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

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

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., Na) 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?

4. Electromagnetic Radiation
Light: one of the forms of energy
Electromagnetic radiation
electromagnetic radiation travels in Waves
Wave properties
Wave speed Height (amplitude)
Wave length Frequency

5. Electromagnetic Waves
How fast Light travels?
Velocity c = speed of light = x 108 m/s in vacuum
all types of light energy travel at the same speed
Frequency n = #peaks pass a point in a second
generally measured in Hertz (Hz),
Low frequency High Frequency

6. 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
Ephoton = hv

7. 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

8. Man-made Rainbow?
Prism: An optical element that refract (“bend”) light.
Sunlight 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! 

9. Electromagnetic Spectrum

10. 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

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

12. 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

13. Emission Spectrum of Hydrogen

14. Emission spectra: Fingerprint of atoms

15. Critical Thinking: Neon Lights

16. Emission Spectrum Absorption Spectrum Sample “White light” Sample
Emission light
Absorption Spectrum of H2
Emission Spectrum of H2

17. 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.

18. 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, …)

19. 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

20. 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

21. 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)

22. Bohr Model of the Atom Hydrogen Spectrum

23. Quantum-Mechanical Orbitals
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%

24. 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

25. 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

26. Shapes of Subshells
s Orbital
p Orbitals: px , py , pz
d Orbitals

27. f orbitals
Tro: Chemistry: A Molecular Approach, 2/e 27

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

29. Shells & Subshells

30. 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: 2px, 2py and 2pz

31. 6s 6p 6d 7s 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d
Energy 2s 2p 1s

32. Order of Subshell Filling in Ground State Electron Configurations
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) 1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 7s
2. draw arrows through
the diagonals, looping back
to the next diagonal
each time

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)

34. 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.

35. 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
unoccupied
orbital
orbital with
1 electron
orbital with
2 electrons

36. 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

37. 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: 2p3 _ _ _ instead of  ____

38. 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 = 1s22s22p6____________________
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 = 1s22s22p63s23p64s23d104p65s1 = ______

39. Example: Ground State Orbital Diagram and Electron Configuration of Magnesium
Determine the number of electrons:
Atomic number = #protons = #electrons = _____
Draw boxes to represent the subshells
Add one electron to each box in a set, then pair the electrons before filling the next subshell
When pair, put in opposite arrows: ___
Use the diagram to write the electron configuration (1s22s2 …)

40. More example: Write Electron Configuration and Orbital Diagram for a chlorine atom
chlorine: ____ electrons

41. Valence Electrons
Definition: the electrons in all the subshells with the highest principal energy shell
Example: electrons in bold
Mg = [Ne]3s2 O = [He]2s22p4
Br = [Ar]4s23d104p5
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

42. Valence Electrons Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1
the highest principal energy shell is the 5th :
___ valence electron + ___ core electrons
Kr = 36 electrons = 1s22s22p63s23p64s23d104p6
the highest principal energy shell is the 4th :
___ valence electrons + ___ core electrons

43. Electrons Configurations and the Periodic Table

44. Electron Configurations from the Periodic Table
Example: Be 2s2 B 2s22p1 C 2s22p2 N 2s22p3 O 2s22p4
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 2s2 Ca 3s2 Sr 4s2 Ba 5s2
VIIA: F 2s22p5 Cl 3s23p5 Br 4s24p5 I 5s25p5
Elements in the same group have the same _______ _______ and same kind of subshell

45. 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: ns1 ;
Group IIIA: ns2np1
Group VIIA: ns2np5

46. Electron Configuration & the Periodic Tables1 s2
p1 p2 p3 p4 p5 s2 1 2 3 4 5 6 7 p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14

47. 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

48. Electron Configuration from the Periodic Table1 2 3 4 5 6 7 2A 3A 4A 5A 6A 7A Ne
3s2 P
3p3
P = [Ne]3s23p3
P has 5 valence electrons

49. Electron Configuration from the Periodic Table1 2 3 4 5 6 7 2A 3A 4A 5A 6A 7A
3d10 Ar As
4s2
4p3
As = [Ar]4s23d104p3
As has 5 valence electrons

50. 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

51. 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

52. Everyone Wants to Be Like a Noble Gas! Alkali Metals (Group 1A)
have one more electron than the previous noble gas, [NG]ns1
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+

53. Everyone Wants to Be Like a Noble Gas! Halogens (Group 7A)
one fewer electron than the next noble gas: [NG]ns2np5
Reactions with Metals: tend to gain an electron and attain the electron configuration of the next noble gas: [NG]ns2np5 + 1e  [NG]ns2np6
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

54. 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

55. 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

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

57. Trends in Atomic Size

58. Trends in Atomic Size Down a group: ___crease
valence shell farther from nucleus, weaker attraction
Across a period (left to right): ___crease
More protons to attract valence shell electrons
Electrons added to same valence shell
valence shell held closer

59. Metallic Character Metals Nonmetals 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

60. Trends in Metallic Character

61. 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

62. Be (4p+ & 4e-) Mg (12p+ & 12e-) Ca (20p+ & 20e-) Example: Group IIA

63. 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

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

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

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