1. Electron Configurations

2. Bohr Bohr proposed that the hydrogen atom has only certain allowable energy states. Bohr suggested that the single electron in a hydrogen atom moves around the nucleus in only certain allowed circular orbits.

3. De Broglie DeBroglie stated that electrons had wave like characteristics Came up with the equation:

4. The Heisenberg Uncertainty Principle Heisenberg concluded that it is impossible to make any measurement on an object without disturbing the object—at least a little. The Heisenberg uncertainty principle states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.

5. The Schrödinger wave equation Combined the ideas of Bohr & DeBroglie The atomic model in which electrons are treated as waves and particles is called the wave mechanical model of the atom or, more commonly, the quantum mechanical model of the atom.

6. Electron Configuration Electron configuration – the arrangement of electrons in an atom There are several rules to writing electron configurations

7. Aufbau Principle Each electron must occupy the lowest energy level first

8. Pauli Exclusion Principle In order for 2 electrons to share an orbital, they must have opposite spins In chemistry we designate spins with arrows. Therefore, if 2 electrons enter an orbital, they must enter 

9. Hund’s Rule A single electron with the same spin must occupy each equal energy orbital before additional electrons will pair up with opposite spins You must fill before you pair For example, say you have 3 orbitals (don’t forget that each orbital can hold at most 2 electrons) You want to add 3 electrons

10. How would they enter? __ __ __   _ __ or   . The answer is …   .

11. Arrow Diagrams Before we begin writing arrow diagrams there are a few things you need to know s – can hold a max of 2 electrons p – can hold a max of 6 electrons d – can hold a max of 10 electrons f – can hold a max of 14 electrons

12. Arrow Diagrams Also s – has 1 orbital p – has 3 orbitals d – has 5 orbitals f – has 7 orbitals What do you notice about the number of orbitals compared to the number of electrons?

13. Arrow Diagrams Lastly, you need to know the sequence of orbitals. I will NOT give you this on a test!

14. Arrow Diagrams Write the arrow diagram for sodium Arrows represent electrons, so before you can start, you need to find the number of electrons. How many electrons does sodium have? 11

15. Arrow Diagrams You will keep going until you reach 11 electrons Where do you always start? 1s 1s  Where to next? 2s 1s  2s 

16. Arrow Diagrams How many electrons have you used so far? 4 (only 7 more to go) What’s next? 2p 1s  2s  2p   

17. Arrow Diagrams Only one more left. Where does it go? 3s 1s  2s  2p    3s 

18. Arrow Diagrams Draw the arrow diagram for Br 1s  2s  2p    3s  3p    4s  3d      4p   ,

19. Electron Configurations Writing electron configurations is just a shorter way to write an arrow diagram You start with 1s and continue the configuration until you get the correct number of elecrtons

20. Electron Configurations Write the full electron configuration for K K (19) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 The number of electrons in each orbital are written as superscripts Remember s = 2, p = 6, d = 10, f = 14 Just follow the chart

21. Electron Configurations Write the full electron configuration for Kr Kr (36) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 Write the full electron configuration for P P (15) 1s 2 2s 2 2p 6 3s 2 3p 3

22. Noble Gas Configurations Noble Gas configuration is just a short hand way to write an electron configuration Steps 1.Find the element 2.Find the noble gas before that element (Group 8A) and place it in [brackets] 3.Move one spot 4.Start the configuration from there and keep going until you get to your element

23. Reading the periodic table s block – the first 2 columns of the periodic table (starts with 1s) p block – Groups 3A-8A, six columns (starts with 2p) d block – the center portion of the periodic table consisting of 10 columns (starts with 3d) f block – the two bottom rows of the periodic table consisting of 14 columns (starts with 4 f)

24. Noble Gas Configurations Write the noble gas configuration for Na Find the Noble Gas before Na [Ne] Move one spot and start the configuration from there [Ne]3s 1

25. Noble Gas Configurations Write the noble gas configuration for Br [Ar]4s 2 3d 10 4p 5 Write the noble gas configuration for Mn [Ar]4s 2 3d 5

26. Final Entry Configuration Final entry configuration – the last thing in an electron configuration It’s like a road map to the element Can Identify the element

27. Final Entry Configuration What is the final entry configuration for Si? Find Si What block are you in (s, p, d, f)? 3p How many spots are you into p? 2 3p 2

28. Final Entry Configuration What is the final entry configuration for Ag? 4d 9 What is the final entry configuration for Cl? 3p 6 What is the final entry configuration for Na? 3s 1

29. Final Entry Configuration What element has the final entry configuration of 4p 3 ? As What element has the final entry configuration of 4d 1 ? Y

30. Periodic Trends & the Periodic Table

31. Periodic Table Periodic Table – arrangement of elements in order of increasing atomic number with elements having similar properties in vertical columns –Groups – vertical columns –Periods – horizontal rows

32. Group Names Group 1A 2A 6A 7A 8A Name Alkali Metals Alkaline Earth Metals Chalcogens Halogens Noble Gases

33. Groups Representative elements – group A elements Transition elements – group B elements

34. Groups The group tell you the number of valence electrons that the element has Valence electrons are electrons in the outermost shell of the atom All group 1A elements have 1 valence electron. Likewise, all group 8A elements have 8 valence electrons.

35. Characteristics Elements in the same group exhibit similar chemical characteristics due to the fact that they all have the same number of valence electrons. The most stable number of valence electrons is 8 This is called an octet

36. Charges Every element wants 8 valence electrons to become stable. They will gain or lose valence electrons to form an octet For example…Group 1A elements have 1 valence electron. They can either gain 7 electrons to have an octet or lose 1. What is easier? Lose 1 If an element loses 1 electron (1 negative charge) what charge will the resulting ion have? +1

37. Charges Let’s go to group 7A. This group has 7 valence electrons It can either loose 7 or gain 1 What is the easiest? Gain 1 What will be the resulting charge if the element gain 1 electron (1 negative charge)?

38. Physical States and Classes of the Elements The majority of the elements are metals. They occupy the entire left side and center of the periodic table. Nonmetals occupy the upper-right-hand corner. Metalloids are located along the boundary between metals and nonmetals

39. Metals Metals are elements that have luster, conduct heat and electricity, and usually bend without breaking.

40. Transition Metals The elements in Groups 3 through 12 of the periodic table are called the transition elements. All transition elements are metals. Many transition metals can have more than one charge

41. Inner Transition Metals In the periodic table, two series of elements, atomic numbers 58-71 and 90-103, are placed below the main body of the table. These elements are separated from the main table because putting them in their proper position would make the table very wide. The elements in these two series are known as the inner transition elements.

42. Inner Transition Metals The first series of inner transition elements is called the lanthanides because they follow element number 57, lanthanum. The second series of inner transition elements, the actinides, have atomic numbers ranging from 90 (thorium, Th) to 103 (lawrencium, Lr).

43. Non Metals Although the majority of the elements in the periodic table are metals, many nonmetals are abundant in nature Most nonmetals don’t conduct electricity, are much poorer conductors of heat than metals, and are brittle when solid. Many are gases at room temperature; those that are solids lack the luster of metals.

44. Properties of Metals and Nonmetals

45. Metalloids Metalloids have some chemical and physical properties of metals and other properties of nonmetals. In the periodic table, the metalloids lie along the border between metals and nonmetals.

46. Electron Dot Structures An electron dot structure consists of the elemental symbol surrounded by dots which represent valence electrons

47. Examples Draw the electron dot structure for Na Draw the electron dot structure for Al Draw the electron dot structure for Br

48. Periodic Trends Periodic Trends are trends that occur across the periodic table and down the periodic table They include: atomic radius, Ionization energy, electro negativity, metallic character, and ionic radius

49. Atomic Radius Atomic Radius – size of the atom  Increases  Decreases

50. Ionization Energy Ionization energy – the ability to pull off 1 electron Increases  Decreases 

51. Electro negativity Electro negativity – the ability of an atom to attract another atom Decreases  Increases 

52. Metallic Character Metallic character – how much like a metal the element is Increases  Decreases 

53. Ionic Radius When you talk about ionic radius, you are comparing an atom and its ion When an atom has a negative charge, you have added electrons Which makes it bigger For example, which will be larger: Cl or Cl -1

54. Ionic Radius When an atom has a positive charge, you have taken away electrons Which makes it smaller For example, which will be larger: Na or Na +1

55. Oxidation Numbers

56. Determining Oxidation Numbers 1.The oxidation number for any uncombined elements or diatomic molecule is ZERO –For example: –C = 0 –O 2 = 0 –S = 0 –Cl 2 = 0

57. 2.The oxidation number for a monatomic ion is its charge –For example: –Ca +2 = +2 –Br -1 = -1 –MgBr 2 = Mg = +2 Br = -1 Determining Oxidation Numbers

58. 3.The oxidation number of Hydrogen is usually +1. The exception is in a metal hydride where the oxidation number will be -1 –For example: –HCl  H = +1 –H 2 SO 4  H = +1 –NaH  H = -1 Determining Oxidation Numbers

59. 4.The oxidation number of oxygen is usually -2 EXCEPT in peroxides. Then it is -1 –For example: –MgO  O = -2 –KClO 3  O = -2 –H 2 O 2  O = -1 Determining Oxidation Numbers

60. 5.In binary compounds (nonmetal + nonmetal) the more electronegative element gets a negative oxidation number. This usually means the positive one is first and the negative one is second –For example: –PCl 3  P = +3 Cl = -1 –CO 2  O = -2 C = +4 Determining Oxidation Numbers

61. 6.The sum of the oxidation numbers for all atoms in a neutral compound is zero –For example: +1 -2 H2OH2O +2 -2 = 0 Determining Oxidation Numbers

62. 7.The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the polyatomic ion For example: +1 X -2 HCO 3 -1 +1+x -6 = -1 X = +4 Determining Oxidation Numbers

63. Oxidation Numbers Determine the oxidation numbers of all species in the following examples: H 2 H = 0 CaCl 2 Ca = +2 Cl = -1 KClO 4 O = -2 K = +1 Cl = +7