Q3 : Week 2 – 2/6 – 2/10

<- Week 1   –   Week 3 ->                                                                                                                             Jump to:  Tuesday,   Wednesday,  Thursday,   Friday 

Refresh this page every time you arrive!

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

Please Make sure Your read all instructions for all assignments.   

Make Sure Powerschool grades are accurate!!     

                                                                                                                                                                                                                                                         Jump to: Monday Homework  

^{_____________________________________________________}                                                                                                                                                              2/6 – Monday – B Day – 2, 3b Lab/4      

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

                                                 m_s =  individual orbitals (actual 3 -d shape where pairs of electrons can exist)

 

Every Electron has a unique set of 4 quantum numbers (including spin m_s) = Pauli 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 = -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, m_l = 0)

p = 2 orbitals (l = 1, m_l = -1,  0, +1)

d = 5 orbitals (l = 2, m_l = -2, -1, 0, +1, +2)

f  = 7 orbitals (l  = 3, m_l = -3, -2, -1, 0, +1, +2, +3)

g =  ?

 

Electron configuration of Al:           Nucleus:  1s²2s²2p⁶3s²3p¹

 

Orbital notation for Al:

Each box represents a single orbital.

 

Which electrons are the most stable? Unstable?

 

*Connections – Our quantum numbers refer to specific energy levels allowable by each quantized element.  The arrangement of electrons in quantized energy orbitals that were solutions to the Erwin Schrodinger equation (quantum numbers) are actually hidden in the periodic table that was already arranged according these energy levels (unbeknownst to Mendeleev and Moseley!!) The periodic table that we use is a condensed version that does not insert the  f block ( l = 3) because it would not fit on most pages!!!!

 

_____________________________________________________                                                                                                                     jump to:  top

2/6 – Monday’s Homework: – 

_______________________________________________________________________                                  Jump to: Tuesday Homework / top

2/7 – Tuesday – A Day – 2/3a Lab, 4   

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

                                                 m_s =  individual orbitals (actual 3 -d shape where pairs of electrons can exist)

 

Every Electron has a unique set of 4 quantum numbers (including spin m_s) = Pauli 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 = -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, m_l = 0)

p = 2 orbitals (l = 1, m_l = -1,  0, +1)

d = 5 orbitals (l = 2, m_l = -2, -1, 0, +1, +2)

f  = 7 orbitals (l  = 3, m_l = -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.

 

*Connections – Our quantum numbers refer to specific energy levels allowable by each quantized element.  The arrangement of electrons in quantized energy orbitals that were solutions to the Erwin Schrodinger equation (quantum numbers) are actually hidden in the periodic table that was already arranged according these energy levels (unbeknownst to Mendeleev and Moseley!!) The periodic table that we use is a condensed version that does not insert the  f block ( l = 3) because it would not fit on most pages!!!!

 

_____________________________________________________                                                                                                                       jump to:  top

2/7 – Tuesday’s Homework: – 

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

__________________________________________________________________________                       Jump to: Wednesday Homework / top

2/8 – Wednesday – B Day – 2, 3b/4 Lab

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!)

 

Z_eff = 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 Z, Z_{eff, and n you can explain almost everything in periodicity and electron}  

                                                                                                                                           _{configurations.}

 

_{****Since we are continuously evaluating the energy levels of electrons that are bound in 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 Z_{eff                               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) :   1s²2s²2p⁶3s¹                                             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

                           

                            IE₁  =   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!

                                                                                                                              

_____________________________________________________.                                                                                                 jump to:  top

2/8 – Wednesday Homework: –

________________________________________________________________________                              Jump to: Thursday Homework / top

2/9 – Thursday – A Day – 2/3a Lab, 4 

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.

 

         

___________________________________________________                                                                                                   jump to:  top

2/9 – Thursday Homework: –

_________________________________________________________________                                               Jump to: Friday Homework / top

2/10 – Friday – B Day – 2, 3b/4 Lab

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 KMnO₄ solution.

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.

____________________________________________                                                                                                                     jump to:  top

2/10 – Friday Homework: –

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