Evidence for shells

History of the atom

 

·        The model of the atom has changed as our observations of its behaviour and properties have increased.

·        A model is used to explain observations.  The model changes to explain any new observations.

·        The stages in the development of the atom:-

 

George Johnstone Stoney (1891)                   Electrolysis.  The charge of an electron.

 

Joseph J Thompson (1897)                             The cathode ray tube and e/m deflection.  The mass / charge of an electron.

 

Robert Milikan (1909)                                         Oil drop experiment.  The mass / charge of an electron.

 

Joseph J Thompson                                          ‘Plum – Pudding’ model of an atom.

 

Geiger, Rutherford and Marsden (1909)        Alpha particle deflection.  The nuclear model.

 

Henry Moseley (1913)                                        Atomic number

 

Plasma displays

 

  • A plasma unit has 100,000's of tiny cells (pixels) filled with a mixture of neon and xenon gasses.
  • A single pixel is made up of three coloured sub-pixels, red, green and blue.
  • Each sub pixel is driven by its own electrode, stimulates the gas to release ultraviolet light photons.
  • The photons interact with a phosphor material coated on the inside wall of the cell.
  • Phosphors are substances that give off light when they are exposed to other light eg. Ultraviolet light. The phosphors in a pixel give off coloured light when they are charged.
  • The cells are situated in a grid like structure.
  • The varying intensity of the current can create millions of different combinations of red, green and blue across the entire spectrum of colour.

 

Ne(g) à Ne+(g) + e-    
Xe(g) à Xe+(g) + e-    

 

Ionisation energies:        Click for DEMO (click demo once)

·        To form positive ions, electrons must be completely removed i.e. ionisation.

·        To do this the electron must completely escape the attraction of the atom. i.e. reach n=¥.

·        At n=¥ the electron has sufficient energy to escape the attraction from the nucleus.

 

 

   Definition:    

·     The 1st ionisation energy of an element is the energy required to remove 1 electrons from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions

1st ionisation energy:

                                    Ca(g)             à        Ca+(g)        +          e-                        1st IE = +590 KjMol-1

2nd ionisation energy:

                                    Ca+(g)           à        Ca2+(g)       +          e-                        2nd IE = +1150 KjMol-1

Factors affecting ionisation energy - always refer to all 3 in any explainations:

 

1)     The distance of the electron from the nucleus

                    F  a  1/d2

2)     Size of the positive nuclear charge

3)     The ‘shielding’ effect by full inner shells

Ionisation energies down a Group:

Ionisation energies across a Period:

Successive ionisation energies        Click for DEMO

·         The successive ionisation energies can tell us which group an element is in.

·         For potassium the easiest electron(s) to remove before a large jump tells us the group.

·         This can also be done by looking at data. 

 

        Click for SUMMARY DEMO

 

IE's & shells

 

Questions: 1 - 3 p41/ 13 p73

 

Shells and orbitals

 

Flame colours and emission spectra

 

Practical

 

See some line emission spectra

 

Energy levels or shells - Explanation

                1)  That the lines are specific colours representing specific energies / electron shells.

 

                2)  That as the move up in energy (towards the violet end) they get closer together (converge)

 

Shells to levels

 

Atomic orbitals

 

Orbit problem

                 

s - Orbitals:

                   

p - Orbitals:

à

see sp orbitals

d and f orbitals:

A summary of orbitals:

Orbital theory

Complete the table below:  Answer

Electron shell Principle Quantum Number Types of orbital No electrons in s shell No electrons in p shell No electrons in d shell No electrons in f shell Total
1              
2              
3              
4              

 

Representing electrons in orbitals

Text Box:

 

 

 

 

 

 

Orbitals

 

Questions 1 - 2 p43

 

Sub - shells and energy levels

 

Sub - shells

n = 1 shell: maximum 2 electrons

Sub - shell 1s    
Orbital    
Electrons 2e    
       

n = 2 shell: maximum 8 electrons

Sub - shell 2s 2p  
Orbital  
Electrons 2e

2e

2e

2e

 
       

n = 3 shell: maximum 18 electrons

Sub - shell 3s 3p 3d
Orbital
Electrons 2e

2e

2e

2e

    2e

2e

2e

2e

  2e

 

 

Electrons and the Periodic Table

Need for a more complex model:

a) 2 - 3 split:  s and p orbitals

 

                       

 

b) 5 - 6 split:  filling the p orbital

Electron shells overlap

  1. The larger the value of n, the higher the energy level and the further from the nucleus.

  2. That the electron shells get closer together as you move away from the nucleus.

  3. As we move from n = 1 to n = 4, a new type of sub shell is added.

  4. Within a shell the energies of the sub - shells increases in the order s, p, d and f.

  • The 4s energy level is below the 3d energy level.

  • This means that the 4s orbital fills before the 3d orbitals. (according to the Aufbau Principle)

     

 

 

Electron energy levels

 

Filling shells and sub - shells:

The Aufbau Principle:

  1. Electrons are added one at a time to 'build up' the atom.

  2. The lowest available energy level fills first.

  3. Each energy level must be full before the next, higher energy level can be filled.

  4. Each orbital in a sub - shell is filled by single electrons before pairing up.

  5. Each orbital can hold 2e of opposite spin

Filling the orbitals - Using the Aufbau Principle:

 

 

Electron configuration

Element Orbitals occupied Electron configuration
B 1s22s22px1 1s22s22p1
C 1s22s22px12py1 1s22s22p2
N 1s22s22px12py12pz1 1s22s22p3
O 1s22s22px22py12pz1 1s22s22p4

 

Questions 1 - 2 p45

 

Elements filling

 

Test yourself

 

Sub shells and the Periodic Table:

Element No electrons Electron configuration What Period is this element in? What is its highest Principle Quantum Number What is the highest sub - shell What Groups is the element in
H            
He            
Li            
Be            
B            
C            
N            
O            
F            
Ne            
Na            
Mg            
Al            

 

Things to notice:

Extending further:

 

 

Using the Periodic Table for electron configurations

 

 

Other examples:-

 

1)     Na   1s2 2s2 2p6 3s1

 

2)     Sc  1s2 2s2 2p6 3s2 3p6 3d1 4s2  (remember 4s fills before 3d)

 

Shortening an electron configuration

                    Sc          1s2 2s2 2p6 3s2 3p6 3d1 4s2

                    Sc          [1s2 2s2 2p6 3s2 3p6]3d1 4s2

                                           Inner shells

                                    [Ar]  has the same electron configuration as the inner shells

                                    [Ar] 3d1 4s  where [Ar] represents the electronic configuration of argon.

 

Electronic configurations in ions

 

·        Follow the same principle as for atoms but add / remove electrons:-

 

Eg Sodium ion Na+

 

                                                         Na atom =   1s2 2s2 2p6 3s1

 

To make a sodium ion 1 electron is removed

 

                                                         Na ion =   1s2 2s2 2p6

 

Eg Chlorine ion Cl-

 

                                                         Cl atom =   1s2 2s2 2p6 3s22p5

 

To make a chlorine ion 1 electron is added

 

                                                         Cl- ion =   1s2 2s2 2p6 3s22p6

Things to note:

Questions  1 - 4 p47 / 1,2 and 13 p73 / 1 p74