1. Ch. 5: Electrons in the Atom

2. Development of atomic models

3. Rutherford’s “Nuclear Atom” Model
Rutherford’s atomic model could not explain the chemical properties of elements
Why do objects change color when heated?
Example: heating and iron horseshoe
The observed behavior could only be explained if the atoms gave off light in specific amounts of energy

4. Bohr’s Model
Planetary Model
Bohr proposed that an electron is found in specific circular paths, or orbits, around the nucleus
Each possible electron orbit in Bohr’s model has a fixed energy.
The fixed energies an electron can have are called energy levels
A quantum of energy is the amount of energy required to move an electron from one energy level to another

5. Like the rungs of the strange ladder, the energy levels in an atom are not equally spaced.
The higher the energy level occupied by an electron, the less energy it takes to move from that energy level to the next higher energy level.

6. Quantum Mechanical Model
Erwin Schrödinger- Devised and solved a mathematical equation describing the behavior of the electron in a hydrogen atom
The modern description of the electrons in atoms, the quantum mechanical model, comes from the mathematical solutions to the Schrödinger equation.

7. Quantum Mechanical Model
The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.
Limits electrons to certain energy levels
Does not attempt to predict the path of the electron

8. Describing the location of an electron
Energy Levels:
Principle quantum number (n)
Can be values 1-7
Defines the size
As n increases, the energy level gets larger
Sublevels
Energy levels broken down (s, p, d, f, g, h, i)
Atomic Orbital:
3D region around the nucleus that describes the electrons probable location
Shape of the sublevel
Spin

9. Energy Sublevel
Each energy level has “n” number of sublevels (Level 1 has 1 sublevel, level 7 has 7 sublevels)
The sublevels have labels
1st sublevel on in each level …s
2nd…p
3rd…d
4th…f

10. Atomic Orbitals
Atomic Orbital: region of space in which there is a high probability of finding an electron
Each energy level has n2 number of atomic orbitals
(Energy level 2 has 22 atomic orbitals = 4)
Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found.

11. Orbitals Each sublevel has a fixed number of orbitals s…1 orbital
p…3 orbitals
d…5 orbitals
f…7 orbitals

12. Different atomic orbitals are denoted by letters
Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped.

13. Four of the five d orbitals have the same shape but different orientations in space.

15. The numbers and kinds of atomic orbitals depend on the energy sublevel.

16. Atomic Spins Each orbital can hold a maximum of 2 electrons
Each energy level has 2n2 number of electrons
The two electrons travel with opposite spins

17. The number of electrons allowed in each of the first four energy levels are shown here.

18. Electron Configuration
The arrangement of the electrons in an atom
The lowest energy is the most stable
There are 3 main rules for writing the electron configurations of elements

19. Aufbau Principle Aufbau Principle
Each electron occupies the lowest energy orbit
Each sublevel has a different energy state
e- within an energy level fill in the sub level order…s, p, d, then f
The energy levels overlap so a guideline is needed to establish sublevel order
Diagonal Rule: Sets the order of filling the sublevels

20. Pauli Exclusion Principle/Hund’s Rule
An atomic orbital contains a maximum of two electrons
The two electrons will travel with opposite spins
The direction of the spin will be represented
Each arrow represents an electron (Together is one pair of electrons)
Hund’s Rule
e- will individually occupy equal energy orbitals before forming a pair
All orbitals of a sublevel are of equal energy
Want the number of electrons with the same spin direction as large as possible

21. Three Methods of Notation
Orbital Notation: shows every electron with an arrow
Example: O-16 (8e-)

22. Practice – Orbital Notation
Argon
Sulfur

23. Three Methods of Notation
2) Electron Configuration: shows the total number of electrons in each sublevel as a superscript
Example: Argon ( 18 e-)
1s2 2s2 2p6 3s2 3p6

24. Three Methods of Notation
3) Electron Dot (Lewis Dot): shows each outer level electron as a dot **Maximum number of 8 dots**

25. Valence Electrons
Valence Electrons: Electrons in the highest number energy level
Found in the highest number s and p sublevels
The electrons used in electron dot notation
Very Important Concept in Chemistry:
- Atoms in the same group (column) on the periodic table have similar chemical properties because they have the same number of valence electrons

26. Notice that the electron dot structures repeat as you move down the table:
1A: 1 dot 2A: 2 dots 3A: 3 dots 4A: 4 dots **Middle Section has 1- 2 VE
5A: 5 dots
6A: 6 dots
7A: 7 dots
8A: 8 dots
**He: 2 dots

28. Exceptions to Electron Configuration
Some electron configurations for some elements differ from those assigned using the aufbau principle.
Because the half-filled sublevels are not as stable as filled sublevel, but they are more stable than other configurations
Examples: Copper, Chromium, Silver, Molybdenum, & Gold

29. 3 3 3 3

30. Levels/Sublevels on Periodic Table
S block- first 2 columns (and He)
D block- next 10 columns
P block- next 6 columns
F block- 14 columns below table

31. Extended Periodic Table

32. Noble Gas Notation:
Noble Gas Notation: a shorter version of electron configuration
Since the inner level electron configuration doesn’t change the noble gas is used as a shortcut
Find the noble gas before it (column 8). Write the noble gas in bracket, and then fill in the remaining configuration
Example: Al - 1s2 2s2 2p6 3s2 3p1
Ne- 1s2 2s2 2p6
Shorter Version: Al- [Ne] 3s2 3p1
Same

33. Noble Gas Notation Practice
Antimony:
Calcium: [Ne] 3s2 3p6 4s2 -- fix it