Electrons exhibit a magnetic field We think of them as spinning They can spin only two ways think of it as left or right Spin quantum number ms can be 1 2 or 1 2 Magnetic Properties come from additive effects of electron spins Diamagnetic all electrons are paired Paramagnetic 1 or more unpaired electrons Ferromagnetic real magnets unpaired electrons all lined up in the same direction Pauli Exclusion Principle No two electrons in an atom can have the same 4 quantum numbers n m define an orbital Therefore an orbital can hold two electrons with opposite spins because m s can only be 1 2 or 1 2 Orbital Energies Why For a single electron it only depends on how far from the nucleus For many electrons e e repulsions also play a role and differ depending on orbital shape Orbital Energies For most atoms Energy increases as n increases 1 2 3 4 Energy increases as subshells go from s p d f Electron Configurations General Rule electrons fill lowest energy orbitals first Sodium Na as an example Na has 11 electrons Fill 2 electrons per orbital till you run out A box represents an orbital A arrow represents an electron Electron Configurations Three Notation Types 1 2 spdf or spectroscopic notation List subshells and how many electrons they contain 1s22s22p63s1 3 Noble gas notation short Ne 3s1 Where Ne 1s22s22p6 Electron Configurations and the Periodic Table Examples using Electron Configuration Simulation Periodic Blocks Hund s Rule using the p block n value increases as you move down table Anomalies Cr and Cu Electron Configurations and the Periodic Table I Electron Configurations and the Periodic Table II Electron Configurations and the Periodic Table III What would the periodic table look like if the rules were different For example what if electrons could only have a spin of 1 2 and not 1 2 Sketch it Notes There s no known reason electrons have spin or have only two of them The other stuff about orbitals is theoretically derived from Schrod Equn but the whole spin thing is just something we see Can t explain it just know it s true Like gravity or Coulombic attractions The reason why different subshells have different energies for example The energy of the 2s subshell has to do with how well the 2s electrons are attacted to the nucleus minus how much they are repelled by the 1s electrons Same thing for the 2p electrons Difference is the 1s electrons repel the 2p electrons more than the 2s electrons so the 2p electrons are less stable and higher energy Same reasoning happens when you go to higher subshells e g d p Why does the 4s subshell come before the 3d subshell The reason above about d being higher in energy plus the fact that as you go up in n value the orbital energies all get closer together So 2 is much higher than 1 3 is less higher than 2 4 is not much higher than 3 etc This comes from the En constant n 2 The reason for Hund s Rule there is less e e repulsion if electrons are in different orbitals because they are in different Places That s why they go to different orbitals in a subshell first I don t know why they go with the same spin Have them fill in the blanks for a set of elements as you use the simulation Be sure to include a Hund s rule one B C N or something like it Answer to hard question the pt looks the same but is half as wide for each block because each orbital can only hold a single electron Predicting Electron Configurations Predicting Electron Configurations Predicting Electron Configurations Predicting Electron Configurations Electron Configurations of Cations Electron Configurations of Anions Transition Metal Cations Lose s electrons first Diamagnetic vs Paramagnetic Elements Periodic Properties of the Elements All Depend on energies of outermost orbitals Atomic Size Ionization Energy Electron Affinity Ion Size Trends in Orbital Energies Why do energies decrease moving left to right General Periodic Trends Which atom is the smallest of all 1 2 3 4 H He Cs Rn Which of these atoms is largest 1 2 3 4 K Ca Rb Sr Which atom has the largest ionization energy 1 2 3 4 K Ca Rb Sr Which ionization energy for Mg will see the largest jump 1 2 3 4 1st 2nd 3rd 4th Why the breaks in the line Which atom has the smallest common ion 1 2 3 4 H Na F Cl Which atom has the largest common ion 1 2 3 4 Na K F Cl Which of the following isoelectronic species is smallest 1 2 3 4 5 Mg2 Na Ne FO2 Which of the following isoelectronic species has the lowest ionization energy 1 2 3 4 5 Mg2 Na Ne FO2