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Archive – Q3 week 2 – 19-20

lithiumWeek 2/3 – 2/7


2/3 – Monday –  period 2

1.  Review of the PES Form- 
2.  Properties of transitional metals-
      a) Diamagnetism/Paramagetism  8-3,8-54 to 8-56 slides in powerpoint 1
      b) colored solutions; Crystal field Theory
      c) MULTIPLE VALENCE ELECTRONS from 2 different energy levels!
ClassWork: (complete or start)
Electron Config and Periodicity worksheet .pdf
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Classwork Key:  
Electron configuration worksheet 1 Key p.pdf
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Electron configuration writing with Diamagnetism
Electron Config and Periodicity worksheet 3.pdf
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Electron Configuration worksheet 3 Key p.pdf
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3.  Isoelectric series – 
Atomic structure 5 screening in graph H He.pdf
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Electron Config and Periodicity worksheet 4.pdf
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Electron Configuration worksheet 4 Key p.pdf
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                            – period 3/4
1.  Review of the PES Form- 
2.  Properties of transitional metals-
      a) Diamagnetism/Paramagetism  8-3,8-54 to 8-56 slides in powerpoint 1
      b) colored solutions; Crystal field Theory
      c) MULTIPLE VALENCE ELECTRONS from 2 different energy levels!
3. Classwork: worksheets posted above!
5. Periodicity Excel Lab – only period 3/4 
Properties of transitional metals – notes
 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!

PES: atomic structure 2 presentation slides 47 – 53.
Todays Presentation for PES – Photon Emission Spectroscopy: 

Photoelectron Spectroscopy ‎(PES)‎ basics

Mystery element from PES data:
2/3 – Monday -Homework – Study for Atomic structure / Periodicity Test
1. Complete all classwork worksheets that was posted above and review with keys above.
2. Study for the test.  Posted study materials are in the Atomic structure page in quarter 2:
3. Complete the Atomic Structure/ Periodicity Test Review Form:

Untitled form

2/4 – Tuesday –  period 2/3  
1. Periodic Trend Lab – questions review
2. Start Atomic Structure / Periodicity / Electron configuration Test  – Part 1
                                period 4             
1. Start Atomic Structure / Periodicity / Electron configuration Test – Part 1
2/4 – Tuesday Homework: 
Study for Atomic Structure/Periodicity  Test – 

2/5- Wednesday –  period 2 
1. Complete Atomic Structure / Periodicity / Electron configuration Test – Part 2
                                        period 2/3
1. Complete Atomic Structure / Periodicity / Electron configuration Test – Part 2
2. Metals vs. Nonmetals using the Periodic trends 3-D graph.
– marked a periodic table with metals, nonmetals, and metalliods
– Defined metals and nonmetals based atomic radii, electronegativity and Ionization Energy
– started ionic bonding
did not get to!
3: Periodic Table development 
            Tom Lehr
            Alkali reactivity
            atom shack
4.  Born- Haber Lattice Energy
5. Ionic Bonding – naming /honeymooners
6.  Covalent Bonding setup..
Lewis DOT Diagrams:
Covalent ditto 1 -Electron Dot Diagrams .pdf
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Covalent ditto 1 – Electron Dot Diagrams key .pdf
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Ionic ditto 1 -Electron Dot Diagrams .pdf
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Ionic ditto 1 -Electron Dot Diagrams KEY.pdf
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Support Materials for Periodic Table development- 

Alkali Reactivity


Lattice Energy – Born Haber Cycle

2/5 – Wednesday Homework – 
1.  Please watch the first Video that reviews both types of bonding and how to complete the lewis dt 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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3. Please complete the 1st row 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 key .pdf
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4. Watch lecture 1 up to 8:05 only.
5. Watch the 2nd lecture entirely and follow along with me using the worksheet below.
We will be completing the backside of todays worksheet and learning the following.
a) sigma bonding
b) pi bonding
c)  sp hybridization
d) lewis dot diagrams
e) polarity
SP hybridization KEY p.pdf
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SP hybridization.pdf
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6. Then read Grodski Notes below!
7. Complete Form below.
Todays Lectures:
1st Lecture:


2nd  lectureWatch up to 8:05 of the this one first:


3rd Lecture: Please watch this video in its entirety.

Hybridization Form 1


For those needing more basic review of covalent lewis structures:
 We did this in AP Biology and in the Summer institute or in Regents Chemistry but this is a refresher.
Chemistry text on bonding (page 18 starts hybridization): if want more resources…
Grodski Notes – (from my textbook!) – 
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.
END of Wednesday….

2/6- Thursday –  period 2/3 , 4

1: Periodic Table development 
            Tom Lehr
            Alkali reactivity
            atom shack
2.  Ionic vs. Covalent bonds properties – review covalent and ionic Lewis Diagrams.
3. hybridization (sp) explanation/ molecular diagrams
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!
– completed the alkali reactivity presentation with the reaction between Cs and F.
period 4 needs to complete – alkali reactivity.
2/6- Thursday – Homework
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 (Wednesday’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 sp2 hybridized 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.
1. make one more submission to last nights form! The second to last row of the last question had an issue and was not graded last night. It has been changed and it will now count.
2.  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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sp Hybridization Lecture: 
Complete the sp2 Hybridization Form:

Hybridization Form 2


End of Thursday..

2/7 – Friday – period 2 –
*Connections – Crystal field theory: (Diagrams posted Monday)
Mn+7  in  KMnO4 = purple (in the oxidation titration lab) ——>  Mn+2  (Clear)
Why are the 2 reasons why Mn+2 is clear and does not undergo crystal field theory?
Why are all Zn ions +2 and why are they colorless as well?
1.  Ionic nomenclature – 

 Binary – 2 different ions

        Li & Cl                   Mg & N                 Cu & O          


        Li+1  Cl-1                 Mg+2  N-3          Cu+1 or +2  O-2                 
          LiCl (s)                 Mg3N (s)                  Cu2O (s)
     lithium chloride        magnesium nitride      copper (I) oxide

 Ternary – more 2 atoms or with a polyatomic ion

     Na+ & OH        Zn+2 PO4-3     Fe+3 or +2 & SO4+2
         NaOH (s)        Zn3(PO4)2 (s)     Fe2(SO4)(s)
      sodium hydroxide         zinc phosphate      iron (IIIsulfate
                          Polyatomic name from Table E!


3.  sp, sp2 ,   and sp  – Began by filling out Hybridization Comparison chart for sp and sp2 only.

        – began lecture of sp3 using sp3 hybridization worksheet
4.  VESPR THeory – PHET
Hybridization Comparison Chart Intro.pdf
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SP3 hybridization.pdf
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                          – period 3/4
period 4 needs to complete – alkali reactivity.
2. Born – Haber Cycle – ionic Bonding/ honeymooners
3.  sp, sp2 ,   and sp3
4.  VESPR THeory – PHET
Octet Hybridization Table:
2/7 – Friday -weekend homework –   (4 Parts )
1: Watch Lecture on sp3 and follow along with me with the SP3 hybridization.pdf worksheet.
     Use the octet hybridization family posted above as needed.
SP3 hybridization.pdf
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sp3 hybridization key p.pdf
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The sp3 lecture:  
2:  Please complete the Hybridization Comparison Chart Intro.pdf with me.  You may skip the first 2 rows on the first page as we did them in class today. Do the worksheet up till 24:33 in the video!
Use the octet hybridization family posted above as needed.
Period 2  and  Period 3/4  – watch up till 24:33 – 
Hybridization Comparison Chart Intro.pdf
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Video for Hybridization Comparison Chart Intro.pdfHomework:
3:  Please complete the Ethane, Ethene, Ethyne Comparison.pdf with me with the video with me!.
      Use the octet hybridization family posted above as needed.
Both classes!!
Ethane, Ethene, Ethyne Comparison.pdf
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Ethane, Ethene, Ethyne Comparison worksheet key p.pdf
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Lecture on the  Ethane, Ethene, Ethyne Comparison.pdf worksheet:
Sorry I forgot To hand this out.
Also in the last few minutes of this lecture I review the bonding in Ibuprofen and Aspirin.
You do not have that worksheet yet.
4. Complete the Hybridization 3 form below:

Hybridization Form 3

Optional resource materials:
Hybridization worksheets that we worked on this week:
 s  and 1 p orbital “mixed”
(2 unhybridized p left)
triple bonds possible = 2 
π, 1 𝞼
 SP hybridization.pdf
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 SP hybridization KEY p.pdf
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 s and 2 p orbitals “mixed”
        (1 unhybridized p left)
double bonds possible = 1 
π , 1 𝞼
 sP2 hybridization.pdf
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 SP2 hybridization KEY p.pdf
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 s and 3 p orbitals “mixed”
 (0 unhybridized p left)
single bonds only = 0 
π , 1 𝞼
 SP3 hybridization.pdf
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 sp3 hybridization key p.pdf
View Download


Optional review video:
Hybridization review sp, sp2 and sp3:
Optional review of Polar and Nonpolar molecules.
Polarity of Molecules:
PHET Animation: VSEPR = Valence Shell Electron Pair Repulsion Theory


End of week 2!