3. Bohr Model of the Atom  Bohr’s Atomic Model of Hydrogen  Bohr - electrons exist in energy levels AND defined orbits around the nucleus.  Each orbit corresponds to a different energy level.  The further out the orbit, the higher the energy level

4. Other Scientists Contributions  De Broglie  Heisenburg  Modeled electrons as waves  Heisenberg Uncertainty Principle: states one cannot know the position and energy of an electron  Electrons exist in orbital’s of probability  Orbital - the area in space around the nucleus where there is a 90% probability of finding an electron

5. Other Scientists Contributions  Schrödinger  Schrödinger Wave Equation - mathematical solution of an electron’s energy in an atom  quantum mechanical model of the atom – current model of the atom treating electrons as waves.

6. Solutions to the Wave Equation  Quantum Numbers  Wave Equation generates 4 variable solutions  n - size  l – shape: azimuthal quantum  m – orientation  s – spin  Address of an electron

7. Quantum Numbers  n – Primary Quantum Number  Describes the size and energy of the orbital  n is any positive #  n = 1,2,3,4,….  Found on the periodic table  Like the “state” you live in

8. Quantum Numbers  l – Orbital Quantum Number  Sub-level of energy  Describes the shape of the orbital  l = 0,1,2,3,4,….(n-1)  “City” you live in n = 3 l = 0,1,2 n = 2 l = 0,1 n = 1 l = 0

9. Quantum Numbers  l – Orbital Quantum Number  # level = # sublevels  1 st level – 1 sublevel  2 nd level – 2 sublevels  4 th level = 4 sublevels

10. Energy Sublevels Labeled s, p, d, or f –Based on shape of the atom’s orbitals –Each sublevel can only contain at most 2 e-

11. Quantum Numbers  m – Magnetic Quantum Number  Describes the orientation of the orbital in space  Also denotes how many orbital's are in each sublevel  For each sublevel there are 2l +1 orbital's  m = 0, ±1, ±2, ±3, ±l  “Street” you live on

12. Quantum Numbers  Look at Orbital's as Quantum Numbers l = 0 m = 0 Can only be one s orbital l = 1 m = -1, 0, +1 For each p sublevel there are 3 possible orientations, so three 3 orbital's

13. Assigning the Numbers  The three quantum numbers (n, l, and m) are integers.  The principal quantum number (n) cannot be zero.  n must be 1, 2, 3, etc.  The angular quantum number (l ) can be any integer between 0 and n - 1.  For n = 3, l can be either 0, 1, or 2.  The magnetic quantum number (m l ) can be any integer between -l and +l.  For l = 2, m can be either -2, -1, 0, +1, +2.

14. Orbital Rules Energy Level Possible sub-levels Number of Sub-levels n No. of Orbitals n 2 No. of Electrons 2n 2 4 s, p, d, f 41632 3 s, p, d 3918 2 s, p 248 1s112

15.  Read the scenario  Complete the questions  Completed packet due tomorrow  HW: Finish Packet

16. Energy Level Diagrams

17. Aufbau Principle Electrons occupy the lowest energy level orbital available.

18. Aufbau Principle  Aufbau Principal  Lowest energy orbital available fills first  “Lazy Tenant Rule”

19. Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli Every house has a different address

20. Pauli’s Exclusion Principle  Pauli Exclusion Principle  No two electrons have the same quantum #’s  Maximum electrons in any orbital is two (  )

21. Hund’s Rule  Hund’s Rule  When filling degenerate orbital's, electrons will fill an empty orbital before pairing up with another electron.  Empty room rule RIGHTWRONG

22. Outermost sub-shell being filled with electrons

23. The order of sublevel filling is arranged according to increasing energy level. Electrons first fill the 1s sublevel followed by the 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p and 6s Increasing Energy 1s 2p 6s 4s 5s 3s 2s 4d 3d 5p 4p 3p Think about piggy palace….

24. Periodic Table and Electron Configuration

26. Orbital Energy Diagram and Electron Configuration p ______ ______ ______ 3 s ______ p ______ ______ ______ 2 s ______ 1 s ______ An energy diagram for Neon Increasing Energy Electron Spin 1s 2 2s 2 2p x 2 2p y 2 2p z 2 2p 6 1s 2 2s 2 Electron Configuration Notation

27. Orbital Notation  Orbital Notation shows each orbital  O (atomic number 8) ____ ____ ____ ____ ____ ____ 1s 2s 2p x 2p y 2p z 3s  1s 2 2s 2 2p 4 electron configuration!

28. Orbital Notation  Write the orbital notation for S  S (atomic number 16) ____ ____ ____ ____ ____ ____ ____ ____ ____ 1s 2s 2p 3s 3p  1s 2 2s 2 2p 6 3s 2 3p 4  How many unpaired electrons does sulfur have? 2 unpaired electrons!

29. Electron Configuration  Shorthand way of writing electron configuration of atoms 10 Ne: 1s 2 2s 2 2p 6 Elemental Symbol and atomic number Principal energy level Energy sublevel Number of electrons

30. zShorthand Configuration S 16e - Valence Electrons Core Electrons S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 Valence Electrons  Longhand Configuration

31. [Ar] 4s 2 3d 10 4p 2 Noble Gas Configuration  Example - Germanium X X X X X X X X X X X X X

32. Electron Configuration Let’s Practice  P (atomic number 15)  1s 2 2s 2 2p 6 3s 2 3p 3  Ca (atomic number 20)  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2  As (atomic number 33)  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3  W (atomic number 74)  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 4 Noble Gas Configuration [Ne] 3s 2 3p 3 [Ar] 4s 2 [Ar] 4s 2 3d 10 4p 3 [Xe] 6s 2 4f 14 5d 4

33. Electron Configuration Your Turn  N (atomic number 7)  1s 2 2s 2 2p 3  Na (atomic number 11)  1s 2 2s 2 2p 6 3s 1  Sb (atomic number 51)  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 3  Cr (atomic number 24)  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 Noble Gas Configuration [He] 2s 2 2p 3 [Ne] 3s 1 [Kr]5s 2 4d 10 5p 3 [Ar] 4s 2 3d 4

34.  End of information for the test on Thursday 1/14

36. Valence Electrons  Valence Electrons  As (atomic number 33)  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3  The electrons in the outermost energy level.  s and p electrons in last shell  5 valence electrons

37.  Full energy level  Full sublevel  Half full sublevel Stability

38. Exceptions  Copper  Expect: [Ar] 4s 2 3d 9  Actual: [Ar] 4s 1 3d 10  Silver  Expect: [Kr] 5s 2 4d 9  Actual: [Kr] 5s 1 4d 10  Chromium  Expect: [Ar] 4s 2 3d 4  Actual: [Ar] 4s 1 3d 5  Molybdenum  Expect: [Kr] 5s 2 4d 4  Actual: [Kr] 5s 1 4d 5 Exceptions are explained, but not predicted! Atoms are more stable with half full sublevel

39. Stability  Atoms create stability by losing, gaining or sharing electrons to obtain a full octet  Isoelectronic with noble gases +1 +2 -3 -2 0 +3 +4 Atoms take electron configuration of the closest noble gas

40. Stability  Na (atomic number 11)  1s 2 2s 2 2p 6 3s 1  1s 2 2s 2 2p 6 = [Ne] Na 1 Valence electron Metal = Loses Ne

41. Try Some  P -3 (atomic number 15)  1s 2 2s 2 2p 6 3s 2 3p 6  Ca +2 (atomic number 20)  1s 2 2s 2 2p 6 3s 2 3p 6  Zn +2 (atomic number 30)  1s 2 2s 2 2p 6 3s 2 3p 6 3d 10  Last valence electrons (s and p) Full Octet

42. ElementConfiguration notation Orbital notationNoble gas notation Lithium1s 2 2s 1 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 1 Beryllium1s 2 2s 2 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 Boron1s 2 2s 2 p 1 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 1 Carbon1s 2 2s 2 p 2 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 2 Nitrogen1s 2 2s 2 p 3 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 3 Oxygen1s 2 2s 2 p 4 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 4 Fluorine1s 2 2s 2 p 5 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 5 Neon1s 2 2s 2 p 6 ____ ____ ____ ____ ____ 1s 2s 2p [He]2s 2 p 6

43. Half of the distance between nucli in covalently bonded diatomic molecule "covalent atomic radii" Periodic Trends in Atomic Radius Radius decreases across a period Increased effective nuclear charge due to decreased shielding Radius increases down a group Addition of principal quantum levels Determination of Atomic Radius

44. Table of Atomic Radii

45.  Increases for successive electrons taken from the same atom  Tends to increase across a period Electrons in the same quantum level do not shield as effectively as electrons in inner levels Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove  Tends to decrease down a group Outer electrons are farther from the nucleus Ionization Energy: the energy required to remove an electron from an atom

46.  Affinity tends to increase across a period  Affinity tends to decrease as you go down in a period Electrons farther from the nucleus experience less nuclear attraction Some irregularities due to repulsive forces in the relatively small p orbitals Electron Affinity - the energy change associated with the addition of an electron

47. A measure of the ability of an atom in a chemical compound to attract electrons  Electronegativities tend to increase across a period  Electronegativities tend to decrease down a group or remain the same Electronegativity

48. Cations Positively charged ions Smaller than the corresponding atom Anions Negatively charged ions Larger than the corresponding atom Ionic Radii

49. Table of Ion Sizes

50. Summary of Periodic Trends