Electron Configuration

1. Electron Configuration Topic 2.2 Scandium 3-D video (2:31) MIT 5:15

2. Review

3. Jumping Electrons normally electrons exist in the ground state, meaning they are as close to the nucleus as possible when an electron is excited by adding energy (heat/electrity) to an atom, the electron will absorb energy and "jump" to a higher energy level

4. after a short time, this electron will spontaneously fall back to a lower energy level, giving off a quantum of energy • the key to Bohr's theory was the fact that the electron could only "jump" and "fall" to precise energy levels, thus emitting a limited spectrum of light. • a quantum is a discrete packet of energy or matter • the minimum value of a physical property involved in an interaction • plural of quantum is quanta

8. Quantum Numbers (however, actual numbers are often not used) • each electron in an atom is described by four different quantum numbers • think of the 4 quantum numbers as the address of an electron… country > state > city > street • electrons fill low energy orbitals before they fill higher energy ones

9. Quick intro, more later. • Principle quantum number (n) • describes the SIZE of the orbital or ENERGY LEVEL (shell) of the atom. • Angular quantum number (l) • a SUB-LEVEL (shell) that describes the type or SHAPE of the orbital • Magnetic quantum number (m) • the NUMBER of orbitals • describes an orbital's ORIENTATION in space • Spin quantum number (s) • describes the SPIN or direction (clockwise or counter-clockwise) in which an electron spins

11. 4f 4d 4p 4s 14(7) 10(5) 6(3) = level and sub-level = max. # of electrons = # of electrons = number of orbitals 2 (1) 32 3d 3p 3s 10(5) 6(3) 2(1) 18 2p 2s 6(3) 2(1) 8 1s 2(1) 2

12. Principle Quantum Number (n) or Energy Level • values 1-7 (ground state) used to specify the level the electron is in • describes how far away from the nucleus the electron level is • the lower the number, the closer the level is to the atom's nucleus and generallyless energy • maximum # of electrons that can fit in an energy level is given by formula 2n2(n = energy level)

15. Angular Quantum Number (l) or Sub-Levels just know this • determines the shapeof the sub-level • number of sub-levels equal the level number • ex. the second level has two sub-levels • they are numbered but are also given letters referring to the sub-level type • l=0 refers to the s sub-level • l=1 refers to the p sub-level • l=2 refers to the d sub-level • l=3 refers to the f sub-level

18. Magnetic quantum number (m) or Orbitals Electron Orbitals YouTube 1:37 the third of a set of quantum numbers tells us how many sub-levels there are of a particular type and their orientation in space of a particular sub-level only two electrons can fit in an orbital = electron

19. S sub-levelhas only 1 orbitalonly holds two electrons

20. P sub-levelhas 3 orbitalsholds up to six electrons

22. D sub-levelhas 5 orbitalsholds up to 10 electrons

24. F sub-levelhas 7 orbitals holds up to 14 electrons

27. Spin quantum number (s) • the fourth of a set of quantum numbers • number specifying the direction of the spin of an electron around its own axis. • only two electrons of opposite spin may occupy an orbit • the only possible values of a spin quantum number are +1/2 or -1/2.

31. “Rules” for Writing Electron Configurations • a method of writing where electrons are found in various orbitals around the nuclei of atoms. • three rules in order to determine this (don’t to know the name of the “rules”, just know the rules: • Aufbau principle • Pauli exclusion principle • Hund’s rule

32. Aufbau Principle • electrons occupy the orbitals of the lowest energy first • each written represents an atomic orbital (such as or oror ….) • electrons in the same sublevel/shell have equal energy ( same energy as ) • energy levels (1,2,3,4..) can overlap one another • ex: 4s orbital has less energy than a 3d orbital

33. Pauli Exclusion Principle Hamster video 1:00 actually incorrect as well, see next slide • only two electrons in an orbital • must have opposite spins • represents one electron • represents two electrons in an orbital

34. Hund’s Rules every orbital in a sublevel must have one electron before any one orbital has two electrons all electrons in singly occupied orbitals have the same spin.

35. Writing Orbital Diagrams

36. Energy

40. Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental) s orbitals d orbitals p orbitals f orbitals

44. Boron Atomic # 5 http://colossus.chem.umass.edu/genchem/whelan/class_images/Orbital_Energies.jpg

45. Boron ion (3+) Atomic # 5 http://colossus.chem.umass.edu/genchem/whelan/class_images/Orbital_Energies.jpg

46. Neon Atomic # 10 http://colossus.chem.umass.edu/genchem/whelan/class_images/Orbital_Energies.jpg

47. Bromine Atomic # 35 http://colossus.chem.umass.edu/genchem/whelan/class_images/Orbital_Energies.jpg

48. Bromine ion (1-) Atomic # 35 http://colossus.chem.umass.edu/genchem/whelan/class_images/Orbital_Energies.jpg

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