Bohr Model of the Atom

2. 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

3. 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’sof probability • Orbital - the area in space around the nucleus where there is a 90% probability of finding an electron

4. 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.

5. 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

6. 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

7. Quantum Numbers n = 3 l = 0,1,2 n = 2 l = 0,1 n = 1 l = 0 • 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

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

9. 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-

10. 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 are2l +1 orbital's • m = 0, ±1, ±2, ±3, ±l • “Street” you live on

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

12. 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, lcan be either 0, 1, or 2. • The magnetic quantum number (ml) can be any integer between -l and +l. • For l = 2, m can be either -2, -1, 0, +1, +2.

13. Orbital Rules

14. Hog Hilton Time • Read the scenario • Complete the questions • Completed packet due tomorrow • HW: Finish Packet

15. Energy Level Diagrams

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

17. Aufbau Principle Aufbau Principal • Lowest energy orbital available fills first • “Lazy Tenant Rule”

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

19. Pauli’s Exclusion Principle • No two electrons have the same quantum #’s • Maximum electrons in any orbital is two () Pauli Exclusion Principle

20. Hund’s Rule • When filling degenerateorbital's, electrons will fill an empty orbital before pairing up with another electron. • Empty room rule RIGHT WRONG Hund’s Rule

21. Outermost sub-shell being filled with electrons

22. 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 6s Think about piggy palace…. Increasing Energy 5s 5p 4s 4p 4d 3s 3p 3d 2s 2p 1s

23. Periodic Table and Electron Configuration

24. Periodic Table and Electron Configuration

25. Orbital Energy Diagram and Electron Configuration 2px2 2py2 2pz2 Electron Spin Increasing Energy p ______ ______ ______ 3 s ______ p ______ ______ ______ 2 s ______ 1 s ______ 1s2 2s2 1s2 2s2 2p6 Electron Configuration Notation An energy diagram for Neon

26. Orbital Notation • Orbital Notation shows each orbital • O(atomic number 8) ____ ____ ____ ____ ____ ____ 1s2s2px2py2pz3s • 1s22s22p4electron configuration!

27. Orbital Notation 2 unpaired electrons! • Write the orbital notation for S • S(atomic number 16) ____ ____ ____ ____ ____ ____ ____ ____ ____ 1s2s2p3s 3p • 1s22s22p63s23p4 • How many unpaired electrons does sulfur have?

28. Electron Configuration Shorthand way of writing electron configuration of atoms 10Ne: 1s2 2s2 2p6 Number of electrons Energy sublevel Elemental Symbol and atomic number Principal energy level

29. Valence Electrons Longhand Configuration Core Electrons Valence Electrons 2p6 S 16e- 2s2 1s2 3s2 3p4 • Shorthand Configuration S 16e- [Ne]3s2 3p4

30. Noble Gas Configuration Example - Germanium X X X X X X X X X X X X X [Ar] 4s2 3d10 4p2

31. Electron Configuration Let’s Practice • P (atomic number 15) • 1s22s22p63s23p3 • Ca (atomic number 20) • 1s22s22p63s23p64s2 • As (atomic number 33) • 1s22s22p63s23p64s23d104p3 • W (atomic number 74) • 1s22s22p63s23p64s23d104p65s24d105p66s24f145d4 Noble Gas Configuration [Ne]3s23p3 [Ar]4s2 [Ar]4s23d104p3 [Xe]6s24f145d4

32. Electron Configuration Your Turn • N (atomic number 7) • 1s22s22p3 • Na (atomic number 11) • 1s22s22p63s1 • Sb(atomic number 51) • 1s22s22p63s23p64s23d104p65s24d105p3 • Cr (atomic number 24) • 1s22s22p63s23p64s23d4 Noble Gas Configuration [He]2s22p3 [Ne]3s1 [Kr]5s24d105p3 [Ar]4s23d4

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

35. Valence Electrons • Valence Electrons • As (atomic number 33) • 1s22s22p63s23p64s23d104p3 • The electrons in the outermost energy level. • s and p electrons in last shell • 5 valence electrons

36. Full energy level Full sublevel Half full sublevel Stability

37. Exceptions Exceptions are explained, but not predicted! Atoms are more stable with half full sublevel • Copper • Expect: [Ar] 4s2 3d9 • Actual: [Ar] 4s1 3d10 • Silver • Expect: [Kr] 5s2 4d9 • Actual: [Kr] 5s1 4d10 • Chromium • Expect:[Ar] 4s2 3d4 • Actual: [Ar] 4s1 3d5 • Molybdenum • Expect: [Kr] 5s2 4d4 • Actual: [Kr] 5s1 4d5

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

39. Stability 1 Valence electron Metal = Loses Ne Na • Na (atomic number 11) • 1s22s22p63s1 • 1s22s22p6 = [Ne]

40. Try Some Full Octet • P-3(atomic number 15) • 1s22s22p63s23p6 • Ca+2(atomic number 20) • 1s22s22p63s23p6 • Zn+2(atomic number 30) • 1s22s22p63s23p63d10 • Last valence electrons (s and p)

42. Determination of Atomic Radius 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

43. Table of Atomic Radii

44. Ionization Energy: the energy required to remove an electron from an atom • 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

45. Electron Affinity - the energy change associated with the addition of an electron • 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

46. Electronegativity 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

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

48. Table of Ion Sizes

49. Summary of Periodic Trends