Chapter 8 Chapter 8 1.1 Orbital Energies 1.2 Orbital Energies First three quantum numbers uniquely describe orbitals 1s 2s 3s 4s 2p 3p 4p 3d 4d What general principle explains orbital energies? Which orbital has higher energy, 1s or 2s? Why? 4f level: contains orbitals with the same n Which orbital has higher energy, 2s or 2p? Why? y sublevel: contains orbitals with the same n and l What determines the relative energies of these orbitals? Which are lower in energy? Which are higher in energy? Chapter 8 Which orbital has higher energy, 2px, 2py or 2pz? Why? Chapter 8 1.3 Orbital Energies 1.4 Orbital Energies Which orbitals have higher energy, 3s, 3p or 3d? Why? (what do you need to know?) Will the five 3d orbitals have the same energy? Why or why not? 1 Chapter 8 Chapter 8 1.5 Energy of Orbitals: Summary For the same type of orbital (same ______), energy increases as n increases ((1s < 2s < 3s < 4s…)) For the same n, energy increases s < p < d < f (3s < 3p < 3d) All orbitals of the same sublevel have the same energy (they are _________________)) (3px = 3py = 3pz) 2.1 Magnetic Properties: Electron A physical phenomenon: spinning, p g charged g p particles produce magnetic fields Spinning electrons produce tiny magnetic fields Electrons can spin in one of two directions 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p Chapter 8 2.2 Magnetic Properties: Electron Diamagnetic Paramagnetic Chapter 8 2.3 The 4th Quantum Number Electron spin, ms: ms = ½ or -½ Pauli exclusion principle: 2 Chapter 8 Chapter 8 3.1 Electron Configuration Rules Electrons fill the lowest energy orbital first (Aufbau principle) 3.2 Electron Configurations Two notations for the arrangement of electrons in atoms Two electrons (max) per orbital spdf notation Maximize parallel spins when filling a sublevel Fill orbitals of equal energy with one electron each before pairing up Why don’t we double-up first? Chapter 8 noble gas notation Chapter 8 3.3 Electron Configurations Hydrogen 3.4 More Examples Provide the electron configurations (in spdf and noble gas notation) phosphorus Lithium vanadium Oxygen iodine 3 Chapter 8 Chapter 8 3.5 More Examples: Ions 4.1 Periodic Table Organization (a) S2– So does S2– = Ar? (b) Br – (c) Al3+ Chapter 8 Chapter 8 4.2 Periodic Table Organization s-block p-block d block d-block atoms where an s sublevel is being filled atoms where a p sublevel is being filled atoms where a d sublevel is being filled 4.3 Periodic Properties You will need to know the following: 1 Definitions and chemical equations where 1. appropriate Valence electrons 2. Periodic trends moving up and down and left to right across the periodic table Core electrons 3. Explanations of the trends Same group = same number and type of valence electrons 4. How the atomic properties affect chemical properties 4 Chapter 8 Chapter 8 4.4 Effective Nuclear Charge Valence electrons don’t “feel” the full charge of the nucleus Valence electrons are shielded from nuclear charge (Z) Take the case of Li 1s22s1 4.5 Effective Nuclear Charge Zeff: the positive charge “felt” by a valence electron An crude approximation: Zeffff = Z – core electrons where Z = atomic number Zeff increases across the periodic table Chapter 8 Lithium: Zeff = 3 – 2 = 1 Carbon: Zeff = 6 – 2 = 4 Fluorine: Zeff = 9 – 2 = 7 Sodium: Zeff = Silicon: Zeff = Chapter 8 4.6 Atomic Size 4.7 Atomic Size The distance from the nucleus to the edge of the outermost electron Periodic trend: Explanation: 5 Chapter 8 Chapter 8 5.1 Ionization Energy (IE) The energy required to remove an electron from a gaseous atom A(g) + energy A+(g) + e- 5.2 Sign Conventions Energy absorbed (in) = a positive value + 165 kJ Energy required (input, raw material) Energy released (out) = a negative value - 165 kJ Energy produced (output, product) Energy input required The sign tells us which way energy is going The magnitude tells us how much energy is involved Chapter 8 Chapter 8 5.3 Ionization Energies 5.4 Ionization Energies: Summary 6 Chapter 8 Chapter 8 5.5 Successive Ionizations IE1 IE2 IE3 Na 495 4560 Mg 735 1445 IE4 IE5 IE6 IE7 5.6 Successive Ionizations For Mg, 2nd IE > 1st IE For Al, 3rd IE > 2nd IE > 1st IE Why? 7730 Al 580 1815 2740 11600 Si 780 1575 3220 4350 16100 P 1060 1890 2905 4950 6270 21200 S 1005 2260 3375 4565 6950 8490 27000 For Mg, 3rd IE >>> 2nd IE For Al, 4th IE >>> 3rd IE Example: Na(g) + IE1 Na+(g) + e- Why? Na+(g) + IE2 Na2+(g) + e- Chapter 8 Chapter 8 5.7 Electron Affinity The energy released when an electron is added to a gaseous atom A(g) + e- A-(g) + energy A free electron is not a stable beast. It would rather be associated with an atom. 5.8 Electron Affinity Predictions A(g) + e- A-(g) + energy Across a period (left to right): Should it get easier or harder to add an electron? Down a group: Should it get easier or harder to add an electron? If it’s “easy” to add an electron, is the EA a large negative number or a small negative number? Deviations from the general trends 7 Chapter 8 Chapter 8 5.9 Electron Affinity Trends Chapter 8 5.10 Electron Affinity Summary Chapter 8 6.1 Ionization: Change in Size Why does the size decrease? 6.2 Ionization: Change in Size Why does the size increase? 8
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