Chapter 8§8.2 Organization of the Periodic Table§8.3 Electron ConfigurationPowerPoint PresentationSlide 5§8.3 Orbital Diagrams§8.3 Electron ConfigurationsSlide 8Slide 9Slide 10Slide 11§8.3 Electron Spin§8.3 Pauli Exclusion PrincipleSlide 14§8.3 Multielectron Atoms§8.3 Multielectron Atoms (≠ H)§8.3 Shielding and PenetrationSlide 20§8.3 Penetrating and ShieldingSlide 22§8.3 Subshell Order in Electron ConfigurationsSlide 24Slide 25§8.4 Classifying Electrons Based on ReactivitySlide 27§8.6 Periodic Trends§8.6 Atomic Radius§8.6 Atomic Radii (pm) of Main Group ElementsExample 8.5 – Choose the Larger Atom in Each Pair§8.6 Atomic Radii of Transition Metals§8.6 Effective Nuclear Charge§8.6 Screening & Effective Nuclear Charge§8.7 Electron Configuration & Ion Charge§8.7 Electron Configuration of Anions§8.7 Electron Configuration of Cations- Groups I, II and III§8.7 Electron Configuration of Cations- Transition MetalsWhich element’s +3 ion would have a [Kr]4d10 electron configuration?§8.3 Adding Electrons to Subshells§8.7 Magnetic Properties Atoms & IonsIs the Fe3+ ion paramagnetic?§8.7 Trends in Ionic Radius§8.7 Isoelectronic Ions§8.7 Cation Size§8.7 Anion SizeSlide 48Slide 49Slide 50Example 8.8 – Choose the Atom in Each Pair with the Lower First Ionization Energy§8.7 Irregularities in the IE TrendSlide 53Slide 54Slide 55§8.7 Successive Ionization EnergiesSlide 57§8.7 Electron AffinitySlide 59§8.7 Metals vs. NonmetalsSlide 61Example 8.9 – Choose the More Metallic Element in Each PairChapter 8PERIODIC PROPERTIES§8.2 Organization of the Periodic TableHow are elements arranged on the periodic table? Is it arbitrary? There are several less common versions of the periodic table.Ground state - all electrons are in lowest energy orbitals possibleElectron configuration – a notation for distribution of electrons into the orbitals of a ground state atom.The number is n; the letter is l ; the superscript is the number of electrons in that sublevel.§8.3 Electron ConfigurationLi: 1s22s14orbital with1 electronunoccupiedorbitalorbital with2 electrons§8.3 Orbital DiagramsOrbital diagram – electrons drawn as arrows, orbitals as boxesSame information as electron configurations plus electron spinAn orbital can hold 0, 1 or 2 electrons. If it holds 2, the spins must be opposed.5P = [Ne]3s23p3P has 5 valence electrons3p3PNe12345671A2A3A 4A 5A 6A 7A8A3s2§8.4 Electron Configuration and the Periodic Table§8.3 Orbital DiagramsMg: Z = 12Write the ground state electron configuration and orbital diagram of Mg.1.The atomic number gives the number of protons and electrons in a neutral atom.2.Add one electron to each box in a subshell then pair the electrons before going to the next set. For pairs, draw opposing arrows.1s 2s 2p 3s 3p 1s22s22p63s2 = [Ne]3s2§8.3 Electron ConfigurationsElectron configuration and orbital diagrams give the same basic information, but orbital diagrams give electron spins. Electron Orbital diagram configuration 1s 2s 2pBe 4 1s22s2N 7 1s22s22p3F 9 1s22s22p5Ne 10 1s22s22p6Electronconfiguration§8.3 Electron ConfigurationsChromium and copper are exceptions to the Aufbau Principle: 4s 3dCr 24 [Ar]4s13d5Cu 29 [Ar]4s13d10Don’t worry about writing electron configurations for lanthanides and actinides—there are too many exceptions.§8.3 Electron ConfigurationsKnow the electron configurations in pink:Main group electron configurations don’t have exceptions.Transition metal electron configurations are rife with exceptions.§8.3 Electron SpinThere is a 4th quantum number called the electron spin quantum number (ms).ms has two values: +½ and -½ (or ↑ and ↓). These values represent an electron’s two spins states.§8.3 Pauli Exclusion PrincipleTwo ways to phrase this principle:1) In an atom, no electrons have the same 4quantum numbers (n, l, ml , ms).2) At most an orbital can hold two electrons, and they must have opposite spins.Wolfgang Pauli (1900 – 1958)·1945 Nobel Laureate·Austrian-German-American-Swiss·Proposed the neutrino·Obsessed with the number 137 ·Carl Jung’s patient·Fled Germany to U.S. during WWII·Princeton faculty member§8.3 Pauli Exclusion PrincipleThe number of orbitals in a sublevel determines the maximum number of electrons it can hold:The s sublevel has 1 orbital, ∴ it can hold 2 electronsThe p sublevel has 3 orbitals, ∴ it can hold 6 electronsThe d sublevel has 5 orbitals, ∴ it can hold 10 electronsThe f sublevel has 7 orbitals, ∴ it can hold 14 electronsFor hydrogen: the sublevels in each principal energy level are degenerate: ns = np = nd = nf All other atoms: sublevel energies are split due to electron-electron repulsion. A sublevel with a lower l value has lower energy: ns < np < nd < nf §8.3 Multielectron Atoms= H ≠ H§8.3 Multielectron Atoms (≠ H)As Z increases:Electron-nucleus attraction increases.The number of electrons increases ∴ electron-electron repulsions increase.Outer electrons are “shielded” by inner electrons.§8.3 Shielding and Penetration§8.3 Multielectron AtomsWhy is a 1s orbital lower in energy than a 7s? 1s electrons are closer to the nucleus.1s electrons experience no shielding. 7s electrons have 85 lower energy electrons between them and the nucleus.Which orbital has a greater probability of being closer to the nucleus, a 2s or 2p electron? 2p, even though it’s higher in energy§8.3 Penetrating and ShieldingDistance from nucleusRadial Probability2s2pOn average, 2s electrons are further from the nucleus than 2p’s, but 2s electrons have a small chance of getting very close to the nucleus.Energy1s7s2s2p3s3p3d6s6p6d4s4p4d4f5s5p5d5f1.For non-H atoms, sublevels (same n, same l) are not degenerate.2.Penetration by ns electrons is so strong that for n ≥ 4, ns electrons are lower in energy than (n – 1)d electrons.3.The energy difference between levels becomes smaller at higher n values.§8.3 Subshell Order in Electron Configurations1s2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d7s1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d•List each energy shell on a row listing the subshells (s, p, d, f) for each shell in order of energy.•To see the order the subshells are filled, draw parallel upper right to lower left diagonals.•Too