1. Chapter 5: Electrons in Atoms The chemical properties of atoms, ions, and molecules are related to the arrangement of the electrons within them

2. Periodic Table Video http://www.learner.org/resources/series61.html#

3. John Dalton (1803) Atom was considered a solid indivisible mass

4. J. J. Thomson (1897) Discovered the electron Plum-pudding model: negatively charged electrons stuck into a lump of positively charged material

5. Plum-pudding Model

6. Ernest Rutherford (1911) Discovered the nucleus (+ charge) Proposed electrons surround a dense nucleus

7. Rutherford’s Model

8. The Bohr Model of the Atom Niels Bohr (1913) Niels Bohr pictured electrons orbiting the nucleus much like planets orbiting the sun. But he was wrong! They’re more like bees around a hive.

9. Bohr’s Model

10. Edwin Schrödinger (1926) Quantum Mechanical Model- mathematical model describing descriptions of electrons in an atom Schrödinger equation- complicated calculus Estimates the probability of finding an electron in a certain position

11. Quantum Mechanical Model

12. Heisenberg Uncertainty Principal Werner Heisenberg: Can find out where the electron is, but not where it is going– or vice versa. Cannot determine both position and momentum

14. Energy Level Region around the nucleus where an electron is likely to be moving Analogy: rungs of a ladder Quantum of energy: amount of energy required to move an electron from its present energy level to the next one

15. Energy Level Diagram Nucleus

16. Energy Levels Represented by principle quantum number (n) n = 1, 2, 3, 4, … Within each energy level, electrons can occupy sub-levels

17. Sub-shells and Orbitals Probability of finding an electron within a specific energy level 4 types: s, p, d, f orbitals Rule: At most per energy level: 1- s orbital, 3-p orbitals, 5-d orbitals, 7-f orbitals (Note: not every energy level has all types of sublevels)

18. s Orbital Shapes The s orbital has a spherical shape centered around the origin of the three axes in space.

19. p Orbital Shapes There are three dumbbell-shaped p orbitals in each energy level above n = 1, each assigned to its own axis (x, y and z) in space.

20. d and f Orbital Shapes

21. Electron Configuration Way in which the electrons are arranged in an atom 3 rules to follow:  Aufbau principle- electrons fill in order of increasing energy  Hund’s rule- when electrons occupy orbitals of equal energy, 1 electron must go in each before pairing up  Pauli-Exclusion principle- 2 electrons max. per orbital

22. Three Rules for Writing Electron Configuration Aufbau Principle: Electrons occupy orbitals of lower energy levels first 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p Increasing Energy

23. Hund’s Rule: Electrons occupy orbitals of the same energy to maximize the number of electrons with the same spin Ex. Filling up a school bus with strangers Three Rules for Writing Electron Configuration

24. Pauli Exclusion Principle: Orbitals may (at most) have 2 electrons and they must have opposite spins

25. Orbital Filling Table

26. Electron Configurations of First Three Series

27. Energy Levels, Sublevels, and Electrons Energy Level (n) Sublevels in main energy level (n sublevels) Number of orbitals per sublevel Number of Electrons per sublevel Number of Electrons per main energy level (2n 2 ) 1s122 2spsp 1313 2626 8 3spdspd 135135 2 6 10 18 4spdfspdf 13571357 2 6 10 14 32

28. Sublevels and Orbitals s-sublevels hold 2 electrons. They have just one orbital. p-sublevels hold 6 electrons, divided between 3 orbitals d-sublevels hold 10 electrons, divided between 5 orbitals s-sublevels hold 14 electrons, divided between 7 orbitals

29. Electron Spin Describes the direction of spin of an electron within a magnetic field Possibilities for electron spin : +1/2, -1/2

30. Sub-levels in each Energy Level The energy level number is equal to the number of sub-levels Question:  How many orbitals are in the third energy level?  How many electrons are in the fourth energy level?

31. 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

32. Forming Ion’s Ion’s are either negatively or positively charged An atom gaining electrons is an anion (-)  Nonmetals An atom losing electrons is a cation (+)  Metals Transition metals have a roman numeral that indicates the charge

33. Ion Examples Sodium (Na): 1s 2 2s 2 2p 6 3s 1 Sodium ion (Na 1+ ): 1s 2 2s 2 2p 6 Fluorine (F): 1s 2 2s 2 2p 5 Fluorine ion (F 1- ): 1s 2 2s 2 2p 6

34. Valence Electrons Electrons that are available for bonding Electrons in the outer most shell (energy level)

35. Group IA (alkali metals) have 1 valence electron

36. Group IIA (alkaline earth metals) have 2 valence electrons

37. Group IIIA elements have 3 valence electrons

38. Group IVA elements have 4 valence electrons

39. Group VA elements have 5 valence electrons

40. Group VIA elements have 6 valence electrons

41. Group VIIA (halogens) have 7 valence electrons

42. Group VIIIA (Noble gases) have 8 valence electrons, except helium, which has only 2

43. Transition metals (“d” block) have 1 or 2 valence electrons

44. Lanthanides and actinides (“f” block) have 1 or 2 valence electrons Lanthanides and actinides (“f” block) have 1 or 2 valence electrons

45. Lewis Dot Structures Diagrams that show the number of valence electrons in an atom or ion Octet Rule  Atoms try to achieve stability by gaining/losing e- to have a filled outer energy level  Filled s & p sublevels (2 + 6 = 8)

46. Dot Notations An atom’s valence electrons can be represented by Lewis dot notations. 1 valence e - X 2 valence e - X 3 valence e - X 4 valence e - X 5 valence e - X 6 valence e - X 7 valence e - X 8 valence e - X

47. Dot Notations – Period 2 Lewis dot notations for the valence electrons of the elements of Period 2. lithium Liberyllium Beboron Bcarbon C nitrogen Noxygen Ofluorine Fneon Ne

48. Lewis Dot Structures of Ion’s Make sure to include brackets and the charge!

49. Exceptions Chromium: [Ar] 4s 1 3d 5 Copper: [Ar] 4s 1 3d 10 Most stable:  Completely filled > Half filled > All others