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4. 3 Orbitals and the Periodic Table Orbitals grouped in s, p, d, and f orbitals (sharp, principle, diffuse, and fundamental(fine))Orbitals grouped in s, p, d, and f orbitals (sharp, principle, diffuse, and fundamental(fine)) s orbitals p orbitals d orbitals f orbitals

5. 4 Arrangement of Electrons in Atoms Electrons in atoms are arranged as LEVELS (n) SUBLEVELS (l) ORBITALS (m l )

6. 5 Energy Levels n = 1 n = 2 n = 3 n = 4 To determine max # e- in energy level: 2n 2

7. 6 Sublevels The most probable area to find these electrons takes on a shapeThe most probable area to find these electrons takes on a shape So far, we have 4 shapes. They are named s, p, d, and f.So far, we have 4 shapes. They are named s, p, d, and f. No more than 2 e- assigned to an orbital – one spins clockwise, one spins counterclockwiseNo more than 2 e- assigned to an orbital – one spins clockwise, one spins counterclockwise

8. 7 Types of Orbitals s orbital p orbital d orbital

9. 8 Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

10. 9 p Orbitals this is a p sublevel with 3 orbitals These are called x, y, and z this is a p sublevel with 3 orbitals These are called x, y, and z There is a PLANAR NODE thru the nucleus, which is an area of zero probability of finding an electron 3p y orbital

11. 10 p Orbitals The three p orbitals lie 90 o apart in spaceThe three p orbitals lie 90 o apart in space

12. 11 d Orbitals d sublevel has 5 orbitalsd sublevel has 5 orbitals

13. 12 The shapes and labels of the five 3d orbitals.

14. 13 f Orbitals For l = 3, ---> f sublevel with 7 orbitals

15. 14 ARRANGEMENTARRANGEMENT The shape, size, and energy of each orbital is a function of 3 quantum numbers which describe the location of an electron within an atom or ion 1-7 energy level spdf shape of sublevel spdf shape of sublevel p x, p y, p z orientation of orbital

16. 15 s orbitals d orbitals Number of orbitals Number of electrons p orbitals f orbitals √ LEARNING CHECK How many electrons can be in a sublevel? (Remember: A maximum of two electrons can be placed in an orbital.)

17. 16 Electron Configurations A list of all the electrons in an atom (or ion) Must go in order (Aufbau principle) 2 electrons per orbital, maximum We need electron configurations so that we can determine the number of electrons in the outermost energy level. These are called valence electrons. The number of valence electrons determines how many and what this atom (or ion) can bond to in order to make a molecule 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 … etc.

18. 17 Electron Configurations 2p 4 Energy Level Sublevel Number of electrons in the sublevel 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 … etc.

19. 18 √ LEARNING CHECK Write the electron configuration for the following elements: H Li N

20. 19 √ SOLUTIONS H 1 S 1 Li 1 S 2 2 S 1 N 1 S 2 2 S 2 2 P 3

21. 20 Why are d and f orbitals always in lower energy levels? d and f orbitals require LARGE amounts of energy to get into it.. It’s better (lower in energy) to skip a sublevel that requires a large amount of energy (d and f orbtials) for one in a higher level but lower energy

22. 21 Shorthand Notation A way of abbreviating long electron configurations Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration

23. 22 Shorthand Notation Step 1: It’s the Showcase Showdown! Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ]. Step 2: Find where to resume by finding the next energy level. Step 3: Resume the configuration until it’s finished.

24. 23 Shorthand Notation Chlorine You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s 2 2s 2 2p 6 The next energy level after Neon is 3 –Longhand is 1s 2 2s 2 2p 6 3s 2 3p 5 So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17 [Ne] 3s 2 3p 5

25. 24 √ LEARNING CHECK Write the shorthand notation for each of the following atoms: Cl K I

26. 25 √ SOLUTIONS Cl [Ne] 3s 2 3p 5 K [Ar] 4s 1 I [Kr] 5s 2 4p 10 5p 5

27. 26 ORBITAL DIAGRAMS/NOTATIONS

28. 27 Orbital Diagrams Graphical representation of an electron configuration One arrow represents one electron Shows spin and which orbital within a sublevel Same rules as before (Aufbau principle, d 4 and d 9 exceptions, two electrons in each orbital, etc. etc.)

29. 28 LithiumLithium Group 1A Atomic number = 3 1s 2 2s 1 ---> 3 total electrons

30. 29 CarbonCarbon Group 4A Atomic number = 6 1s 2 2s 2 2p 2 ---> 6 total electrons 6 total electrons Here we see for the first time HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible.

31. 30 REMEMBER: 3 Basic Rules Apply ► Aufbau Principle (“A to Z”) States: Each electron occupies the lowest orbital available ► Pauli Exclusion Principle (“Pair”) States: A maximum of 2 electrons may occupy a single atomic orbital, but only if the electron have opposite spins ► Hund Rule (“Half way”) States: a single electron with the same spin must occupy each equal energy orbital before additional electrons with opposite spins can occupy the same orbital

32. 31 Aufbau Principle The physical and chemical properties of elements is determined by the atomic structure. The atomic structure is, in turn, determined by the electrons and which shells, subshells and orbitals they reside in. The rules of placing electrons within shells is known as the Aufbau principle.

33. 32 Pauli’s Exclusion Principle The second major fact to keep in mind is the Pauli Exclusion Principle which states no two electrons can have the same four quantum numbers first three may be similar but the four quantum number must be different one orbital has a maximum of two electrons if the two electrons have opposing spins one would spin up ( +1/2) and the other would spin down (-1/2)

34. 33 Orbital Diagrams Hund’s Rule –In orbitals of EQUAL ENERGY (p, d, and f), place one electron in each orbital before making any pairs –All single electrons must spin the same way I nickname this rule the “Monopoly Rule” In Monopoly, you have to build houses EVENLY. You can not put 2 houses on a property until all the properties has at least 1 house.

35. 34 Electron Spin and Magnetism Diamagnetic : NOT attracted to a magnetic field; actually repelled somewhat; no unpaired e- Paramagnetic : substance is attracted to a magnetic field. Substances with unpaired electrons are paramagnetic Ferromagnetic: I Ferromagnetic: If the atomic structure allows the cooperative parallel alignment of the magnetic dipole moments.

36. 35 Examples Mg Cl Write orbital notation: if it has an unpaired e - it is paramagnetic.

37. 36 Examples Mg Cl Write orbital notation: if it has an unpaired e - it is paramagnetic.

38. 37 Fe orbital notation Iron exhibits FERROMAGNETISM. It is an internal force produced by many unshared electrons that will form alignment and stronger magnetism than paramagnetic forces.

39. 38 OUTMOST ELECTRONS HAVE THE VERY MOST …ENERGY They are called….. valence electrons Valence electrons are sooooo important:::::

40. 39 Draw these orbital diagrams! Oxygen (O) Chromium (Cr) Mercury (Hg)

41. 40 Valence electrons determine all of the following properties 1) Chemical 2) Electrical 3) Thermal 4) Optical

42. 41 Lewis Dot Diagrams (Electron Dot Diagrams) These diagrams use the element symbols surrounded by 1 dot for every valence electron.

43. 42 Valence Electrons Electrons are divided between core and valence electrons B 1s 2 2s 2 2p 1 Core = [He], valence = 2s 2 2p 1 Br [Ar] 3d 10 4s 2 4p 5 Core = [Ar] 3d 10, valence = 4s 2 4p 5

44. 43 Rules of the Game No. of valence electrons of a main group atom = Group number (for A groups) Atoms like to either empty or fill their outermost level. Since the outer level contains two s electrons and six p electrons (d & f are always in lower levels), the optimum number of electrons is eight. This is called the octet rule.

45. 44 ION CONFIGURATIONS

46. 45 Keep an Eye On Those Ions! Electrons are lost or gained like they always are with ions… negative ions have gained electrons, positive ions have lost electrons The electrons that are lost or gained should be added/removed from the highest energy level (not the highest orbital in energy!)

47. 46 Keep an Eye On Those Ions! Tin Atom: [Kr] 5s 2 4d 10 5p 2 Sn +4 ion: [Kr] 4d 10 Sn +2 ion: [Kr] 5s 2 4d 10 Note that the electrons came out of the highest energy level

48. 47 Keep an Eye On Those Ions! Bromine Atom: [Ar] 4s 2 3d 10 4p 5 Br - ion: [Ar] 4s 2 3d 10 4p 6 Note that the electrons went into the highest energy level

49. 48 √ LEARNING CHECK Write the longhand notation for these: F - Li + Mg +2 Write the shorthand notation for these: Br - Ba +2 Al +3

50. 49 √ SOLUTIONS Write the longhand notation for these: F -1s22s2 Li + Mg +2 Write the shorthand notation for these: Br - Ba +2 Al +3

51. 50 Lanthanide Element Configurations 4f orbitals used for Ce - Lu and 5f for Th - Lr

52. 51 Ion Configurations To form anions from elements, add 1 or more e- from the highest sublevel. P [Ne] 3s 2 3p 3 + 3e- ---> P 3- [Ne] 3s 2 3p 6 or [Ar]

53. 52 What is the ultimate goal Elements are most stable when all electron shells are filled. The outer shell is filled when 8 electrons are in it.

54. 53 The Octet rule So, an element will either empty its outer shell (get to 0) Or fill its shell (get to 8)