized family (Thursday’s homework), the electron domain geometry ALWAYS was the same as the molecular geometry because the sp molecular shape are always linear BECAUSE THERE ARE NO LONE ELECTRONS to repel the electrons in the central atom.  Because there can be no lone pairs in the central atom for sp hybridization the shape will always be linear for the bonded electrons and the atoms that are bonded to these electrons. 
Take a look at the diagram below to view how the sp2 family has a molecular geometry that is different from the  electron domain geometry for one of the possible geometries of molecules that are sp2 hybridized.  Keep in mind that the sphybridized molecules have one shape where the electron domain geometry and the molecular geometry are the same (trigonal planar) but also has one other shape where they are not the same (unlike the sp hybridization family).
This will be made clear in tonight’s lecture.



3/4- Weekend Homework: – 

1.  Please watch the lecture on sp2 and follow along with me with worksheet below and complete the Hybridization Form 2 Below: 
sp2 hybridization.pdf
sp2 hybridization KEY p.pdf
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      1 :sp2 hybridization worksheet Lecture :


2 : Hybridization Form 2 – 21-22