1. Chapter 5 Notes Electron Models

2. Evolution of Electron Models The first model of the electron was given by J.J. Thompson—the electron’s discoverer. His was the “plum pudding” model.

3. The Rutherford Model With Rutherford’s discovery of the nucleus of an atom, the atomic model changed.

4. The Bohr Model Niels Bohr introduced his model, which answered why e- do not fall into the nucleus. e- orbit the nucleus in energy levels similar to the way the planets orbit the sun.

5. Bohr Model and Energy Levels energy levels are like the rungs of a ladder— you cannot be between rungs, just like an e- cannot be between energy levels. A quantum of energy is the amount of energy it takes to move from one energy level to the next.

6. Quantum Mechanical Model In 1926, Erwin Schrodinger used the new quantum theory to write and solve mathematical equations to describe e- location.

7. The Quantum Mechanical Model, cont. Previous models were based on physical models of the motion of large objects. This model does not predict the path of e-, but estimates the probability of finding an e- in a certain position.

8. Where are the electrons? In an atom, principal energy levels (n) can hold e-. These principal energy levels are assigned values in order of increasing energy (n=1,2,3,4...). Within each principal energy level, e- occupy energy sublevels. There are as many sublevels as the number of the energy level (i.e., level 1 has 1 sublevel, level 2 has 2 sublevels, etc.)

9. Where are the electrons? There are 4 types of sublevels (s,p,d and f). Inside the sublevel are atomic orbitals that hold the e-. Every atomic orbital can hold 2 e-. Principle Energy Levels: 1, 2, 3, 4... Sublevels: s, p, d, f Orbitals: 1 for s, 3 for p, 5 for d, 7 for f. Electrons: 2 for each orbital.

10. Orbital Shapes s orbital = s sublevel (spherical) + + p x orbitalp y orbital p z orbital = p sublevel (peanut)

11. D orbitals – fyi only (daisy)

12. Where are the electrons? So how many electrons can each energy level hold? –Level 1 has an s sublevel=2 e - –Level 2 has an s and a p sublevel=8e - –Level 3 has an s, p and d sublevel=18e - –Level 4 has an s, p, d and f sublevel=32e -

13. Electron Configuration

14. In the atom, electrons and the nucleus interact to make the most stable arrangement possible. The ways that electrons are arranged around the nucleus of an atom is called the electron configuration.

15. Aufbau Principle Electrons occupy orbitals of the lowest energy first. Pauli’s Exclusion Principle An atomic orbital may have no more than two electrons. Hund’s Rule When electrons occupy the same kind of sublevel, one electron goes in each orbital before the second one goes in.

16. Pauli Exclusion Principle –Each orbital can hold only TWO electrons with opposite spins.

17. Aufbau Principle “Sports Spectator Rule” (fill the lower stands first) Electrons fill the lowest energy orbitals first.

18. RIGHT WRONG Hund’s Rule “Empty Bus Seat Rule” Within a sublevel, place one electron per orbital before placing a second electron.

19. Heisenberg Uncertainty Principle Heisenberg concluded that it is impossible to make any measurement on an object with out disturbing the object (at least a little). The principle states: “It is fundamentally impossible to know precisely both the velocity and the position of a particle at the same time.”

20. O 8e - »Orbital Diagram Electron Configuration 1s 2 2s 2 2p 4 B. Notation 1s 2s 2p

21. Abbreviated S 16e - S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 Unabbreviated

22. © 1998 by Harcourt Brace & Company s p d (n-1) f (n-2) 12345671234567 6767 C. Periodic Patterns

23. s-block1st Period 1s 1 1st column of s-block C. Periodic Patterns Example - Hydrogen

24. C. Periodic Patterns Shorthand Configuration –Core e - : Go up one row and over to the Noble Gas. –Valence e - : On the next row, fill in the # of e - in each sublevel.

25. [Ar]4s 2 3d 10 4p 2 C. Periodic Patterns Example - Germanium

26. Electron Configuration Exceptions –Copper EXPECT :[Ar] 4s 2 3d 9 ACTUALLY :[Ar] 4s 1 3d 10 D. Stability Cr - [Ar]4s1 3d5 Cu - [Ar]4s1 3d10

27. Valence Electrons Electrons in the atom’s outermost orbitals. These are the electrons that determine the atom’s chemical properties. The fewer valence electrons an atom holds, the less stable it becomes and the more likely it is to react.

28. Octets When an atom has 8 electrons in its largest energy level it is said to have an octet. This is the most stable and least reactive atom.

29. Lewis Dot Structures An atom’s dot structure consists of the element’s symbol, (represents the atomic nucleus and inner-level electrons), surrounded by dots representing the atom’s valence electrons. G. N. Lewis devised this method while teaching a college chemistry class in 1902. Example: Lithium Atomic Number: 3 Electron Configuration: 1s 2 2s 2 Electron Dot Structure: Li·

30. Practice Carbon: Fluorine: Neon: