1. ELECTRON CONFIGURATION Why are ions more stable than some neutral atoms?

2. WHY DO ELEMENTS FORM IONS?  The reason that elements form ions is because they want to be stable  REMEMBER: The noble gases are nonreactive (inert)  This is because they are stable  Other elements want to be stable like the noble gases  This is why they gain or lose e-

3. ELECTRON CONFIGURATION  For many years, chemists didn’t understand why chemicals released certain color spectra (remember the flame test lab)  Niels Bohr was the first to assume light emission was due to electrons moving to different energy states

4. ELECTRON CONFIGURATION  If electrons gain energy (for example by heating) they can move from a low energy state ( ground state ) to a high energy state ( excited state )  This high energy state is unstable and the electron drops down to the more stable energy state by releasing energy as light (different color spectra)  SEE board for model

5. ELECTRON CONFIGURATION  Niels Bohr was able to predict the color spectra released from only Hydrogen  Problem was he assumed the electrons orbited the nucleus like planets orbit the sun (Bohr model example on board)  He was wrong, the model that explained emission spectra was the QUANTUM-MECHANICAL MODEL

6. ELECTRON CONFIGURATION  In this model, the orbits of electrons are explained using ORBITALS  Orbitals: a region of an atom in which there is a high probablility of finding one or more electrons  In other words, we don’t know exactly where atoms are. We only know where they are likely to be

7. ELECTRON CONFIGURATION  Example of the electron orbital (cloud) for hydrogen (1s 1 ):

8. ELECTRON ORBITALS  To explain orbitals, you need three descriptors:  Principle quantum number (n)  This is the energy level of the electron. The bigger the number, the more energy.  Letter designating the shape of the orbital  s (can hold 2 electrons)  p (can hold 6 electrons)  d (can hold 10 electrons)  f (can hold 14 electrons)  The number of electrons in each orbital

9. EXAMPLE  1s 2 : What does this mean?  1 : This means that the orbital is in the principal quantum number of 1. This is the lowest energy level.  s : This means the orbital has the s shape (spherical)  2 : This means the orbital contains 2 electrons (e-)

10. ELECTRON ORBITALS  The first principal quantum level (n=1) contains only one orbital (1s)  The second principal quantum level (n=2) contains two orbitals  2s  Spherical like the 1s, but larger and with more energy  Holds a total of 2 electrons (e-)  2p  Has 3 subshells with the same shape, but different orientation  Holds a total of 6 electrons (e-), 2 in each orbital

11. (P) ORBITAL

12. ELECTRON ORBITALS  The next principal quantum level (n=3) holds 3 orbitals  3s  Same shape as 1s and 2s, but larger and more energy  Holds 2 e-  3p  Same shape as 2p, but larger and more energy  Holds 6 e-  3d  Has 5 subshells that each hold 2 e- for a total of 10 e-

13. (D) ORBITAL

14. ELECTRON ORBITALS  The next principal quantum level (n=4) holds 4 orbitals  4s: Same shape as 1s,2s and 3s, but larger and more energy  4p: Same shape as 2p and 3p, but larger and more energy  4d: Same shape as 3d, but larger and more energy  4f  Has 7 subshells  Holds a total of 14 e-, 2 in each orbital

15. (F) ORBITAL

16. HOW DO WE PUT ALL OF THIS TOGETHER  Each atom has a certain number of electrons  Example: Hydrogen has 1 e-  Therefore the electron configuration for H is 1s 1  This means hydrogen has 1 electron in the principal quantum number (n=1) in the s orbital

17. HOW DO WE PUT ALL OF THIS TOGETHER  Example: Helium has 2 e-  Therefore the electron configuration for He is 1s 2  This means helium has 2 electrons in the principal quantum number (n=1) in the s orbital

18. ELECTRON SPIN  Electron spin is a property of all e-  Electron spin is explained by the Pauli Exlusion Principle  A maximum of two electrons can occupy each subshell, and these electrons must have different spin quantum numbers  What this means: You can only have 0,1,or 2 electrons in any given orbital

19. HOW DO WE PUT ALL OF THIS TOGETHER  In order to figure out electron configuration, you fill in each orbital from the lowest energy level to the highest energy level  That’s why H is always 1s 1. H has the fewest number of e-.  He is 1s 2 because He has 2 e- and it must completely fill the s orbital before going to a higher energy level

20. ORDER OF ORBITAL ENERGY LEVEL  Arrangement of orbitals in order of increasing energy  1s  2s  2p 2p 2p  3s  3p 3p 3p  4s  3d 3d 3d 3d 3d  4p 4p 4p  5s  4d 4d 4d 4d 4d, etc.  NOTE : Energy levels do not increase in numerical order

21. EXAMPLE  What is the orbital notation for potassium?  Sodium (Na) has 11 electrons  Fill electrons in order of lowest to highest energy levels  1s (2 electrons)  2s (2 electrons)  2p (6 electrons)  3s (1 electron)  The notation for Na = 1s 2 2s 2 2p 6 3s 1

22. THERE MUST BE AN EASIER WAY  How do we figure out electron configuration without using that weird diagram?  THE PERIODIC TABLE  Let’s take a look at it

23. TRY THESE  Magnesium (Mg)  Argon (Ar)  Oxygen (O)

24. ANSWERS 1. Mg = 1s 2 2s 2 2p 6 3s 2 2. Ar = 1s 2 2s 2 2p 6 3s 2 3p 6 3. O = 1s 2 2s 2 2p 4

25. ORBITAL DIAGRAMS  Orbital diagrams are another way to show how electrons fill orbitals  Example: Hydrogen (H) and Helium (He) H  the arrow indicates e- spin 1s He  opposite arrows show 1s opposite spin

26. ORBITAL DIAGRAMS  Neon (Ne, 10 e-): Ne 1s 2s 2p 2p 2p

27. HUNDS’ RULE  Hund’s Rule : The most stable arrangement of electrons is that with the maximum number of unpaired electrons, all with the same spin  In other words, you fill each orbital with 1 e- before starting to add the second

28. EXAMPLE  Carbon (C, 6e-) C 1s 2s 2p 2p 2p  Each 2p gets an e- before you begin to fill the orbital with the other e-

29. EXAMPLE  Oxygen (O, 8e-) O 1s 2s 2p 2p 2p  Once you fill all the 2p with 1 e-, you go back to fill the rest

30. TRY THESE 1. Nitrogen 2. Phosphorus 3. Sulfur

31. SHORT HAND  The noble gases fill all of their orbital shells  Therefore we can abbreviate elements based on the previous noble gas  The outermost electrons available are called valence electrons  VALENCE ELECTRONS : the electrons in an atom’s outermost orbitals; determines the chemical properties of the element

32. SHORT HAND (CONT)  Neon (Ne, 10 e-): Ne 1s 2s 2p 2p 2p  Therefore we can abbreviate Sodium with 11e- as (Na) = [Ne]3s 1

33. SHORT HAND (CONT)  Calcium (Ca, 20e-)  [Ar]4s 2  Silicon (Si, 14e-)  [Ne]3s 2 3p 2

34. TRY THESE 1. Iron (Fe, 26 electrons) 2. Rubidium (Rb, 37 electrons) 3. Chlorine (Cl, 17 electrons)

35. WHAT ABOUT IONS?  The purpose of all electron orbitals is to be FULL  If an orbital is not full, then it wants to either get rid of an electron or grab an electron  This depends on what is energetically favorable

36. FOR EXAMPLE  Draw the orbital diagram for sodium:

37. EXAMPLE  Therefore the orbital diagram for Na: Na 1s 2s 2p 2p 2p 3s  Notice: the 3s orbital is not full. What do you think happens?

38. FINAL ANSWER  Sodium will lose 1e- so it will look like the noble gas neon  Forms Na +  Draw what the orbital diagram and electron configuration will be for Na +

39. TRY THESE  Write out the orbital diagram and electron configuration for the following ions: 1. Mg +2 2. F - 3. Al +3 4. Cs + 5. Cu +2

40. HOW THE ELEMENTS COMBINE TO FORM COMPOUNDS  To have full orbital shells, some elements must give up (donate) and electron and some must receive (accept) an electron  Therefore elements combine to share electrons  Therefore they: FORM IONS

41. SUMMARY  Electrons fill orbitals from lowest e- energy level first  No more than 2 e- per each orbital (must have opposite spins) – Pauli Exclusion Principle  When orbitals of identical energy are available, e- occupy orbitals singly before pairing up (Hund’s rule)