1. Chapter 5: Electrons in the Atom
5.1 Models of the Atom
5.2 Electron Arrangement in Atoms

2. Bell Question How is the Bohr Model related to electrons and energy?
The Bohr model represented how electrons that jump energy levels when energy is absorbed and released
This is not true for all elements… We need a better model!

3. The Wave Mechanical Model of the Atom
Picture a firefly
Orbitals
Nothing like orbits
Probability of finding the electron within a certain space
Does not describe how an electron moves or when it will be in a given space.

4. Section 5.1 – Atomic orbitals

5. Atomic Orbitals
Atomic orbitals represent how the electrons are arranged on an atom.

6. The Hydrogen Orbitals Atomic orbitals do not have an exact edge.
We arbitrarily designate the orbital as the shape that contains the electron 90% of the time.
This spherical orbital represents the ground state of a hydrogen atom

7. In the Bohr model a higher energy level was just an orbit with a larger radius.
In the wave mechanical model, the higher energy states are different kinds of orbitals with different shapes.

8. Orbitals and Energy
Each principal energy level is divided into sublevels.
Labeled with numbers and letters
Indicate the shape of the orbital
ex. 1s

9. Each sublevel has a different number of orbitals
s contains 1 orbital
p contains 3 orbitals
d contains 5 orbitals
f contains 7 orbitals

10. The s and p types of orbitals

11. Section 5.2 – Electron arrangement in the atoms

12. Electron Configurations
Electron configurations are like the electron’s address
Indicates which sublevels and orbitals the electrons are most likely going to occupy.

13. Energy and Subshells 6p 5d 4f 6s 5p 4d 5s 4p
Subshells are filled from the lowest energy level to increasing energy levels. 3d 4s 3p 3s 2p 2s
Energy 1s

14. Aufbau Principle
Aufbau Principle: Electrons fill sublevels (and orbitals) so that the total energy of atom is the minimum
What does this mean?
Electrons must fill the lowest energy sublevels and orbitals before moving on to the next higher energy sublevels/orbital.

15. Hund’s Rule
Hund’s Rule: Place electrons in unoccupied orbitals of the same energy level before doubling up.
How does this work?
If you need to add 3 electrons to a p sublevel, add 1 to each before beginning to double up.

16. Pauli Exclusion Principle
Pauli Exclusion Principle: Two electrons that occupy the same orbital must have different spins.
“Spin” describes the angular momentum of the electron
“Spin” is designated with an up or down arrow.
How does this work?
If you need to add 4 electrons to a p subshell, you’ll need to double up. When you double up, make them opposite spins.

17. Orbital Diagrams
Orbital diagrams use boxes for orbitals and arrows for electrons.
Pay close attention to the number of orbitals for each sublevel. This will determine the number of boxes that need to be drawn

18. Example
Nitrogen
1s2 2s2 2p3 2p 2s 1s

19. Practice
Bromine
Sodium
Scandium

20. Electron Configurations
Electron configurations describe where the electrons are located in the atom by using the number of the principal energy level and the letter of the sublevel.
1s2
Principal energy level
Orbital (sublevel)
Number of electrons in that sublevel

21. 1s2 2s2 2p6 ……

22. Diagonal Rule
Follow the arrows to know the order of each sublevel

23. Periodic Table

24. Practice
Calcium
Bromide ion
Sodium ion

25. Noble Gases and Noble Gas Notation
Noble Gas – Group 8 of the Periodic Table. They contain full valence shells.
Noble Gas Notation – Noble gas is used to represent the core (inner) electrons
Spectroscopic 1s 2s 2p 3s 3p 2 6 4s 3d 10 4p 5 Br
Noble gas
[Ar] 4s 2 3d 10 4p 5
The “[Ar]” represents the core electrons and only the remaining electrons are shown

26. Writing Noble Gas Notation
Find the element on the periodic table
Find the nearest noble gas that has less electrons than the element.
Write the noble gas in [ ].
Write the valence electrons for the element.