1. CH. 2 atomic models electronic configuration oxidation numbers
The Elements
CH. 2
atomic models
electronic configuration
oxidation numbers

2. The Models of the Atom

3. Line Spectra 22

4. What is causing the light to appear when the atom is excited by electricity?
Why does the line spectrum of hydrogen have very distinct lines of color and not the complete rainbow, which includes all colors?

5. The Atomic Energy and the Bohr Model of the Atom25

6. The Quantum Theory
The main idea that comes from the Quantum Mechanical Theory and the Schrodinger Equation is Quantum numbers that define the surfaces of the electron densities in which the electrons reside.
There are 3 main quantum numbers, n, l and ml, and a fourth quantum number, ms

7. Principle quantum number, n
Schrödinger’s equation provides a number of possible functions that define the shape of the electron cloud about the nucleus, these shapes are the orbitals. The functions can be arranged in sets and subsets. Orbitals are classified in three ways:
An electron shell is a collection of orbitals, identified by an integer, n, which ranges in value infinitely up from 1. Electrons with the same value are in the same shell. In chemistry, important shells have n values between 1 and 7.

8. Angular Momentum, l
Subshells are groups of orbitals within an electron shell. They are identified by the letters s, p, d, f and through the alphabet beginning with g. In chemical systems, only s, p, d, and f subshells are important. The number of subshells within an electron shell is equal to the n value of the shell.

9. Magnetic, ml
Individual orbitals are identified by their direction in space using Cartesian coordinates. The number of orbitals within a subshell depends on the subshell type: s subshells have a single orbital, p subshells have 3 orbitals, d subshells 5 and f subshells 7 orbitals.

10. Values of n, l and ml

11. Example

12. The orbital shape and l

14. The Energy Levels of the orbitals
In hydrogen the energy of the orbitals and the order in which they fill are dependant only upon their principle quantum number; however, in atoms with more electrons the order changes.

15. Energy levels in atoms of more than one electron.
Notice the shifting of the energy levels, as n increases, energy increases, as l increases, energy increases, but mixing of the energy levels still occur.

16. Orbitals Increasing Energy 10s 9s 9p 9d 9f 9g 9h 9i 9j 9k

17. Electrons fill different orbitals in a subshell until the subshell is half - filled. (one in each orbital before pairing begins)
If an atom has a filled valence shell, the atom is more stable.
Each electron or the space it occupies can be defined by its four quantum numbers (n, l, ml, ms)
No two electrons can occupy exactly the same space or have the same four quantum numbers.
Some electron configurations defy expectations created by the periodic table, particularly for heavier elements.

18. Periodic Blocks

19. Why do we have those rows at the bottom?Rb H Li Na Cs K La Ba Fr Be Mg Sr Ca Y Ac Ra Sc Tl Hg Au Hf Pt Ir Os Re W Ta He Rn At Po Bi Pb Cd Ag Zr Pd Rh Ru Tc Mo Nb Zn Cu Ti Ni Co Fe Mn Cr V In Xe I Te Sb Sn Ga Kr Br Se As Ge Al Ar Cl S P Si B Ne F O N C Gd Cm Tb Bk Sm Pu Eu Am Nd U Pm Np Ce Th Pr Pa Yb No Lu Lr Er Fm Tm Md Dy Cf Ho Es
This arrangement takes too
much space and is hard to read.

20. Hund’s Rule
A full energy level is very stable, of lowest energy, a half filled shell is almost as stable as a fully filled shell and a partially filled shell is the least stable of highest energy. The orbitals will go from partially filled to half filled by putting one electron in each orbital before pairing the spins.

21. Electron Configurations
Complete electron configurations
noble gas electron configurations showing the noble gas and the remaining valence electrons
valence electron configuration

24. Fe 4s 3d
Fe+2 4s 3d
Fe+3 4s 3d

25. Rules for assigning oxidation numbers
The sum of the oxidation number of all atoms must equal the net charge of the species.
In compounds:
Group IA are +1.
Group IIA are +2.
B and Al are +3, and F is -1.
Hydrogen is +1 except when combined with a metal. Then it is -1.
Oxygen is -2 except for peroxides and superoxides.
Elements in their elemental state have an oxidation number of zero.

26. Common oxidation numbersLi +1 Na Cs Rb K Fr Ba +2 Be Mg Sr Ca Ra B +3 C +4 -2 -4 N
+5 +4
+3 +2
+1 -3 O -1 -2 F -1 Ne H +1 He Al +3 Si +4 -4 P +5 +3 -3 S
+6 +4 +2 -2 Cl
+7 +5
+3 +1 -1 Ar Sc 3+ Ti +4 +3 +2 V
+5 +4 +3 +2 Cr +6 +3 +2 Mn
+7 +6
+4 +3 +2 Fe +3 +2 Co +3 +2 Ni +2 Cu +2 +1 Zn +2 Ga +3 Ge +4 -4 As 5+ 3+ 3- Se 6+ 4+ 2- Br +5 +1 -1 Kr +4 +2 Y +3 Zr +4 Nb +5 +4 Mo +6 +4 +3 Tc +7 +6 +4 Ru
+8 +6 +4 +3 Rh +4 +3 +2 Pd +4 +2 Ag +1 Cd +2 In +3 Sn +4 +2 Sb +5 +3 -3 Te +6 +4 -2 I
+7 +5 +1 -1 Xe +6 +4 +2 Lu +3 Hf +4 Ta +5 W +6 +4 Re +7 +6 +4 Os +8 +6 Ir +4 +3 Pt +4 +2 Au +3 +1 Hg +2 +1 Tl +3 +1 Pb +4 +2 Bi +5 +3 Po +2 At -1 Rn Lr +3
Common oxidation numbers

27. Oxidation numbers and the periodic table
Some observed trends in compounds.
Metals have positive oxidation numbers.
Transition metals typically have more than one oxidation number.
Nonmetals and semimetals have both positive and negative oxidation numbers.
No element exists in a compound with an oxidation number greater than +8.
The most negative oxidation numbers equals 8 - the group number

28. Electron Configuration Activities
Simple Atomic Structure
Atomic Structure - Electron Configuration
Structured Curriculum Lesson Plan