1. Understanding Matter Part II Beyond the Bohr model

2. ELECTRON CONFIGUIRATION SHELL DIAGRAM

3. Quantum Mechanical Model  As we saw earlier, the Bohr Model had several short comings  Krypton does not follow the 2, 8, 8 pattern.  In order for Krypton to have enough electrons for 36 it needs an extra 18 electrons.  The model currently used to describe the atom is the Quantum Mechanical Model of the atom  This is the current theoretical framework that is used to describe all of the information we have about atoms and how they function

4. Definitions  Quantum (plural ‘quanta’)  A finite amount of energy  i.e. – an energy level in an atom  The amount of energy required to move an electron from its present energy level to the next higher one  Electrons can only have specific energy levels and nothing inbetween.  Mechanical  Movement of parts in relation to a whole  i.e. – electrons in an atom  Hence the Quantum Mechanical Model deals with the movement and location of electrons in an atom

5. Uncertainty Principle  We cannot know where an electron is and where it is going  Because of this, we use probability to determine where an electron is most likely to be  Using the electron probabilities, we find areas where electrons are most likely to be  These areas are called electron clouds where the probabilities of finding electrons is very high  The shapes and distance from the nucleus of these electron clouds depends on several factors

6. Quantum Numbers  To describe electron clouds and where electrons probably are, we use quantum numbers  There are a total of four (4) quantum  Principal Quantum Number  Angular Quantum Number  Magnetic Quantum Number  Spin Quantum Number We will be concerned with theses 2

7. Principal Quantum Number  Energy level  Distance away from the nucleus  As # increases, distance from the nucleus also increases  As the number increases, so does the energy of the electrons in those orbitals  Represented by integers 1,2,3,4,5,6,7 that correspond to the seven horizontal rows on the periodic table  Determined by counting as you move down (top to bottom) the periodic table

8. Angular Quantum Number  Also known as “sub-shells”  Refer to the shape of the orbital  There are four (4) different shapes  S, P, D, F  These correspond to the s, p, d, f blocks on the periodic table

9. Periodic table shows Quantum Structure Energy increases as you go down the periods Subshell s Subshell p Subshell d Subshell f

10. Sub-Shells  S” Sub-shell  Spherical shape Only one (1) orbital per energy level  The 1 sub shell can hold 2 electrons  “P” Sub-shell  Dumbbell shape  Three (3) orbitals per energy level  Each shell can hold 2 electrons 3 orbitals mean the p-shell can hold up to 6 electrons

11. Sub-Shells Continued  “D” Sub-shell  Tend to have a clover-leaf shape  Five (5) orbitals per energy level  Each can hold a maximum of two (2) electrons  Can hold a max of 10 electrons  “F” Sub-shell  Shape contains 6 lobes for the most part  Seven (7) orbitals per energy level  Each can hold a maximum of two (2) electrons  Fourteen 14 electrons total at each energy level

12. To Summarize

13. Modeling the Quantum Atom  Krypton

14. Potassium

16. Manganese

17. Three principles for electrons filling Shells  Aufbau Principle:  Electrons enter sub-shells of lowest energy first  1 st energy level fills up before the next  Pauli Exclusion Principle:  All atomic sub-shells contain a maximum of two (2) electrons. Each MUST have a different spin  Hund’s Rule:  when electrons occupy sub-shells of equal energy, ONE electron enters EACH sub-shell until all the sub-shells contain one electron with identical directions  Electrons are added to sub-shells so that a maximum number of unpaired electrons result

18. ELECTRON CONFIGUIRATION ORDER OF FILLING ORBITALS  Orbitals are filled in increasing order of energy  Different blocks on the periodic table (shaded in different colors in this chart) correspond to different types of orbitals.  The periodic table, from left to right, shows the ‘basic’ pattern of sub- shell filling.

19. Cheat Note

20. Orbital Notation and Electron Configuration ORDER OF FILLING ORBITALS - HUND’S RULE  The lowest energy stability of atom is attained when the number of electrons with the same spin is maximized

21. Lets Try  Oxygen

22. Oxygen

23. Lets Try  Aluminum

24. Aluminum

25. Lets Try  Chlorine

26. Chlorine

27. Nobel Gas notation  An even more simplified and shorthand method for representing electron configuration.  Emphasizes the outermost energy level only  Instead of listing every energy level and amount of electrons individually, it utilizes the nearest noble gas element of the energy level below as a representation of the inner energy levels

28. Nobel Gas Notation Example  For Example: Sulfur  Electron configuration would be:  1s 2 2s 2 2p 6 3s 2 3p 4  Its Noble Gas Notation would be:  [Ne] 3s 2 3p 4  …this is because we know that the electron configuration of Ne is: 1s 2 2s 2 2p 6, therefore there is no need to write it all out.

29. Nobel Gas Examples

30. Examples Continued

31. Nobel Gas Notation  Noble Gas Configurations are especially useful for elements with a large atomic number, as their complete electron configurations become tiresome & redundant to write out each time.

32. You Try  Complete Orbital and noble gas practice sheet.  For additional practice you could try:  Write out the orbital diagram and electron configuration of all even number elements up to 18  Write out the Nobel gas notation for each of the following elements.  Ca, Br, Nd, U, Co and Au