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SUMMER INSTITUTE – Module 9 – Periodic Table and Bonding

______________________________________________________________________________________________________________________________________                                       Jump to Activity 1,  Activity 2,  Activity 3,  Activity 4,  Activity 5,  Activity 6,  Activity 7,  Module Test

Module 9 – Periodic Table and Bonding – 

5 hours – due August 24th!  

Link to the Module 9 worksheets:
***By the way we are entering the 5th week of the 3rd quarter of Regents Chemistry
Module 9 – Periodic Table Notes.pdf
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Module 9 – Bonding Notes.pdf                             


* Connections The electron arrangement and the transitions of electrons in atom is really what determines the chemical properties of elements and compounds.  In large part is it what chemistry is really about.  We have looked the at the stoichiometry chemical compounds and reactions but none of that would be possible if elements did not bond!  Before we begin our bonding discussion we must first start with a brief look at how the periodic table was made and how the electron arrangement or configurations paralleled the early work of the arrangements of elements in the periodic table.  You see the periodic table was made to help organize the known elements of the day in some logical fashion.  One scientist, Dimitri Mendeleev made a table and said that his table followed a repeating pattern based on increasing atomic mass (it was not known in 1850’s that the proton number was the determining factor in differentiating elements) and changing chemical properties.

This pattern he called Periodic Law and his table was different than all others of the day because he was able to predict the chemical properties of elements elements that were not discovered yet based on missing elements in his table.  Mendeleev did not know what caused the repeating pattern of properties but as you can see from your work in atomic structure, the electron configurations were determined in the same patterns (just 70 or so years later!)  Quantum mechanic’s and electron arrangement that resulted from it were proof of the Periodic Law of Mendeleev that he found in 1860!  
Lets take a look at his early table: 
Now in Mendeleev’s table the columns that we see in the modern periodic table were rows. So if you look at row of Na it will be the 1st column in our current Periodic Table.  
Look at the question mark at atomic mass 68 and 70.  Remember that Mendeleev’s table was created in order of atomic mass and that chemical properties repeated with the atomic mass changes.  In his table, if you were in the same row you had the similar chemical properties thus B (Boron), Al (Aluminum) were similar.  Zn who had a similar chemical properties of Mg (we now know that they both liked to become +2) had a atomic mass of 65 while the next heaviest element known at the time was As (Arsenic), atomic mass 75 and it had chemical properties similar to N (nitrogen) and P (phosphorus) so Mendeleev put As in that row.  So because he jumped to rows with As he predicted there were 2 elements that were not discovered yet and the amazing part is that he could predict what the properties of these elements based on the atomic mass they had.  If is had an atomic mass of about 68 then it would have properties of B and Al but if it had a atomic mass of about 70 it would have properties of C (carbon) and S (silicon). 
Gallium (Ga – 69.7) and Germanium (Ga – 72.6) were soon discovered after the publishing of Mendeleev’s table and they of course had the correct atomic mass and properties that Mendeleev predicted, giving evidence to his periodic law!  I must also report that on many occasions Mendeleev changed the known atomic masses of certain elements (based on the mole concept!!!!) because they were often not very precise values.  Measuring pure gases as the same temperature and pressure was not easy work, especially if some elements were not gases at STP. Not only did he predict the chemical properties of elements not discovered yet he also corrected many atomic masses  that were accepted values but were not accurate.
You may also ask what were the similar chemical properties that Mendeleev was using as his chemical law?  They were the chemical compounds that he could write because of Cannizzaro’s work at the Chemistry Conference!!!   For example: B and Al are related because both will make the same chemical formula with oxygen:
                             B2O3  ,  Al2O3  and guess how Gallium binds with oxygen?    Ga2O 
                      That means that Boron, Aluminum, and Gallium all become +3
                       which means all three have the same number of valence electrons!
Why is this so?  Mendeleev just said it was Periodic Law but we now know because B, Al and Ga both have the same number of valence electrons due to having similar electron configurations (which again was not found for 70 years later with french accent!)
Also we learned about Robert Moseley who used x-rays to determine the unique proton number of each atom.  After his work the periodic table had to be adjusted so that the Periodic table went in order of atomic number NOT atomic mass.  For instance: 
Notice that Co (58.9332) goes in front of Ni (58.693) even though Co has a higher atomic mass!! These 2 elements had to be switched, after Moseley’s work because the periodic table was actually a function of the atomic number NOT ATOMIC MASS that Mendeleev used.
So you can see that all of these historical developments are interrelated!!



 Activity 1:  SKILL – Identify and predict the 3 periodic trends, Electronegativity, Ionization                                   Energy, and Atomic Radius.

1: Table of elements arranged according to chemical properties (that we now know to be valence electrons).
2:  The three trends of the periodic table:
    a) Atomic Radii–  The size of the atom
    b) Electronegativity–  The attraction that an element has for (other atoms) electrons.
    c) Ionization Energy – The energy needed to pull electrons away from an atom.
3:  These three trends are responsible for all of the properties of elements and thus defines elements as Metals, Non- metals and Metalloids.


Activity 1 Assignment:

1.  Watch Lecture 3.10R – Periodic Table properties with the BlankPeriodicTable[1].pdf worksheet.   
2:  Fill in the worksheet as directed in the lecture .    
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    1 : Lecture 3.10R

Bonus optional classroom lecture Using “Elvira the outermost electron” teaching the three periodic trends.  



Activity 2:  SKILL : Metals/Nonmetals/Metalloids periodic table geography table and the                                                   reactivities .

Concepts on this lecture:
a) Metalliods (Semimetals) separate the metals and nonmetals.
b) Metalliods are located on the Staircase in the periodic table
c) TRansitional metals are locate in the middle of the periodic table and are less reactive than the metals          that are in group to the left.
d) Metal reactivity increases as the atomic radius increases.
e) Nonmetal reactivity increases as the atomic radius decreases.
g) Most reactive metals are in the Alkali (group or column 1) and the Alkaline Earth (group or column 2).
h) Most reactive nonmetals are in group (column) 17 called the Halogens.
i)  The least reactive group is in group (column) 18 called the Noble Gases.
1. Please watch the lecture 3.11R below and continue placing notes on the blank periodic table you used for activity 2.
2.  View the all slides of the Alkali (group 1) metal reactivity presentation.  Basically, I have videos of all metals of the family dropped into water.  Each metal will lose its single valence electron to hydrogen that is bonded to the water.  Notice what happens to the reactivity (the ability to lose electrons) of each metal as we move down the group (or as the atomic radius increases)
There are no questions for the presentation but you need to view the demos to get a feel of what I am trying to teach!



    1 : Lecture 3.11R – Families and Reactivities


2 : Alkali Reactivity –  




Activity 3:  SKILLAtomic Radius vs Ionic Radius, Transitional Metals, Phases

1.  Please watch lecture 3.12R, 3.13R, 3.13R ( THEY ARE SMALLER LECTURES).
2.  Complete the combined 2 periodic table worksheet.pdf and put the answers into the form below.
combined 2 periodic table worksheet.pdf
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    1 : Lecture 3.12R – Atomic and Ionic Radius

    1 : Lecture 3.13R – Transitional Metals

    1 : Lecture 3.14R – Phases of the elements


2: Periodic Table Worksheet Form 1


Activity 4:  SKILL: Observe the periodic trends through Graphs of Data from Elements

1. Please download the three graphs that plot Electronegativity, Ionization Energy, and Atomic Radii with atomic number.
In your regents chemistry class you would have made 3 graphs that plotted the three trends listed above with increasing atomic number to visually see the periodicity of elements that Mendeleev first discovered.
Atomic Radius, Electronegativity, Ionization Energy graphs vs Atomic Number.pdf
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2. Write the element’s symbol above each major peak and valley on of the three graphs like this:

Please label the major peaks and valley’s.  I did not label the  first and second element, but you should.  You can just use a  marker on your graph.  This will make the questions easier to answer.


3.  Please answer the following questions in the form below: 



3 : Periodic Graphs Form: 


Activity 5:  SKILL: Intro to Bonding, IONIC Bonds and a little Covalent bonding.

1: Please carefully read the following:
Connections: At this point we have seen how the arrangement of electrons determines stability in electrons in atoms and we know understand how much an electron(s) attracts nucleus determines whether that atom has the ability to become stable
For instance Na is an unstable atom that badly wants to lose an electron to become Na+ 
              1s22s22p63s1                                      to                                             1s22s22p6
                             by losing 1 electron (its valence electron) from the 3s orbital .
The very important point here is that the atom Na and all alkali metals are very unstable (high energy) because they are 1 electron away from attaining a very stable electron configuration ( 1s22s22p6 = Ne). Ne is a noble gas (inert gas( that does not react because it is already stable!!! Atoms that are reactive bond to get stable!!!!
Remember, the reason that Na becomes Na+ is to attain a noble gas electron configuration. BUT THEY CAN ONLY GET STABLE IF THEY HAVE THE ABILITY TO LOSE AN ELECTRON!  You see the large atomic radii gives the Na the pathway to get stable and thus react. Everything is interconnected!!!
Think of boulder high up on a mountain.  It is unstable and has the potential to roll down the mountain with tremendous force like the atom Na when it reacts (or rubidium in the Alkali presentation you watched in this module).  This boulder can only “react” when it is near the edge of the mountain. Being at the edge gives the boulder the opportunity to roll or react, just as having a large atomic radius gives Na the ability to react.  It is not just stability that drives reactions but a pathway! (We will learn that a pathway is called Entropy!
So a metal like Na loses an electron to become stable AND A NONMETAL will gain an electron to become stable! Let’s look at the nonmetal Chlorine.
                        A Chlorine atom Cl , a nonmetal is also very unstable and it badly wants to become Cl
                               1s22s22p63s23p5                               to                                   1s22s22p63s23p6 
                                      by gaining 1 electron (to fill its valence shell) in the 3p orbital .
                **Notice that Cl attains an electron configuration of Ne: 1s22s22p63s23p6   by gaining electrons!
 So if Na gets stable by losing an electron and Cl gets stable by gaining an electron then if you put these 2 atoms together they will react by transferring an electron (redox reaction) from the Na to the Cl.  This results in a ionic bond!
        Lets look at an animation: 

Notice that the Na gets smaller when it loses electrons and the Cl gets bigger when it gains!

The transfer of electrons causes each Atom to become charged (ION) and the opposite charged ions electrostatically attract each other into an IONic bond.  BECAUSE EACH ION is charged in each direction, it can form an ionic  bond in many directions creating a crystal lattice which is the most stable   structure.

Notice the crystal has no net charge because the formula has a 1 Na to 1 Cl (lowest ratio)!! 

      The Formula is NaCl and we call it sodium chloride   Connections!


Now let’s take a look at the real chemical reaction between Na and Cl.  Remember that Cl is diatomic (HOFBRINCl’s) so the reaction is really between Na and Cl2.
                           2Na (s)  +   Cl2 (g)        2NaCl (s)  +  Heat 

This is a very exothermic reaction due to the great amount of heat that is released overall.

The net change of heat released means that there is more energy released in forming the stable ionic solid (NaCl) than energy needed to break the initial bonds in Cl2.


WHY?  BONDING, like in NaCl results in stable electron configurations thus to break a stable (low energy state) energy is required, while energy is released when a bond (a low energy state) is made.
So in chemical reaction bonds must be broken (endothermic) and rearranged or formed (exothermic).  How is it that the reaction of making NaCl was overall exothermic (the heat was written in the product side). The answer is that it more energy was released in forming the stable NaCl ionic crystal than it was in breaking the bond in chlorine ( Cl2  ).

                           The enthalpy or  Heat of reaction when NaCl is formed from its elements is

                     ΔHf (kJ/mol) = 411 , meaning 411 Kilojoules of heat are released (negative sign)                                                       for every  mole of NaCl formed.
                    Notice if we did the reverse reaction we would have a positive ΔH because the reverse reaction is endothermic and we would need more energy to break the stable bonds and crystal lattice formation of the ionic compound.
So what we just did was view how IONIC BONDS are formed and they must be formed with elements that have Low Electronegativity with ones that have HIGH Electronegativity.
                                                                                       easily lose electrons                                                                                                         easily gain electrons
 Ionic Bonds:                                Metals                             &                                Nonmetals
You can think of an ionic bond as one where the two atoms are in a tug of war and because in the case of Ionic Bonds, one of the atoms is very weak (metal) holding onto the rope (valence electrons) and the other is very strong (nonmetal).  In this scenario the nonmetal will pull the rope away from the metal and a transfer of electrons occur.  This occurs because the difference in electronegativity is very large.
Covalent bonds:  as you will learn covalent bonds occur between nonmetals and nonmetals. In this scenario, since both atoms are nonmetals and they have both high electronegativity the rope must be shared. This is due to a small difference in electronegativity between bonding atoms.  Notice the name Covalent! co sharers of the valence electrons.  There is no transfer of electrons in this type of bonding. This is the type of bonding that occurred in your AP Biology course
Here is a comparison:
                            Bonds between ions!                                    Bonds between nonmetals!
                                  Metal and Nonmetal                                               Nonmetal and Nonmetals
                       A electron transfer makes ions                                                    Forced to share!
             Notice Brackets and Charges to show no sharing!                              No Brackets, they are sharing electrons!
                           THESE BONDS CREATE CRYSTALS                                         These bonds create mostly MOLECULES!
                  In lowest Ratio (Empirical): Na2O                           In actual ratio that exists (molecular): H2O2   
                               High Melting points                                                                Low melting points
          Ionic compounds are in very stable crystals                      Molecular compounds attract each weaker IMF’s
                         Na2O   :  1182 degrees Celsius                                              H2O2    :   -0.43 degrees Celsius
     Conduct electricity as a liquid or solution only                              Never Conducts electricity! (insulator)
         In a liquid or solution the IONs ARE FREE!                               Not made of ions nor are the electrons free to
                                                                                                                  move because of high electronegativity
2.  Please watch the following on the properties of these compounds:

3. Please view the 2 lectures below on writing Lewis Dot diagrams for Ionic compounds and complete the worksheet by reviewing with the key.
Ionic ditto 1 -Electron Dot Diagrams 1112.pdf
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Ionic ditto 1 -Electron Dot Diagrams 1112 key.pdf
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Forget the notebook file I tell you to download.  Just listen and view how I write Lewis dot diagrams for Ionic compounds.   

 Lecture 2.17 – Lewis Diagrams of Ionic Compounds.

                                        Simple review of ionic lewis diagrams: (utilizing the criss-cross method)
                                        We did this sort of did this with writing chemical formulas. 



Activity 6: SKILL: Covalent Bonds and Lewis Dot Diagrams

                                   and determining Bond and Molecular Polarity
1.  Please view the lecture 4.16R and 4.14R (in that order)below and fill out the backside of the worksheet with me. You can keep reviewing with me question by question or you can stop video and go ahead and complete. 
2. Complete the worksheet front and back and review with the key.
Covalent ditto 1 -Electron Dot Diagrams 1112.pdf
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Covalent ditto 1 – Electron Dot Diagrams key 1112.pdf
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    1 : Lecture 4.16R – Polar Bonds vs Polar molecules

    1 : Lecture 4.14R – Covalent Lewis Structures and polarity

     Optional Mini Lectures – If you need more clarity.

                      Polar and Nonpolar Bonds :                                                      Polar and Nonpolar Molecules:



Activity 7:  Skill –  Metallic Bonds, identification of types of compounds and properties

1. Please view the short lecture below on metallic bonds ( Its not mine but it is very       
     good!). Mine are not that short! I am not good at doing short!
2. Complete with me the worksheet below using the lecture 2.23 and then complete with key
Bodacious bonding with properties chart.pdf
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Bodacoius bonding with properties chart key.pdf
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    1 : Metallic Bonding Lecture

Metallic Bond lecture:


    2 : Lecture 2.23 – Metallic Bonding and Properties


Module Test 9Periodic Table Assessment and Bonding Assessment   

– available below

Once you are done with all activities please complete Module Test 9 – by using the files below and the form below:
Module 9 Test – Periodic Table.pdf
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Module Test 9 – Bonding.pdf
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Remember none of this is graded but I am getting a number to see how well you understand the topics.  If you are enjoying this madness then you are at the right place.  If this is very painful, I am sorry but you need these skills for AP Chemistry!!!
 You must have completed all assignments above first before you begin your module Test 9.



Module Test 9 – Periodic Table

Module Test 9 – Bonding –