Atomic Structure Electron Configuration

1. Atomic StructureElectron Configuration Scandium 3-D video (2:31) 3-D Graphic Examples of Atomic Orbitals Topic 12.1.3 – 12.1.6

3. Review 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 to an atom, the electron will absorb energy and "jump" to a higher energy level • heating a chemical with a Bunsen burner is enough energy to do this

5. Review • after a short time, this electron will spontaneously "fall" back to a lower energy level, giving off a quantum of light energy called a photon • 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. • quantum is the amount of energy required to move an electron from one energy level to another

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 • the first three of these quantum numbers (n, l, and m) represent the three dimensions in which an electron could be found • the fourth quantum number (s) refers to a certain magnetic quality called spin

9. 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 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 less energy • maximum # of electrons that can fit in an energy level is given by formula 2n2

14. Angular Quantum Number (l) or Sub-Levels • 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

16. 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

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

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

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

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

25. 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.

30. “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: • Aufbau principle • Pauli exclusion principle • Hund’s rule

31. 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 ) • principle energy levels/shells (1,2,3,4..) can overlap one another • ex: 4s orbital has less energy than a 3d orbital

32. 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

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

34. Writing Orbital Diagrams

35. Energy

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

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

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

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

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

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

48. ?

50. Writing Electron Configurations • To write out the electron configuration of an atom: • use the principal quantum number/energy level (1,2,3, or 4…) • use the letter term for each sub-level (s,p,d, or f); • don’t worry about orientation such as x,y,z axis but you do have to be able to draw these for IB • use a superscript number indicates how many electrons are present in each sub-level • hydrogen =1s1. • Lithium =1s22s1. • don’t write anything for spin