1. Unit 3 The Quantum Model

2. The Quantum Mechanical Model
Erwin Schrodinger ( ) In 1926 Schrodenger proposed our modern model for the atom, the quantum
mechanical model, which
is based mathematical
calculations.

3. The Quantum Mechanical Model
Like the Bohr model, electrons are restricted to certain energy levels.
Unlike the Bohr model, electrons do NOT have to take an exact path around the nucleus.

4. The Quantum Mechanical Model
The importance of the quantum mechanical model is that it determines:
allowed energies an electron can have
how likely it is to find the electron in a particular location

5. The Quantum Mechanical Model
-We cannot know both the location and velocity of an electron (Heisenberg’s uncertainty principle), thus Schrodinger’s equation does not tell us the exact location of the electron, rather it describes the probability that an electron will be at a certain location in the atom.

6. The Quantum Mechanical Model

7. Because it behaves as a particle, a baseball follows a well-defined path as it travels from the pitcher to the catcher. Because of their wave nature, an electron's path cannot be precisely known. The best we can do is to calculate the probability of the electron following a specific path.

8. If the baseball displayed wave-particle duality, the path of the baseball could not be precisely determined. The best we could do would be to make a probability map of where a "pitched" electron will cross home plate.

9. The Quantum Mechanical Model
Today we say that the electrons are located in a region outside the nucleus called the electron cloud.
The nucleus is not a little “planet”- rather the way we look at atoms is now like a probability map. Where the cloud makes up regions where electrons may be found.

10. A. Quantum
To move from one energy level to another, an electron must gain or lose just the right amount of energy.
The exact amount of energy required to move from one energy level to another is called a quantum of energy.

11. IV. Quantum Mechanics Model of the Atom and Quantum Numbers
Unlike the rungs of a ladder, the energy levels in an atom are NOT equally spaced. In fact, the energy levels become more closely spaced the farther they are from the nucleus.

12. B. Photon
An electron falling from a higher energy level to a lower energy level gives off an exact amount of light (called a photon).

13. A. Energy Levels
energy levels (represented by the letter n; also called the principal quantum number) are assigned values in order of increasing energy: n=1,2,3,4, and so forth.
Which energy level is furthest
away from the nucleus and has
electrons with the highest energy
- 1,2,3, or 4? _______ 4

14. B. Sublevels
Within each energy level, the electrons are located in various sublevels – there are 4 different sublevels s, p, d, and f.
s, p, d, and f describe the shapes of the orbitals.)
Remember we’re talking about what the probability map looks like!

15. C. Orbitals Orbitals are regions of the electron cloud.
Each orbital can only hold 2 electrons at a time.
For example, within the s sublevel there is 1 orbital (which is spherical); it is called the s orbital.

16. Each s ortibal holds 2 electrons
1 orbital - 2 electrons

17. 3 orbitals – 6 total electrons
Within the p sublevel there are 3 orbitals (which are dumbbell shaped) called the p orbitals.
3 orbitals – 6 total electrons

18. Within the d sublevel there are 5 orbitals (4 of which are clover leaf shaped) called the d orbitals. 5 orbitals -10 total electrons

19. Within the f sublevel there are 7 orbitals (that are flower-shaped) called the f orbitals orbitals - 14 total electrons

20. The electron cloud

22. s p d f Sublevel Number of orbitals this sublevel contains
Shape of orbitals
Total number of electron this sublevel can hold s p d f

23. Electron Arrangement in Atoms
Electron Configurations: the way in which electrons are arranged in various orbitals around the nucleus of the atom

24. Three rules of electron configurations
1. Aufbau Principle – electrons fill in order from lowest to highest energy.

26. Three rules of electron configurations
2. The Pauli exclusion principle
– An orbital can only hold a maximum of two electrons. Two electrons in the same orbital must have opposite spins.

28. Three rules of electron configurations
3. Hund’s rule – the lowest energy configuration for an atom is the one having the maximum number of unpaired electrons for a set of degenerate orbitals. By convention, all unpaired electrons are represented as having parallel spins with spin “up”.

29. 1. Orbital Notation Example: We Do: Potassium You Do: Silicon
Hydrogen ____ Helium _ Lithium ___ ___
1s s s 2s
We Do: Potassium
You Do: Silicon

30. Remember to fill each orbital in the correct order

31. 2. Electron Configuration Notation
The number of electrons in a sublevel is shown by adding a superscript to the sublevel designation. The superscript indicates the number of electrons present in that sublevel.
Example:
We Do: Cobalt ___________________________
You Do: Arsenic__________________________

32. 3. Noble Gas Notation (Shorthand Notation)
In order to write less, a noble gas can stand for that element’s number of electrons
Example: Sodium [Ne] 3s1
The [Ne] stands for the first 10 electrons
because neon has 10 electrons = 11 electrons
We Do: Cobalt ________________________
You Do: Arsenic _______________________

33. Elements in the same block of the periodic table have their final electrons in the same sublevel.