2. Electrons in Atoms

3. 2 Blimp Chemistry: Hydrogen vs. Helium Blimps float: filled with a gas that is less dense than the surrounding air Early blimps used the gas Hydrogen: flammability lead to the Hindenburg disaster Hydrogen Blimps now use Helium gas: not flammable nor any chemical reactions Why Hydrogen and Helium are so different in Chemical Reaction?

4. 3 Why Patterns for Charges of Common Cations and Anions? Why metals always tend to lose electron(s) whereas nonmetals tend to gain electron(s)? Why Group IA metals (e.g., Li, K) always form cation with +1 charge, Group IIA cations always with +2? Why Group VIIA (halogens) always form anions with –1 charge, Group VIA monatomic anions always –2 charge?

5. 4 Chapter outline From Light to Energy of Electrons in Atom Quantum mechanical description of Atom 1.Principal quantum number: Shell 2.Orientation (shape) of : Subshell 3.Orbitals hold electrons with opposite spin Electron configuration and Orbital Diagram Periodic patterns relate to Electron configuration

6. 5 Electromagnetic Radiation Light: one of the forms of energy Electromagnetic radiation electromagnetic radiation travels in Waves Wave properties Wave speed Height (amplitude) Wave lengthFrequency

7. 6 Electromagnetic Waves How fast Light travels? Velocity c = speed of light = 2.998 x 10 8 m/s in vacuum all types of light energy travel at the same speed Frequency = #peaks pass a point in a second generally measured in Hertz (Hz), Low frequency High Frequency

8. 7 Energy of Electromagnetic Radiation Max Planck: German Physicist Light consists of numerous, individual “packets” of electromagnetic energy, called “photon”. The energy of each photon is proportional to its frequency E photon = hv

9. 8 Visible Light Colors in visible light: Different frequency of electromagnetic radiation: Frequency for visible light: Red < Yellow < Green < Blue < Violet Photons with different frequency have a different amount of energy frequency  Energy of the photons  Energy: Red < Yellow < Green < Blue < Violet

10. 9 Man-made Rainbow? Prism: An optical element that refract (“bend”) light. Light passed through a prism is separated into all its colors - this is called a continuous spectrum Nowadays, a silver-colored CD disc can generate your homemade spectrum!

11. 10 Electromagnetic Spectrum

12. 11 Types of Electromagnetic Radiation By the frequency from low to high Radiowaves : low frequency and energy Microwaves Infrared (IR) Visible: ROYGBIV Ultraviolet (UV) X-rays Gamma rays  high frequency and energy

13. 12 Everyday spectra Common street light (containing mainly Na vapor) Indoor fluorescent light (containing Hg vapor)

14. 13 Light’s Relationship to Matter Atoms can acquire extra energy, but they must eventually release it When atoms emit energy, it always is released in the form of light However, atoms don’t emit all colors, only very specific wavelengths in fact, the spectrum of wavelengths can be used to identify the element

15. 14 Emission Spectrum of Hydrogen

16. 15 Emission spectraEmission spectra: Fingerprint of atoms

18. 17 656.3 486.1 434.1 410.2 Absorption Spectrum of H 2 Emission Spectrum of H 2 Absorption Spectrum “White light”SampleEmission light Emission Spectrum Sample

19. 18 Why Line Spectra? Another way for the same question: Why atoms can only emit or absorb certain amount of energies? An simple guess would be that an atom could only have very specific amounts of energy; When they absorb or release energies (photon), the change in the energies they possess would be certain amount.

20. 19 Bohr Model of the Atom The energy of the atom was quantized The amount of energy in the atom was related to the electron’s position in the atom (Electron Orbit) The atom could only have very specific amounts of energy (n = 1, 2, 3, …)

21. 20 Electron Orbits Electrons travel in orbits around the nucleus more like shells than planet orbits the farther the electron is from the nucleus, the more energy it has

22. 21 Orbits and Energy each Orbit has a specific amount of Energy Energy of each orbit is characterized by an integer n The larger n, the more energy an electron in that orbit has, the farther it is from the nucleus n: quantum number

23. 22 Energy Transitions when the atom gains energy, the electron leaps from a lower energy orbit to higher energy orbit: “Excitation” (_______ spectra) when the electron leaps from a higher energy orbit to lower energy orbit, energy is emitted as a photon of light: “Relaxation” (_______ spectra)

24. 23 Bohr Model of the Atom Hydrogen Spectrum

25. 24 Quantum Physicists including Schrödinger: Electrons show up with a particular probability at certain location of the atom Orbital: A region where the electrons show up a very high probability when it has a particular amount of energy generally set at 90 or 95% Quantum-Mechanical Orbitals

26. 25 Quantum-Mechanical Model: Quantum Numbers Three quantum numbers: quantize the energy Principal quantum number, n, specifies the main energy level for the orbital the higher n value, the higher energy of the electrons, the further away electrons are located from the nucleus

27. 26 Quantum-Mechanical: Quantum Numbers Principal energy shell has one or more Subshells the number of subshells = the Principal quantum number n = 1, one subshell; n = 2, two subshells; n = 3, three subshells Subshell Quantum numbers: s, p, d, f each Subshell has orbitals with a particular shape the shape represents the probability map  90% probability of finding electron in that region

28. 27 Shapes of Subshells p Orbitals: p x, p y, p z s Orbital d Orbitals

29. 28 f orbitals 28 Tro: Chemistry: A Molecular Approach, 2/e

30. 29 How does the 1s Subshell Differ from the 2s Subshell? (colors: signs of wavefunction)

31. 30 Shells & Subshells

32. 31 Subshells and Orbitals Among the subshells of a principal shell, slightly different energies: s < p < d < f each subshell contains one or more Orbitals s : 1 orbital p : 3 orbitals d : 5 orbitals f : 7 orbitals within one subshell, different orbitals have the same energy. Example: 2p x, 2p y and 2p z

33. 32 Energy 1s 7s 2s 2p 3s 3p 3d 6s 6p 6d6d 4s 4p 4d 4f 5s 5p 5d 5f

34. 33 Electron Configurations Definition: The distribution of electrons into the various energy shells (n = 1,2,3,…) and subshells (s, p, d, f) in an atom in its ground state Each energy shell and subshell has a maximum number of electrons it can hold Subshell s = ___, p = ___, d = ___, f = ___  Shell n: 1 = 2e, 2 = 8e, 3 = 18e, 4 = 32e Electrons fill in the energy shells and subshells in order of energy, from low energy up Aufbau Principal (“Construction” in German)

35. 34 Order of Subshell Filling in Ground State Electron Configurations 1s2s2p3s3p3d4s4p4d4f5s5p5d5f6s6p6d7s1s2s2p3s3p3d4s4p4d4f5s5p5d5f6s6p6d7s 1. Diagram putting each energy shell on a row and listing the subshells, (s, p, d, f), for that shell in order of energy, (left-to-right) 2. draw arrows through the diagonals, looping back to the next diagonal each time

36. 35 Spinning Electron(s) in Orbital Experiments (Stern and Gerlach) showed Electrons spin on an axis generating their own magnetic field Pauli Exclusion Principle each Orbital may have a maximum of 2 electrons, with opposite spin Two electrons sharing the same orbital must have Opposite spins so their magnetic fields will cancel analogous to two bar magnets in parallel: only opposite alignment could stabilize each other.

37. 36 Orbital Diagrams often an orbital as a square the electrons in that orbital as arrows the direction of the arrow represents the spin of the electron orbital with 1 electron unoccupied orbital orbital with 2 electrons

38. 37 How electrons in an atom are filled into orbitals 1. How Electrons fill subshells with multiple orbitals 2. How Electrons fill subshells with higher n number first

39. 38 Filling the Orbitals in a Subshell with Electrons Energy shells fill from lowest energy to high 1 → 2 → 3 → 4 Subshells fill from lowest energy to high s → p → d → f Orbitals of the same subshell have the same energy. Three 2p orbitals; Five 3d orbitals  Electrons prefer “spreading out” in orbitals of same subshell before they pair up in orbitals. Hund’s Rule Example: 2p 3  _  _  _ instead of   ____

40. 39 Electron Configuration of Atoms in their Ground State Electron configuration: a listing of the subshells in order of filling with the number of electrons in that subshell written as a superscript Kr = 36 electrons = 1s 2 2s 2 2p 6 ____________________ a shorthand way : use the symbol of the previous noble gas in [] for the inner electrons, then just write the last set Rb = 37 electrons = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 = ______

41. 40 1.Determine the number of electrons: Atomic number = #protons = #electrons = _____ Example: Ground State Orbital Diagram and Electron Configuration of Magnesium 2.Draw boxes to represent the subshells 3.Add one electron to each box in a set, then pair the electrons before filling the next subshell When pair, put in opposite arrows: ___ 4.Use the diagram to write the electron configuration (1s 2 2s 2 … )

42. 41 More example: Write Electron Configuration and Orbital Diagram for a Phosphorus atom Phosphorus: ____ electrons

43. 42 Valence Electrons Definition: the electrons in all the subshells with the highest principal energy shell Example: electrons in bold Mg = [Ne]3s 2 O = [He]2s 2 2p 4 Br = [Ar]4s 2 3d 10 4p 5 Core electrons: electrons in lower energy shells Chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the Number of Valence electrons

44. 43 Valence Electrons Rb = 37 electrons = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 the highest principal energy shell is the 5 th : ___ valence electron + ___ core electrons Kr = 36 electrons = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 the highest principal energy shell is the 4 th : ___ valence electrons + ___ core electrons

45. 44 Electrons Configurations and the Periodic Table

46. 45 Electron Configurations from the Periodic Table Example: Be 2s 2 B 2s 2 2p 1 C 2s 2 2p 2 N 2s 2 2p 3 O 2s 2 2p 4 Elements in the same period (row) have Valence Electrons in the same principal energy shell. #Valence electrons increases by 1 from ____ to ___ Example: IIA: Be 2s 2 Ca 3s 2 Sr 4s 2 Ba 5s 2 VIIA: F 2s 2 2p 5 Cl 3s 2 3p 5 Br 4s 2 4p 5 I 5s 2 5p 5 Elements in the same group have the same _______ _______ and same kind of subshell

47. 46 Electron Configuration & the Periodic Table Elements in the same Group have similar chemical and physical properties  their valence shell electron configuration is the same No. Valence electrons for the main group elements is the same as the Group Number Example: Group IA: ns 1 ; Group IIIA: ns 2 np 1 Group VIIA: ns 2 np 5

48. 47 s1s1 s2s2 d 1 d 2 d 3 d 4 d 5 d 6 d 7 d 8 d 9 d 10 p 1 p 2 p 3 p 4 p 5 s2s2 p6p6 f 1 f 2 f 3 f 4 f 5 f 6 f 7 f 8 f 9 f 10 f 11 f 12 f 13 f 14 12345671234567 Electron Configuration & the Periodic Table

49. 48 Electron Configuration from the Periodic Table Inner electron configuration = Noble gas of the preceding period Outer electron configuration: from the preceding Noble gas  the next period (Subshells)  Element the valence energy shell = the period number the d block is always one energy shell below the period number and the f is two energy shells below

50. 49 Electron Configuration from the Periodic Table P = [Ne]3s 2 3p 3 P has 5 valence electrons 3p33p3 P Ne 12345671234567 1A 2A 3A4A5A6A7A 8A 3s23s2

51. 50 Electron Configuration from the Periodic Table As = [Ar]4s 2 3d 10 4p 3 As has 5 valence electrons As 12345671234567 1A 2A 3A4A5A6A7A 8A 4s24s2 Ar 3d 10 4p34p3

52. 51 Electron configuration & Chemical Reactivity Chemical properties of the elements are largely determined by No. Valence electrons Why elements in groups? Since elements in the same column have the same #valence electrons, they show similar properties

53. 52 Electron Configuration: Noble Gas Noble gases have 8 valence electrons except for He, which has only 2 electrons Noble gases are especially nonreactive He and Ne are practically inert  The reason: the electron configuration of the noble gases is especially stable

54. 53 Everyone Wants to Be Like a Noble Gas! Alkali Metals (Group 1A) have one more electron than the previous noble gas, [NG]ns 1 tend to lose their extra ONE electron, resulting in the same electron configuration as a noble gas forming a cation with a 1+ charge Na  Na + Li  Li +

55. 54 Everyone Wants to Be Like a Noble Gas! Halogens (Group 7A) one fewer electron than the next noble gas: [NG]ns 2 np 5 Reactions with Metals: tend to gain an electron and attain the electron configuration of the next noble gas: [NG]ns 2 np 5 + 1e  [NG]ns 2 np 6 forming an anion with charge 1-: Cl  Cl - Reactions with Nonmetals: tend to share electrons so that each attains the electron configuration of a noble gas

56. 55 Everyone Wants to Be Like a Noble Gas! Summary Alkali Metals as a group are the most reactive metals they react with many things and do so rapidly Halogens are the most reactive group of nonmetals one reason for their high reactivity: they are only ONE electron away from having a very stable electron configuration the same as a noble gas

57. 56 Stable Electron Configuration And Ion Charge Metals:  Cations by losing enough electrons to get the same electron configuration as the previous noble gas Nonmetals:  Anions by gaining enough electrons to get the same electron configuration as the next noble gas

58. 57 Example: Write Electron Configuration for the following ions Sulfide ion: charge = ___, #electrons = ___ Aluminum ion: charge = ___, #electrons = ___

59. 58 Trends in Atomic Size

60. 59 Trends in Atomic Size ↑ Down a group valence shell farther from nucleus effective nuclear charge fairly close ↓Across a period (left to right) adding electrons to same valence shell effective nuclear charge increases valence shell held closer

61. 60 Metallic Character Metals malleable & ductile shiny, lusterous, reflect light conduct heat and electricity most oxides basic and ionic form cations in solution lose electrons in reactions – oxidized Nonmetals brittle in solid state dull electrical and thermal insulators most oxides are acidic and molecular form anions and polyatomic anions gain electrons in reactions - reduced

62. 61 Trends in Metallic Character

63. 62 Electron Configuration Affects the Size of Atoms and Metallic Character: Within a Group Within the same Group, from top to bottom: As valence shell number n increases  valence electron(s) further away from the nucleus  _________ Atomic Radius  weaker Coulombic force (electrostatic force) withholding valence electrons  electrons easier to be lost  ___________metallic character

64. 63 4 p + 2e - 12 p + 2e - 8e - 2e - Be (4p + & 4e - ) Example: Group IIA Mg (12p + & 12e - ) Ca (20p + & 20e - ) 20 p + 2e - 8e - 2e - 8e -

65. 64 Electron Configuration Affects the Size of Atoms and Metallic Character: Over the Period Within the same Period (row), from left to right: Same valence shell number n As Nucleus has increasing number of protons (p + )  Stronger Coulombic force (electrostatic force) withholding valence electrons  Valence Electrons closer the nuclues  ________ Atomic Radius  Valence electrons harder to be lost  ________ metallic character

66. 65 Li (3p + & 3e - ) Example: Period 2 From Li (3 protons) to Ne (10 protons), attraction increases Be (4p + & 4e - )B (5p + & 5e - ) 6+ 2e - 4e - C (6p + & 6e - ) 8+ 2e - 6e - O (8p + & 8e - ) 10+ 2e - 8e - Ne (10p + & 10e - ) 2e - 1e - 3+ 2e - 4+ 2e - 3e - 5+

67. 66 Practice – Choose the Larger Atom in Each Pair C or O Li or K C or Al Se or I?

68. 67 Practice – Choose the More Metallic Element in Each Pair Sn or Te Si or Sn Br or Te Se or I?