1. Atomic Theory & Periodicity
Chapters 6 & 7

2. Light
How does an incandescent light bulb work?

3. Light Electricity is turned on Metal filament heats up
Hot metal gives off light
Blackbody Radiation

4. Blackbody radiation

5. Blackbody Radiation
Max Planck proposed that the energy of light given off could only be emitted in little bundles, called quanta
Bundles of light were multiples of a smaller unit
E=h𝜈
h = x 10-34Js
Planck’s Constant

6. Photoelectric Effect
Light shining on a photo-sensitive metal plate will emit electrons.

7. Photoelectric Effect
Frequency must be above a minimum (threshold) frequency
Brighter light (higher intensity) produces more electrons, but with the same energy
Light with higher frequency will emit electrons with higher energy

8. Photoelectric Effect

9. Photoelectric Effect ℎ𝜈= 𝐾𝐸 𝑚𝑎𝑥 + 𝜙
Einstein expanded on Planck’s work to explain photoelectric effect
Nobel Prize in 1921
Energy of Photon = Energy of ejected electron + work needed to eject electron (work function, Φ)
ℎ𝜈= 𝐾𝐸 𝑚𝑎𝑥 + 𝜙

10. Atomic Model Review Dalton’s Hard Sphere Model Plum Pudding Model
Atom is indivisible, indestructible sphere
Plum Pudding Model
Uniform positive sphere with negative electrons embedded within
Rutherford Model
Atom is mostly empty space, dense positive nucleus, electrons randomly outside nucleus

11. Bohr Model Rutherford Model could not explain chemical properties
Niels Bohr proposed electrons in energy levels
Valence electrons
Electrons can move between energy levels by gaining or losing energy

12. Electron Transitions
When electrons gain energy they move to higher energy levels
When electrons drop to lower energy levels they release energy as light
Energy is related to the distance between energy levels

13. 𝑐=𝜆𝜈 𝐸=ℎ𝜈= ℎ𝑐 𝜆 Light Waves review Wave Equation Energy Equation
Color is based on frequency, 𝜈
𝑐=𝜆𝜈
𝐸=ℎ𝜈= ℎ𝑐 𝜆

14. Electromagnetic Spectrum

15. Energy Practice
Calculate the energy of a photon of yellow light that has a wavelength of 589 nm.
Calculate the energy of a mole of photons of yellow light with a wavelength of 589 nm.
𝐸=ℎ𝜈= ℎ𝑐 𝜆 = (6.626 𝑥 10 −34 )(2.998 𝑥 ) 589 𝑥 10 −9
𝐸=3.37 𝑥 10 −19 𝐽
𝐸=(3.37 𝑥 10 −19 )(6.022 𝑥 )
𝐸=2.03 𝑥 𝐽/𝑚𝑜𝑙

17. Quantum Mechanics Physics of the small New in the early 20th century
High level math
Thought experiments
Led to many advances in technology

18. Schrödinger’s Cat Thought Experiment
Cat, vial of poison, Geiger counter with radioactive sample in a sealed box.
Can’t know if cat is alive or dead without interrupting the experiment (opening the box)
The cat is considered BOTH alive and dead

19. Heisenberg Uncertainty
Limit to what we can know
The more precisely the momentum of a particle is known the less precisely the position can be known
∆𝑥∆𝑝≥ ℏ 2

20. Schrödinger Wave Equation
Differential equation for wave functions of particles
𝑖ℏ 𝛿 𝛿𝑡 Ψ 𝑟,𝑡 =[ − ℏ 2 2𝑚 𝛻 2 +𝑉(𝑟,𝑡)]Ψ(𝑟,𝑡)

21. Wave Equation Solutions
Solving the wave equation for an electron provides wave functions that include quantum numbers that describes where there is a high probability of finding an electron

22. Principle Quantum Number
Denotes the principle energy level n
Integers equal to or larger than 1
n = 1, 2, 3, …

23. Angular Quantum Number
Denotes the sublevel within an energy level
ℓ (lower case cursive L)
Integers ranging from 0 to n–1
0 = s, 1 = p, 2 = d, 3 = f, 4 = g,….

24. Magnetic Quantum Number
Denotes the orbitals in each sublevel mℓ
Integers ranging from –ℓ to +ℓ
ℓ = 0, mℓ = 0 (1 s orbital)
ℓ = 1, mℓ = –1, 0, +1 (3 p orbitals)
ℓ = 2, mℓ = –2, –1, 0, +1, +2 (5 d orbitals)

25. Spin Quantum Number
Describes the magnetic property of the electron, “spin” ms
½ or –½

26. Pauli Exclusion Principle
2 electrons in an atom can’t have the same set of quantum numbers
Each orbital can hold up to 2 electrons
Must be different “spin”

27. Ψ2 Probability density and radial distribution function
Describes the shape of an orbital

28. Orbitals

29. Orbitals

30. Orbitals

31. Electron Configurations
Describes the number of electrons in an atom and their location
1s2
Number of electrons
in sublevel
Energy level
Sublevel

32. Aufbau Principle Electrons will fill the lowest available orbitals
Na (Regents)
1s22s22p63s1 (AP)

33. Electron Config Examples
Be 1s22s2
C 1s22s22p2
F 1s22s22p5
S 1s22s22p63s23p4

34. Electron Configurations
Sublevels overlap each other
More overlap the farther away from the nucleus
K 1s22s22p63s23p64s1
4s sublevel is lower in energy than the 3d sublevel

35. Electron Orbital Configuration
Sublevel order 1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f 5g
6s 6p 6d 6f 6g 6h
7s 7p 7d 7f 7g 7h 7i

36. Condensed Electron Config’s
Sr 1s22s22p63s23p63d104s24p65s2
Use noble gas for inner electrons
Sr [Kr]5s2

37. Orbital Diagram Visual representation Electrons as arrows H He
Lines or boxes
Electrons as arrows
Different “spin”, different arrow direction H He 1s 1s

38. Orbital Diagram 1s 2s Li Be B 1s 2s 1s 2s 2p

39. Hund’s Rule
Lowest energy is achieved when the number of electrons with the same spin is maximized C N 1s 2s 2p 1s 2s 2p

40. Electron Config Anomalies
Because of how close in energy some sublevels are, some transition metals have partially filled ns sublevels with partially filled (n-1)d sublevels
Sometimes ns = 0
Have to be determined experimentally

41. Electron Config Anomalies
Cr [Ar]4s23d4 [Ar]4s13d5
Cu [Ar]4s23d9 [Ar]4s13d10

42. Magnetism Paramagnetism Diamagnetism
Weak attraction to magnetic field due to unpaired electrons
Diamagnetism
Weak repulsion to magnetic field due to paired electrons

44. Photoelectron Spectroscopy
Experimental method used to determine the electronic structure of an atom
Photoemission Spectroscopy
PES
Based on the Photoelectric Effect
specifically the Work Function

45. Photoelectric Effect (Review)
Light shining on a photo-sensitive metal plate will emit electrons.
Energy of Photon = Energy of ejected electron + work needed to eject electron (work function, Φ)
ℎ𝜈= 𝐾𝐸 𝑚𝑎𝑥 + 𝜙

46. Photoelectron Spectroscopy
Graph shows relative number of electrons and amount of energy required to remove them.
Binding energy
Often energy scale on x axis is reversed
Zero to the right end
Not always, have to pay attention!

47. Photoelectron Spectroscopy
Inner most electrons
Outer most electrons

49. Periodic Table Developed by Mendeleev in 1869
Initially arranged by atomic mass
Certain elements rearranged by chemical properties
Predicted missing elements with very accurate prediction of missing element’s properties

50. Periodic Table
Henry Moseley determined atomic numbers using X-ray Spectroscopy
Proved Mendeleev’s Periodic Table correct
Died in WWI at age of 27
Most likely would have won 1916 Nobel Prize
Not awarded to anyone in 1916

51. Periodic Table

52. Periodicity (Regents Review)
Atomic Radius – Size of an atom
Ionic Radius – Size of an ion
Ionization Energy – energy required to remove outermost electron
Electronegativity – ability to attract electrons while in a compound

53. Periodicity (Regents Review)
Atomic Radius
Decreases left to right
Increases top to bottom
Ionic Radius

54. Periodicity (Regents Review)
Ionization Energy
Increases left to right
Decreases top to bottom
Electronegativity

55. Periodicity (Regents Review)
Metallic Character
Decreases left to right
Increases top to bottom

56. Periodicity (Regents Review)
Valence Electrons
Electrons in outermost occupied energy level
Elements in same group have similar physical and chemical properties because they have the same number of valence electrons

57. Coulomb’s Law
Electrons in an atom are attracted to the protons in the nucleus
Electrons are also repelled by other electrons in the atom
Attraction of outer electrons to nuclear protons is diminished by inner electrons
“Shielding”

58. Effective Nuclear Charge, Zeff
Actual attractive force felt by electrons
Zeff = Z – S
Z = nuclear charge
S = shielding constant
Approx = # of core electrons

60. Atomic Radius 2 types of measurements
van der Waals radius = non-bonding
Covalent radius = bonding radius

61. Atomic Radius Atomic Radius decreases across a period
Additional electrons are in the same energy level
Effective nuclear charge is increasing

62. Atomic Radius Atomic Radius increases down a group
Additional electrons are in next higher energy level
Effective nuclear charge doesn’t change (much)

63. Atomic Radius

64. Ionic Radius Metals tend to lose electrons when they form ions
Valence electrons are lost
The “new” valence electrons experience a larger effective nuclear charge, making them even smaller
Less repulsion between electrons
Na 1s22s22p63s1
Na+ 1s22s22p6

65. Ionic Radius Nonmetals tend to gain electrons when they form ions
More valence electrons are added
The “new” valence electrons experience a smaller effective nuclear charge, making them larger
More repulsion between electrons
F 1s22s22p5
F- 1s22s22p6

66. Ionic Radius For cations, ionic radius decreases across a period
For anions, ionic radius decreases across a period
Additional electrons are in the same energy level
Effective nuclear charge is increasing

67. Ionic Radius Ionic radius increases down a group
Additional electrons are in next higher energy level
Effective nuclear charge doesn’t change (much)

68. Ionic Radius

69. Isoelectronic Series Ions have the same number of electrons
Ionic radius decreases with an increasing nuclear charge

70. Ionization Energy Energy required to remove electron
Each successive electron requires more energy to remove
1st < 2nd < 3rd < 4th…..
After valence electrons, there is a large increase in energy to remove

71. Ionization Energy

72. Ionization Energy Ionization energy increases across a period
Effective nuclear charge is increasing
Atomic radius is decreasing
Ionization energy decreases down group
Valence electron is farther from nucleus
Atomic radius increases

73. Ionization Energy

74. Ionization Energy Anomalies
Transition from group 2 to group 13
Electron is removed from a farther out sublevel, p-sublevel instead of s-sublevel
Transition from group 15 to group 16
Electron being removed from group 16 is from a paired orbital, resulting in a half-filled sublevel
Repulsion due to pairing helps too

75. Electronegativity Electronegativity increases across a period
Effective nuclear charge is increasing
Atomic radius is decreasing
Electronegativity decreases down group
Atomic radius is increasing