1. Chapter 4 Electrons in Atoms Section 4.2

2. Development of the Atom  Originally described as the smallest particles of matter  Discoveries of electrons, protons, and neutrons destroyed the idea of an indivisible atom

3. Development of the Atom  Where are the electrons?  Niels Bohr developed a model for the hydrogen atom  Electrons are in definite fixed orbits Energy levels

4. Development of the Atom The amount of energy the electron possess depends on its distance from the nucleus. The further away from the nucleus the more energy the electron possess.

5. Development of the Atom  Bohr’s Model worked perfectly for hydrogen but failed with any atom that had more than one electron  Back to the drawing board…

6. Schrodinger Wave Equation  Electrons have dual wave-particle nature  Developed an equation that treated electrons like waves  Only waves of specific energies, and therefore frequencies, provided solutions to the equation.

7. Heisenberg Uncertainty Principle  There is a fundamental limitation to how precisely both the position and the momentum of an electron can be known.  Together with Schrodinger’s wave equation a new model of the atom was developed.

8. The Quantum Mechanical Model  Quantum Mechanics describes the motion of small particles confined to a tiny region of space.  The exact position of an electron at any given instant is not specified  The exact path that the electron takes can not be specified

9.  Its all about probability!  The electron is found inside a blurry “electron cloud”  An area where there is a chance of finding an electron. The Quantum Mechanical Model

10.  The electron does not travel around the nucleus in neat orbits of fixed energy like Bohr proposed.  The exist in certain regions called orbitals.  An orbital is a 3-D region around the nucleus that indicates the probable location of an electron.

11. Atomic Orbitals  Atomic orbitals have different shapes and sizes.  More to come!

12. Atomic Orbitals and Quantum Numbers  According to Bohr, electrons of increasing energy occupy orbits farther and farther from the nucleus  Schrodinger’s equation also accounts for quantized energies for electrons

13. Atomic Orbitals and Quantum Numbers  Electrons energy level is not the only characteristic of an orbital that is indicated by solving Schrodinger’s Eq.  Quantum numbers specify the properties of atomic orbitals and the properties of electrons in orbitals

14. Atomic Orbitals and Quantum Numbers  The first three quantum numbers result from solutions to Schrodinger’s equation.  They indicate the main energy level, the shape and orientation of an orbital.  The fourth indicates the spin direction of an electron.

15. Principal Quantum Number  indicates the main energy level occupied by the electron.  Positive integers 1,2,3,…  Symbol n  Larger n = more energy

16. Principal Quantum Numbers  Within each energy level the complex math of Schrödinger's equation describes several shapes.  These are called atomic orbitals  regions where there is a high probability of finding an electron.  The total number of orbital’s that exist in a main energy level is equal to n 2

17. Angular Momentum Quantum Number  indicates the shape of the orbital.  Symbol l  Except at E1, n = 1, orbitals of different shapes (sublevels) exist for a given value of n.

18. Angular Momentum Quantum Number  The number of orbital shapes possible is equal to n.  l = 0, 1, 2, … n-1 (all positive integers)

19. Shapes of Orbitals  n = 1 l = 0 one orbital s  n = 2 l = 0, 1 two orbitals s, p  n = 3 l = 0, 1, 2 three orbitals s, p, d

20.  One s orbital for every energy level  Spherical shaped  Each s orbital can hold 2 electrons  Called the 1s, 2s, 3s, etc.. orbitals. S orbitals

21. P orbitals  Start at the second energy level  3 different directions  3 different shapes (dumbell)  Each can hold 2 electrons

22. P Orbitals All three p-orbitals

23. D orbitals  Start at the third energy level  5 different shapes  Each can hold 2 electrons

24. F orbitals  Start at the fourth energy level  Have seven different shapes  2 electrons per shape

25. F orbitals

26. Summary s p d f # of shapes Max electrons Starts at energy level 121 362 5103 7144

27. Magnetic Quantum Number  the orientation of the orbital in 3-D space. (x, y, z)  symbol m l  the values of m l range from – l to + l  Ex: n=1 l=0 m l =0 n=2l=0, 1m l = -1,0, 1 n=3 l= 0, 1, 2m l = ?

28. Spin Quantum Number  electrons are not stationary particles, they spin  symbol m s  they can only spin in two directions, clockwise and counterclockwise (designations we have assigned them)  the values of m s are +1/2 or – 1/2

29. The Address of an Electron  No two electrons have the same 4 quantum numbers.  what I know from the quantum numbers of an electron:  Ex:  1, 0, 0, +1/ 2

30. The Address of an Electron  n = 1  first principal energy level,  l = 0  s orbital,  m l = 0  encompasses all axis, x, y,& z  m s = + ½  spinning clockwise  Try this one  3, 1, -1, -1/2

31. 3, 1, -1, -1/2  n = 3  third principal energy level,  l = 1  one of the p orbitals,  m l = -1  specifically the p x orbital  m s = - ½  spinning counterclockwise