1. Unit: Periodicity
(Chapter 7)

2. Atomic Models Rutherford
Based on gold foil experiment (alpha, α, 4He+2, particles shot at very thin-2000 atoms wide- gold foil. Most particles went through, few--less than 1/8000– were deflected or reflected. Indicating mostly empty space. Allowed for calculations to determine relative size of nucleus compared to atomic radius)
Identified nucleus
Dense, massive, positive, small
Atom is mostly empty space
With electrons dispersed around nucleus

3. F α q1 q2 r2 Ultimately, Rutherford’s model is based on
Attraction between oppositely charged protons and electrons
Governed by Coulombs Law, which states
Force, F, between two charged particles, q1 and q2, is inversely proportional to the square of the distances between them
When q1 and q2 have the same sign, the forces is repulsive
When they are opposite signs, the force is attractive
i.e. like charges repel and opposite charges attract.
F α
q1 q2 r2

4. Bohr
Suggested electrons could only travel in fixed orbits or shells around the nucleus
Based upon the following observations
Gaseous elements emit electromagnetic radiation when energized.
Light emitted consists of discrete packets of energy (quanta)
Each element a unique pattern of radiation (line emission spectra)
Based on release of radiation caused by absorption of energy by electrons which allowed them to be promoted to shell further from nucleus (ex: from “orbit 1 to orbit 2”)

5. Electron falls back to original, lowest energy shell (ground state), it releases energy previously absorbed (excited state)
Since shells are in fixed positions, differences in energy between them (hence the wavelength of radiation-as seen in line emission spectra) is also fixed.
Gives a unique and identified pattern for each element.
Lyman series (ultraviolet)
Balmar series (visible)
Paschen (infrared)
Others: lower frequency with much overlap (difficult to distinguish)
Question: what will be the difference in energies for each of the series?

6. Ionization energy
Electron may gain enough energy to completely overcome the force of attraction exerted by nucleus
Will get ejected/removed from atom
Energy required to remove electron from atom
Different for each element
Since all electrons sit at different distances from nucleus
Force of attraction is partly dependent on distance between nucleus and electron (closer electron = greater attraction = higher energy required to remove electron)
Described as 1st, 2nd, 3rd … ionization energies.
Consistent with Bohr model of electrons being at certain regular arrangements.

7. Wave-particle behavior of electrons
Although Bohr’s fixed “orbits” in which electrons traveled helped explain some behavior, it was insufficient in explaining other observed properties of atoms and electrons.
Scientists postulated that electrons (and other very small particles), like photons display dual wave-particle behavior.
Wave behavior
Wavelength (lambda: λ)-distance between two repeating points on a sine wave. Units: meters
Frequency (υ)-number of waves that pass a fixed point in s second. Units: sec-1 or Hertz (Hz)
Speed of light (c) – speed of electromagnetic radiation. Constant in vacuum X 108 m/sec
Speed of light = frequency X wavelength c = λ υ
Inversely proportionate-as one increases, the other decreases
Since electrons have mass, de Broglie’s equation can be used to determine λ: λ =
h m υ

8. Value of frequency and value of wavelength helps determine type of electromagnetic radiation.
Also allows for calculations to determine energy associated with type of radiation
E =
h c λ
E = h υ OR
Where h = Planck’s constant, 6.63 X J·sec

10. Quantum Model
Schrodinger—developed and solved wave equations that treat electrons as photons
Makes predictions about where an electron is likely to be found in a atom
Heisenberg Uncertainty Principle
Momentum and position of electron cannot be simultaneously determined.
It is only possible to predict where an electron will probably be at any one time rather than knowing its exact whereabouts
Helped develop quantum mechanical model
Three dimensional probability map (orbitals) where electrons are likely to be found

11. Quantum Numbers Principle quantum number n
Main energy level = shells (as per Bohr) = periods on periodic table
Value = 1-7
Each subsequent shell is
farther from nucleus
More energetic
Has increasing number of sublevels

12. Angular momentum quantum number
Identifies shape of orbitals
Identifies sub-shell
Increasing number of subshells available with increasing main energy level
Possible values:
s (spherical)
p (dumb bell / infinity)
d (4 leaf clover)
f (8 petals)
Together with main energy level, forms the sublevels
1s, 2p, 4f etc.

14. Each sub-shell is further divided into specific number of orbitals
Each orbital holds a MAXIMUM of 2 electrons
therefore, helps determine maximum number of electrons allowed per main energy level or shell, using formula 2n2

15. Summary of shells, orbitals and electrons

16. Rules defining electron properties
Pauli exclusion principle
All electrons in any single atom must be unique (defined by unique set of four quantum numbers)
For pairs of electrons in the SAME orbital, the shell, subshell and orbital are the same, they must be distinguished another method
Intrinsic property known as “spin” (spin quantum number)
2 electrons in same orbital must have opposite spins, often denoted by pair of arrows (one up, one down) or quantum numbers +1/2 and -1/2
For electrons in same subshell but different orbitals, unique values are identified by the alignment of orbitals of same type and energy along x, y and z axes in space (for p type. d and f orbitals have more complicated shapes and axes orientation)

17. Rules for filling orbitals
Aufbau’s process
Determine number of electrons to be placed (atomic number. + or – electrons for ions)
Lowest energy orbitals are filled first
Orbitals have ascending energies with 1s having the smallest energy, 2s, the next smallest etc.
Complications arise for larger atoms with electrons in d and f sublevels.
Generally can be ignored. Recognize that actual electron configurations do not always follow predicted configurations

18. Hund’s Rule Maximum multiplicity
If multiple orbitals of same energy (ex: 3- 4p orbitals or 5 – 3d orbitals), (described as degenerate orbitals) are available/occupied, 1 electron in placed in each orbital before pairing takes place.
Maximum distance between repelling electrons of same energy.
As a result
Completely filled sublevels are most stable
Partially filled sublevels with 1 electron in each orbital are more stable than with
some empty orbitals
some orbitals with one and others with 2 electrons.

19. Periodic table and electronic configurations

20. Period number = shell number
When filling d orbitals, subtract one from period to determine correct shell
Block shows type of orbital
Add one electron for each element until orbital, then sub-shell and ultimately shell is full
Record electron configuration in the format:
Shell # (period), block (orbital), number of electrons (as a superscript)
Example: H : 1s1 He: 1s2 F: 1s2, 2s2, 2p5

21. Noble Gas Core Method Abbreviated method
Determined by writing in previous noble gas in square brackets, then filling orbitals as before.
Ex: phosphorous, atomic number 15: Previous Noble gas: Neon [Ne] 3s2, 3p5

22. Ions and electronic configuration
Gain or loss of electrons from/to outermost shell
Cations (positive ions) through loss of electrons (metals)
For d block elements, s electrons are lost first, then d electrons
Anions (negative ions) through gain of electrons (non-metals)
Magnitude of charge denotes number of electrons gained (-) or lost (+)
To determine electron configuration, simply add or remove necessary number of electrons to/from regular electron configuration.

23. Expanded electron configuration (orbital notation)
Shows ALL electrons and implies all 4 quantum numbers
Follows Pauli, Hund and Aufbau rules
Shows electrons in boxes

24. Application of Hund, Pauli and Aufbau.
Example: Nitrogen could be represented 4 ways
Showing it in orbital notation shows all three rules being applied appropriately

25. Photoelectron spectroscopy (PES)
Uses high energy UV or X-ray photons to energize electrons in an atom
Atoms can absorb sufficient energy for electrons to overcome attraction for nucleus and be ejected from the atom
Known as the photoelectric effect.
Analysis of ejected electrons by PES is used to gather data about specific electronic structure of an atom.
Photons used to eject electrons have energies corresponding to E = h υ
These energies EXCEED the ionization energy. Result:
Electrons overcome nuclear attraction, AND
Gain kinetic energy h υ = IE + KE
Analysis of the KE of the ejected electrons produces a PES spectrum
Gives electronic structure information

26. Interpreting Photoelectron Spectra
Consider the simulated PES plot for Mg (below)
CAREFULLY note—x-axis is ploted “BACKWARDS”
i.e. decreasing energy from left to right
Also note: scale is NON-LINEAR

27. X-axis: Energy associated with EACH SUB-shell. Units MJ/mol
Y-axis: Relative number of electrons
Larger (higher) peaks = greater number of electrons
Look at relative ratio of heights
Coulombs law-predicts that electrons FARTHEST from nucleus are EASIEST to remove
require least amount of energy
Example above (Mg) shows 2 electrons with very low energies (0.74), 8 (6 + 2) with intermediate energies (between 5.31 and 9.07) and 2 with VERY high energies (126)
Corresponds to Mg electronic configuration (in reverse)
3s2, 2p6, 2s2, 1s2
Think of EACH peak as subsequent sub-levels.

28. Breakdown 3s
Farthest, most energetic electrons, least amount of attraction to nucleus. Easiest (least energy required) to remove 2p
Closer, less energetic electrons, greater attraction to nucleus. Requires more energy to remove. 2s
VERY similar to 2p. Same shell, similar energies, but slightly lower energy = slightly greater attraction to nucleus. Slightly higher energy to remove 1s
Closest, LEAST energetic electrons, GREATEST amount of attraction to nucleus. MOST difficult (most energy required) to remove.
The “reversed” x-axis allows the electronic configuration of Mg to be read from left to right (normally).
***NOTE: Don’t ASSUME reversed x-axis. Read carefully (check to see whether energy is increasing or decreasing)

30. PES for mixtures of elements
Consider the PES for a mixture of C, N, O and F.
ALL have atomic numbers greater than 2
Each 1s peak (highlighted in pink) must be produced by 2 electrons
So WHY are all 1s peaks not the same height?
Elements present in the mixture in varying abundances. Therefore TOTAL 1s electrons vary
O most abundant
F least abundant

31. Let’s practice Identify the element Charge on ion?
Formula for compound formed with Calcium?

32. Identify the element
Charge on ion?
Formula for compound formed with chlorine?