1. Electrons Configurations Cartoon courtesy of NearingZero.net

2. Wave-Particle Duality JJ Thomson won the Nobel prize for describing the electron as a particle. His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron. The electron is a particle! The electron is an energy wave!

3. The Wave-like Electron Louis deBroglie The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.

4. …produces all of the colors in a continuous spectrum Spectroscopic analysis of the visible spectrum…

5. This produces bands of light with definite wavelengths. Electron transitions involve jumps of definite amounts of energy.

6. Quantum Numbers Each electron in an atom has a unique set of 4 quantum numbers which describe it.  Principal quantum number  Angular momentum quantum number  Magnetic quantum number  Spin quantum number

7. Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli

8. Principal Quantum Number Generally symbolized by n, it denotes the shell (energy level) in which the electron is located. Number of electrons that can fit in a shell: 2n 2

9. Angular Momentum Quantum Number The angular momentum quantum number, generally symbolized by l, denotes the orbital (subshell) in which the electron is located.

10. Magnetic Quantum Number The magnetic quantum number, generally symbolized by m, denotes the orientation of the electron’s orbital with respect to the three axes in space.

11. Assigning the Numbers  The three quantum numbers (n, l, and m) are integers.  The principal quantum number (n) cannot be zero.  n must be 1, 2, 3, etc.  The angular momentum quantum number (l) can be any integer between 0 and n - 1.  For n = 3, l can be either 0, 1, or 2.  The magnetic quantum number (m) can be any integer between -l and +l.  For l = 2, m can be either -2, -1, 0, +1, or +2.

12. Spin Quantum Number Spin quantum number denotes the behavior (direction of spin) of an electron within a magnetic field. Possibilities for electron spin:

13. Uncertainty Principle  In 1927 by Werner Heisenberg (German theoretical physicist)  electrons can only be detected by their interaction with light  any attempt to locate a specific electron knocks the electron off course

14. Uncertainty Principle  Heisenberg Uncertainty Principle- it is impossible to know both the position and velocity of an electron  We can find the most likely position of an electron

15. Orbitals  each sublevel is broken into orbitals  each orbital can hold a maximum of 2 electrons  orbital- a 3D region around the nucleus that has a high probability of holding electrons

16. Orbital shapes are defined as the surface that contains 90% of the total electron probability. An orbital is a region within an atom where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…

17. Orbitals of the same shape (s, for instance) grow larger as n increases… Nodes are regions of low probability within an orbital. Sizes of s orbitals

18. The s orbital has a spherical shape centered around the origin of the three axes in space. s orbital shape

19. There are three dumbbell-shaped p orbitals in each energy level above n = 1, each assigned to its own axis (x, y and z) in space. P orbital shape

20. Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbells ” …and a “dumbell with a donut”!

21. Shape of f orbitals

22. Number of Orbitals sublevelmax # e-# orbitals s21 p63 d105 f147

23. Energy Number of Types of Level Sublevels Sublevels 11s 22s,p 33s,p,d 44s,p,d,f

24. Energy Level Sublevels#Orbitals Total Orbitals Total Electrons 1 2 3 4

25. Energy Level Sublevels#Orbitals Total Orbitals Total Electrons 1s112 2s148 p3 3s1918 p3 d5 4s11632 p3 d5 f7

26. Orbital filling table

27. Electron configuration of the elements of the first three series

28. Irregular confirmations of Cr and Cu Chromium steals a 4s electron to half fill its 3d sublevel Copper steals a 4s electron to FILL its 3d sublevel

29. Electron Configurations  the arrangement of electrons in an atom  each type of atom has a unique electron configuration  electrons tend to assume positions that create the lowest possible energy for atom  ground state electron configuration- lowest energy arrangement of electrons

30. Rules for Arrangements  Aufbau Principle- an electron occupies the lowest-energy orbital that can receive it  Beginning in the 3 rd energy level, the energies of the sublevels in different energy levels begin to overlap

31. Rules for Arrangements  Pauli Exclusion Principle- no two electrons in the same atom can have the same set of 4 quantum numbers  Hund’s Rule- orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second  all unpaired electrons must have the same spin

32. Rules for Arrangements

33. Writing Configurations  Orbital Notation: an orbital is written as a line an orbital is written as a line each orbital has a name written below it each orbital has a name written below it electrons are drawn as arrows (up and down) electrons are drawn as arrows (up and down)  Electron Configuration Notation number of electrons in sublevel is added as a superscript number of electrons in sublevel is added as a superscript

34. Order for Filling Sublevels

35. Writing Configurations  Start by finding the number of electrons in the atom  Identify the sublevel that the last electron added is in by looking at the location in periodic table  Draw out lines for each orbital beginning with 1s and ending with the sublevel identified  Add arrows individually to the orbitals until all electrons have been drawn

36. Silicon  number of electrons: 14  last electron is in sublevel: 3p 1s 2s 2p 3s 3p  Valence Electrons- the electrons in the outermost energy level

37. Chlorine  number of electrons: 17  last electron is in sublevel: 3p 2p 3s 3p 1s 2s

38. Sodium  number of electrons: 11  last electron is in sublevel: 3s 1s 2 2s 2 2p 6 3s 1 1s 2s 2p3s

39. Calcium  number of electrons: 20  last electron is in sublevel: 4s 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 1s 2s 2p3s 3p 4s

40. Bromine  number of electrons: 35  last electron is in sublevel: 4p 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 1s 2s 2p 3s3p 4s3d 4p 1s 2s 2p 3s3p 4s3d 4p

41. Argon  number of electrons: 18  last electron is in sublevel: 3p 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2s 2p 3s3p

42. Noble Gas Notation  short hand for larger atoms  configuration for the last noble gas is abbreviated by the noble gas’s symbol in brackets