1. Chapter 8 Periodic Properties of the Elements

2. Electron Spin experiments by Stern and Gerlach showed a beam of silver atoms is split in two by a magnetic field the experiment reveals that the electrons spin on their axis as they spin, they generate a magnetic field ◦ spinning charged particles generate a magnetic field if there is an even number of electrons, about half the atoms will have a net magnetic field pointing “North” and the other half will have a net magnetic field pointing “South” 2

3. Electron Spin Experiment 3

4. Spin Quantum Number, m s spin quantum number describes how the electron spins on its axis ◦ clockwise or counterclockwise ◦ spin up or spin down spins must cancel in an orbital ◦ paired m s can have values of ±½ 4

5. Pauli Exclusion Principle no two electrons in an atom may have the same set of 4 quantum numbers therefore no orbital may have more than 2 electrons, and they must have with opposite spins knowing the number orbitals in a sublevel allows us to determine the maximum number of electrons in the sublevel s sublevel has 1 orbital, therefore it can hold 2 electrons p sublevel has 3 orbitals, therefore it can hold 6 electrons d sublevel has 5 orbitals, therefore it can hold 10 electrons f sublevel has 7 orbitals, therefore it can hold 14 electrons 5

6. Allowed Quantum Numbers 6

7. Quantum Numbers of Helium’s Electrons helium has two electrons both electrons are in the first energy level both electrons are in the s orbital of the first energy level since they are in the same orbital, they must have opposite spins nlmlml msms first electron 100+½+½ second electron 100-½-½ 7

8. Electron Configurations Electron Configurations the ground state of the electron is the lowest energy orbital it can occupy the distribution of electrons into the various orbitals in an atom in its ground state is called its electron configuration the number designates the principal energy level the letter designates the sublevel and type of orbital the superscript designates the number of electrons in that sublevel He = 1s 2 8

9. Orbital Diagrams we often represent an orbital as a square and the electrons in that orbital as arrows ◦ the direction of the arrow represents the spin of the electron 9 orbital with 1 electron unoccupied orbital orbital with 2 electrons

10. Sublevel Splitting in Multielectron Atoms the sublevels in each principal energy level of Hydrogen all have the same energy – we call orbitals with the same energy degenerate ◦ or other single electron systems for multielectron atoms, the energies of the sublevels are split ◦ caused by electron-electron repulsion the lower the value of the l quantum number, the less energy the sublevel has ◦ s (l = 0) < p (l = 1) < d (l = 2) < f (l = 3) 10

11. Penetrating and Shielding the radial distribution function shows that the 2s orbital penetrates more deeply into the 1s orbital than does the 2p the weaker penetration of the 2p sublevel means that electrons in the 2p sublevel experience more repulsive force, they are more shielded from the attractive force of the nucleus the deeper penetration of the 2s electrons means electrons in the 2s sublevel experience a greater attractive force to the nucleus and are not shielded as effectively the result is that the electrons in the 2s sublevel are lower in energy than the electrons in the 2p 11

12. Penetration & Shielding 12

13. Energy 1s 7s 2s 2p 3s 3p 3d 6s 6p 6d 4s 4p 4d 4f 5s 5p 5d 5f Notice the following: 1.because of penetration, sublevels within an energy level are not degenerate 2.penetration of the 4 th and higher energy levels is so strong that their s sublevel is lower in energy than the d sublevel of the previous energy level 3.the energy difference between levels becomes smaller for higher energy levels

14. Filling the Orbitals with Electrons energy shells fill from lowest energy to high subshells fill from lowest energy to high ◦ s → p → d → f ◦ Aufbau Principle orbitals that are in the same subshell have the same energy no more than 2 electrons per orbital ◦ Pauli Exclusion Principle when filling orbitals that have the same energy, place one electron in each before completing pairs ◦ Hund’s Rule 14

16. Electron Configurations of Multielectron Atoms n = 1 s orbital (l = 0) 1 electron H: 1s11s1 1s21s2 n = 1 s orbital (l = 0) 2 electrons He: n = 2 s orbital (l = 0) 1 electrons 1s 2 2s 1 Li: Lowest energy to highest energy

17. Valence Electrons the electrons in all the subshells with the highest principal energy shell are called the valence electrons electrons in lower energy shells are called core electrons chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the number of valence electrons 17

18. Examples For the following atom, write: ◦ the Ground State Electron Configuration ◦ Use short hand notation to write orbital Diagram ◦ Determine the core electrons and valence electrons  Carbon  Magnesium  Sulfur  Potassium

19. Electron configuration of transition metal and atoms in higher energy state For the following atom, write: ◦ the Ground State Electron Configuration ◦ Use short hand notation to write orbital Diagram ◦ Determine the core electrons and valence electrons  Cr  Br  Pd  Bi